Chemistry 101L

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1.
Chemistry 101L
Fall 2009
Dr. Calhoun
Dr. Fulmer
Dr. Coker
M 3:00-5:10, T 10:20-12:30, Th 10:20-12:30,
T 12:40-2:50, W 12:40-2:50
Th 3:00-5:10
Chemistry 101 Lab – Spring 2009
Table of Contents
Lab Safety Rules
Good Lab Practices
Introduction
Making a Graph in Excel
Grading Policies
Equipment/Check-in Sheet
Glossary (Primarily for use in lecture)
Supplementary Materials (Primarily for use in lecture)
Dates
Aug. 17-20
Aug. 24-27
Aug. 31-Sept.3
Sept.8-10
Sept. 14-17
Sept. 21-24
Sept. 28-Oct. 1
Oct. 5-7
Oct. 12-15
Oct. 19-22
Oct. 26-29
Nov. 2-5
Nov. 9-12
Nov. 16-19
Nov. 31-Dec. 3
Lab
Check-in; Classifications
Separation of Mixture
Metric Estimates and Volume of Water
Drop
Density and Reliability of Data
Solubility & Conductivity; Week 1 of
zinc chloride
Zinc Chloride, Week 2
Types of Reactions
Writing Chemical Reaction Equations
Vinegar Analysis
Nomenclature Lab
Lewis Structures
Prep and Properties of Sulfur Dioxide
and Nitrogen Oxides
Acid-Base I
Acid-Base II
pH, Buffers, Checkout
2
3
6
7
9
10
11
92
120
Page
12
16
21
28
32 and 39
38
41
49
58
61
62
74
81
85
89
Name:_______________________________ Course & section #:_____________ Date:___________
(print)
LABORATORY SAFETY RULES
1)
Responsible behavior is essential. Think about the consequences of your actions or inactions. Make
note of the safety precautions written in the lab text and told to you by the laboratory instructor. If
you are unsure about a procedure, ASK!
2)
WEAR APPROVED EYE PROTECTION AND AN APRON AT ALL TIMES IN THE LAB.
Goggles and safety glasses are available in the bookstore. If you wear glasses, be sure your eye
protection fits over your glasses. Contact lenses are discouraged. If you must wear them, then you
must wear goggles. Your shoes and clothing should also provide protection, i.e. wear closed-toe
shoes whose "uppers" repel spills; confine long hair and loose clothing; no bare midriffs; do not wear
sleeveless shirts, short-shorts, short skirts, or high-heeled or slick-soled shoes.
3)
In case of fire, accident, or chemical spill, notify the instructor at once.
a) Wet towels can be used to smother small fires.
b) In case of a chemical spill on you body or clothing, wash the affected area with large quantities of
running water. Be sure to first remove the clothing that has been wetted by chemicals to prevent
further reaction with the skin.
c) If you should get a chemical in your eye, wash with flowing water from the eye wash for at least
15 minutes. Get medical attention immediately.
d) Report all injuries to your instructor at once.
4)
Note the locations of the safety equipment listed below. Make sure you know how to use them.
a) Eye Washes b) Safety Shower
c) Fire Extinguisher
d) Blanket
e) Telephone f) Sinks g) Hoods h) Exits
5)
EMERGENCY HELP IS AVAILABLE BY DIALING 911. Telephones are available in the C
hallway between 445 & 446 and in the middle (B) hall outside 423.
6)
Do not eat, drink, or chew on anything in the laboratory. This includes food, gum, tobacco, pens,
and other chemicals. Also, you should not handle many lab chemicals. They may absorb through the
skin or be ingested if you lick your fingers. Wash thoroughly before you leave the lab, even if you
haven't "touched" anything.
7)
Avoid breathing fumes of any kind.
a) To test the smell of a vapor, "waft" the vapor toward you. Test only those vapors specified.
b) Work in a hood if there is a possibility that noxious or poisonous vapors may be produced. If
unexpected vapors are produced during an experiment, move the experiment to a hood
immediately.
8)
Keep your work area clean at all times. Clean up spills and broken glass immediately, whether at
your bench or in the common area. Clean your work space, including wet-wiping the bench top and
putting away all chemicals and equipment, at the end of the period.
9)
Perform no unauthorized experiments. This includes using more reagent than instructed.
10)
Never work alone in the laboratory.
11)
Always pour acids into water when mixing. Otherwise, the acid can violently spatter. "Acid into
water is the way that you oughter."
3
12)
Do not force a rubber stopper onto glass tubing or thermometers. Use a split-hole stopper if possible.
If not, lubricate the tubing and the stopper with glycerol or water; use paper or cloth toweling to
protect your hand; and grasp the glass close to the stopper.
13)
Beware of hot iron rings and hot glass. They "look" cool long before they can be handled.
14)
Dispose of small quantities of excess non-toxic liquid reagents by flushing them down the sink with
copious quantities of water. Dispose of solids and toxic liquids in the waste containers provided.
Never use your eyedroppers to remove liquids from a reagent bottle. Always pour reagents into your
own clean containers. Never return reagents to the bottle.
15)
Spatters are common in general chemistry laboratories. Test tubes containing reacting mixtures or
those being heated should never be pointed at anyone.
16)
Carefully read the experiment before coming to lab. Review techniques to be used and safety
precautions to be taken.
17)
Finally, and most importantly, THINK about what you're doing. If you doubt that what you're doing
is safe, don't do it. Check with your instructor. Ask questions.
Minimization of exposure of all students to the chemicals we work with in lab is an important goal
for both the instructor and the students. The experiments we pick, chemicals we choose to use, and
safety procedures we emphasize are all a part of attaining that goal. However, be aware that the
effects of many chemicals are different for different people. Certainly, we all know about allergies,
and have heard about how different medicines can interact when taken together and cause disastrous
results. Analogously, certain medical conditions such as asthma, allergies, etc. can be worsened by
exposure to certain chemicals. Please inform your instructor about any allergies or medical
conditions of which they should be aware.
Some chemicals which are not generally considered toxins are reproductive toxins. They may affect
the fetus directly or affect the ability of the parent (fathers, too) to produce healthy offspring. Also
realize that not all chemicals have been tested to see if they are reproductive toxins. Anyone who is
sexually active needs to be diligent in their efforts to minimize their chemical exposure in the lab. If
you are or become pregnant, you need to notify your instructor as soon as possible. Notification
forms are available in your manual or from your instructor.
.........................................................................................................................................................................
I have read and understand the Laboratory Safety Rules and have retained a copy for my reference.
________________________________
(signature)
________________
(course & section)
__________
(date)
Please list any allergies or medical conditions (including pregnancy) of which the instructor should be aware.
Please indicate if you wear contacts.
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
Revised: RKC 2/04
4
Notification of Pregnancy
I hereby notify my lab instructor that I am pregnant. I give permission to the laboratory
supervisor to send a list of the chemicals I will be using this semester to my doctor. I also authorize
my doctor to discuss with my instructor or the laboratory supervisor any hazards associated with
exposure to these chemicals. I also authorize my doctor to respond in writing to the lab supervisor
with recommendations for extra precautions they recommend, including exclusion of specific
chemicals from the experiments I do.
Student's Name ________________________________________
(please print)
Student's Signature _____________________________________
Student's Social Security # _________________
Student's Date of Birth ____________________
Doctor's Name _________________________________________
Doctor's Address _______________________________________
_______________________________________
_______________________________________
Doctor's Phone # ________________________
Course _______________
Lab Instructor _________________________________
Date _________________
5
GOOD LABORATORY PRACTICES
The list below is intended to state how we can keep the laboratory a pleasant place in which
to work and in which the results we obtain will be reliable.
1) Read and understand the experiment before coming to lab.
2) Equipment Care. Laboratory glassware is expensive and easily broken. Handle it carefully,
wash it thoroughly with soap and water, and return it to your drawer before leaving the lab.
Ordinary dish-washing detergent and brushes are available at the large sink at the rear of the
lab. Clean glassware immediately after use. Use the brush to scrub it clean, rinse with a lot
of tap water, and then rinse with distilled water.
3) Bench Tops. All work spaces should be wiped clean with a damp sponge at the end of the lab
period.
4) Stock Chemicals. Most experiments require chemicals that are supplied in stock bottles as
solids or liquids. These stock chemicals are placed on the benches in the center of the lab.
Carry a beaker or other container to the island and get what you need.
5) Do not remove the stock bottles from the supply table.
6) Never use your droppers to remove liquids from a bottle. Always pour reagents into your
own clean containers
7) If you accidentally pour out more of a chemical than you need, DON'T put it back into the
supply bottle! Doing so may contaminate the entire bottle and ruin everyone else’s data.
8) Conservation. Most chemicals used in the laboratory are very pure and expensive. Don't get
more than you need. If you don’t know how much you need, don’t take any. If you
accidentally take too much, find out if anybody else can use it before you dispose of this
excess.
9) Disposal. Certain chemicals must not be poured down the drain or placed in the trash cans.
There will be specially-marked bottles for these materials. Be certain to use these containers.
10) Dispose of small quantities of excess non-toxic liquid reagents by flushing them down the
hood sink with large amounts of water.
11) Dispose of paper, match sticks, metals, and non-hazardous solid chemicals in the dry trash
cans. Broken glass should be placed in the specially marked container, not in the trash can.
12) Balances. Balances are expensive and easily damaged. Never place any chemical directly on
the balance pan. The best procedure for weighing a chemical is to weigh it directly into the
container in which it is to be used. Alternately, weigh into a small glass beaker. If any
chemical is spilled on or near the balance, use the paint brush near the balance to clean up the
chemical.
13) Clean Team. Each week students will be appointed to clean the areas used by everyone in the
lab, to ensure that all extra equipment is put away, and to check that individual work areas
are clean. On the day that you are responsible for cleaning and inspection you will be the last
to leave the lab.
6
INTRODUCTION
General Procedure
You will arrive in lab and prepared to start the laboratory work at the scheduled time. Being
prepared means that you have read and understood the general features of the lab, and that you will
have completed any specific pre-lab activity that is assigned as a part of the lab. You will also be
wearing your eye protection and apron.
Your lab instructor may or may not give a short talk of up to 10 minutes that explains the
hazards of the experiment, any special difficulties with the experiment and demonstrate any new
techniques. Your instructor may choose to inspect any pre-lab assignment before allowing you to
work. You will then complete the work assigned in the lab exercise, keeping careful written records
of what you do. Before you clean-up at the end of the lab, you should be sure your data is reliable.
While you work in the lab, you must be conscientious about working safely. Anyone violating
safety rules may be expelled from the laboratory and not allowed to return.
Your final product from the lab will be your lab report, which may be due at the end of the
experiment or at the beginning of the next laboratory period as announced by your instructor. The
graded lab report will be returned at the beginning of the lab after the report is submitted. Look over
the graded report before you submit the next lab report.
Data Records
You will soon realize that collecting laboratory data is extremely time consuming, and, as the
saying goes, "Time is money!" Consequently, laboratory data are valuable and should be treated that
way. Data are not scribbled on loose sheets of paper where they may be misplaced. Rather, they are
recorded in ink directly onto your report sheet. All data are carefully labeled and organized in tables
so that they are easily read and understood months or years after the data are collected. If a mistake
is made, a single line is drawn through the mistaken datum, and the corrected value is written beside
it.
We do not ask that you buy a special laboratory notebook, but we do ask that you record data
in ink directly on the Report Sheet that you turn in and that you write legibly, label all entries
carefully, and draw a single line through mistaken entries rather than obliterating or erasing them.
Notes on Laboratory Report Sheets
Different instructors will likely provide different guidelines, which you are expected to follow when
completing your report sheets. Below are a few general guidelines that should help you complete
your report sheets.
Sample calculations: It is not necessary for you to include every single calculation that you
performed in your report sheet. If a particular type of calculation is reported several times, you
should include one sample calculation. Your sample calculation should include actual numbers from
one of your trials; clearly show all steps performed and include units. Remember to use the proper
number of significant figures. Results from calculations should be summarized in a table(s).
7
Tables:
Tables are used to present information in an organized way in less space. The labels on
columns and/or rows save us from repeating the same information over-and-over in sentences. Also,
a table may be arranged so numbers that we want to compare are close together. It is appropriate to
use a table if the same type of measurement is made several times on different samples, if several
measurements are made on the same sample and if a measured and calculated value are reported for
samples.
To decide how to make a table, start with a list of all measurements and calculations that you
will be making. Remember to include room for multiple trials and average values, if appropriate. If
appropriate, have room for classmates’ data. Now organize your list into things that are similar or
data that should be compared. Tables columns/rows do not have to be listed in the same order that
the measurements were made. For example, it is easier to subtract the initial volume from the final
volume if the final volume is listed first, and calculated and measured values can be compared easily
if they are close to each other on the table, but you might do the calculation at a time very different
than when you did the measurement. For tables with more than 3 or 4 measurements/calculations, it
is usually preferable to have the measurements and calculations as row headings and the individual
trials as column headings.
Your table should have a title and the label on each column and row should clearly
indicate what is in the column/row, including units.
Graphs:
The independent (manipulated) variable is plotted along the x-(horizontal) axis, while the
dependent variable is plotted along the y-(vertical) axis.
Construct the graph so that the data fill as much space on the graph paper as possible.
Remember you do not have to start at (0,0); in fact, we seldom do. When dividing the scale always
choose values for the major divisions that make smaller subdivisions easy to interpret. Be sure each
division along each axis represents equal amounts, but the amount does not have to be the same for
both axes. The size of the minor subdivisions should permit easy estimation of all significant figures
in the data.
Place a dot for each data point at the appropriate place on the graph. Circle data points.
Draw a straight line that best fits your data. The line does not have to pass through all data
points, or through any of them, but it should pass as closely as possible to all of them. There should
be about the same number of points above the line as below the line. Do not connect the dots.
You should title your graph, clearly label both axes (including units) and label the value of
all major divisions.
If different sets of data are plotted on the same graph, different symbols should be used for
each set and a legend prepared identifying which symbol belongs to each data set.
Summary statement:
At the end of most labs you will be asked to write a summary statement. The summary
statement is not a review of what you did but rather a place to discuss what you learned, any
problems you had. It is a place to explain any cross-outs on your data sheets. When things don't go
as expected or your numerical answers aren't close to expected results, you should discuss this in the
summary statement. You should try to figure out what went wrong and explain your conclusions.
8
To Make And Edit A Graph In Excel
Have your data typed into a spreadsheet, with the independent variable in one column and the
dependent variable in the other. The independent variable may be thought of as the one you
manipulate while the dependent responds to the manipulation. The independent should always go on
the x-axis.
Click and drag across the two column headings. From the Insert menu choose Chart and then choose
X-Y(scatter) as a type. For the Standard SubType, choose the icon with only poiints and no lines or
curves. Choosing Next and then Next again. At this window, fill in the Title for the Chart and the
titles for the x and y axes. Choose Next again and then pick As new Sheet and Finish.
EDITING THE GRAPH
Much editing can be done by double/right clicking in region you want to edit. Note that the little
line that crosses an axis (if no grid is shown) is called a tic mark. Also, scientific data is almost
never shown in bar or pie graphs if the audience is another scientist.
Always do the following:
Make the background white/colorless. Double/right click in the gray area, away from any lines or
data points. A Format Plot Area Dialog box should appear. Choose Colors and Lines and then from
the pull-down color menu choose No Fill.
Unless you have more than one set of data shown on the same graph, delete the Legend
(which shows a type of point and the data type that goes with it). Delete this simply by choosing it
and pressing your delete button.
If the data is supposed to be linear, show the best-fit line and its equation. To do so,
choose Add Trendline from the Chart Menu. A dialog box will appear. Choose Linear under Type
and then select Options and mark the “display equation on chart” box and the “display R-squared
value on chart” box. Never set the intercept to zero. You may reposition the equation and R2 values
shown on the graph by dragging them. (Note: a good set of linear data will have an R2> 0.97!)
To edit the following, choose View>Formatting Palette first, then
To Edit the Axis or Graph Title, Under the Title section there is a pull-down menu which lets
you choose the Main Title, x-axis title or y-axis title. You many then edit it in the box just below
the pull-down menu and when you press return, your changes will show. Note that any changes
(such as bold or subscript) apply to the whole title.
You may add grid lines under the Grid section. (x and y gridlines are added separately from
each other, click on both)
To change how many decimal places are shown in the tic mark label. Select the axis whose
labels you want to change and then under the Number section of the Formatting Palette choose
the icon with the left pointing arrow to add another decimal place or the icon with the right arrow
to decrease the number of decimal places.
9
Grading and Grading Policies
1.
You must be at your laboratory station and ready to begin work at the beginning of each class
period. Anyone who misses more than 2 lab periods or doesn't turn in reports for more than 2 lab
periods, whether excused or unexcused, will receive a failing grade and must repeat the course. It is
not possible to make up a missed experiment.
2.
If you know in advance that you will be unable to attend a scheduled laboratory period, you
may be able to attend a different laboratory section as a guest. Arrange with your instructor a week
ahead of time for this.
3.
Your laboratory grade is based on the quality of your written laboratory reports. However,
instructors may reduce your grade because of lateness, unsafe laboratory practice, leaving your work
area unclean or other failure to follow good laboratory practice, or intellectual dishonesty.
4.
Your laboratory grade is reported to your lecture instructor and is counted as ~25% of your
grade for CHEM 101.
Intellectual Dishonesty
I have seen a number of types of intellectual dishonesty in CHEM 101 Lab. These include:
a. Falsified data - This is any report of results that is consciously different from what was observed.
b. Dry Labs - These are synthetic lab reports that have no reference to actual lab work performed.
For this reason, lab reports will never be accepted from students who were not seen to be performing
the lab work.
c. Plagiarism - This is outright copying of someone else’s results with no further comment.
These and similar offenses will be treated as a serious breach of ethics and handled accordingly. The
minimum penalty for any such behavior will be a zero for the experiment at hand, and more serious
penalties such as a directed grade of E for the course or a recommendation for expulsion from the
university may be considered.
Collaboration
The co-operation of students working the lab and outside the lab is encouraged. It is a wellknown technique for improved learning and a useful skill for life after college. Referring to outside
sources and people is also encouraged and is a kind of collaboration. However, when your lab report
includes the results of collaboration, it may do so only under the following circumstances:
a. The source of data you used should be clearly identified, whether it is the class as a whole,
Joanne Doe (your next hood neighbor), or the “CRC Handbook of Chemistry and Physics”.
b. If your calculation methods or your logical methods were merely borrowed from
someone, then there must be an acknowledgment of this source appended. However, if you asked
someone else how to do the calculations or logic, but now understand how and why to do them as a
result of this discussion, then you may treat them as your own.
10
Equipment List for Chemistry 101
Name(s)
Hood & Desk #s
Lab Section
Check your desk to see that you have the following items, clean and unbroken. Your
instructor will tell you where to place extra items. Replace any broken or missing
equipment.
one 20 mL beaker
one 50 mL beaker
one 100 mL beaker
one 150 mL beaker
two 250 mL beakers
one 400 mL beaker
two evaporating dishes
one 125 mL Erlenmeyer flask
one 250 mL Erlenmeyer flask
one funnel
one scoopula
one triangle
one test tube clamp
1 test tube rack
10 test tubes, 13 mm x 100 mm
one test tube, 25 mm x 200 mm (8
inch)
one Florence flask or a second 250
mL Erlenmeyer flask
one pair crucible tongs
one watch glass
one Pasteur pipet (looks like a long
medicine dropper), with bulb
one medicine dropper
one glass stirring rod
Check the large locker by your drawer for the following:
one ring stand
one ring
one utility clamp
one wire gauze
one bunsen burner and tubing
Check in:
Student’s signature(s)
Date
******************************************************
Check out:
Student’s signature(s):
Date
Instructor’s signature
11
Exercise 1: Check-in and Classification Systems
Introduction: Chemistry is about matter--all of it: air, water, dirt, food, rocks, stars--even Rock
Stars and people like you! That's a lot of matter. As you might imagine, the study of matter has
resulted in a lot of information. One way to cope with large amounts of information is to deal with
categories rather than individual facts. That is how chemists deal with facts about matter. Rather
than talking about iron and copper and aluminum and gold, chemists talk about metals and
summarize information about all of these materials. You use the same strategy when you say
something like "Birds lay eggs." You don't say that chickens lay eggs and ducks lay eggs and
robins lay eggs and wrens lay eggs. You categorize all those animals as birds and say "Birds lay
eggs." This process of placing similar things in categories is classification.
The key to classification is noticing regularity in diversity. Iron, copper, aluminum, and
gold are all solid at ordinary temperatures, can be hammered into thin sheets, and are good
conductors of heat and electricity. These shared properties define the category, metal, and when we
call something a metal, we are classifying it.
Purpose: To develop a classification scheme.
Procedure: We will consider ways that animals with similar characteristics might be grouped
together. Answer the questions and extend the classification scheme as you read along.
Consider these animals: robins, bats, spiders, whales, goats, ostriches, giraffes, chipmunks,
cardinals, bees, skunks, trout, bullfrogs, black bears, ants, kangaroos, tuna, deer, nightcrawlers and
tigers. The first step in developing a classification scheme is to look for some characteristic that is
shared by some of the animals but not others. The characteristic can be anything; e.g., it may
describe where the animals normally live, the kind of food they eat, some physical characteristic
such as number of legs or wings, or what covers their body. Before turning the page, list as many
characteristics as you can that could be used to separate these animals into groups.
1._____________________________________________________________________
2._____________________________________________________________________
3._____________________________________________________________________
4._____________________________________________________________________
12
Some of the characteristics I thought of are a) number of stomachs the animal has, b) size, c)
whether the animal has a bony skeleton, d) number of teeth the animal has, e) whether the animal
lays eggs, f) the animal's color, and g) whether the animal can fly.
After identifying characteristics that can be used to separate the animals into groups, the
next step is to pick one characteristic and use it to make a first separation. The characteristic used
first is somewhat arbitrary. Any characteristic could be used for a first separation, but some are
more useful than others. One consideration in selecting a characteristic to use for classification is
how easily you can make a distinction on the basis of the characteristic. For example, I decided that
the animal's color wouldn't be very useful as a basis for classification because I know that cows,
ducks, chickens, and goats come in a variety of colors. I'll settle for "whether the animal lays eggs."
Once you settle on which characteristic to start with, use that characteristic to separate the
animals into groups or subclasses. The resulting data are normally organized into a branching
pattern called a "tree diagram." We'll build one as we go:
All animals in the list
Does it lay eggs?
Yes
No
Bullfrogs, tuna, ants,cardinals, bees,
robins, spiders, ostriches, trout,
nightcrawlers
bats, black bears, tigers,deer, whales, goats,
giraffes, chipmunks, skunks, kangaroos
From this point on we again pick a characteristic and use it to split a subclass (we do NOT
start over with the original list of animals). For each subclass formed, repeat the process of a)
identifying characteristics that are shared by some of the animals, but not all, b) selecting one of
these characteristics to use, and c) use the characteristic to separate the animals into groups.
Let's try a further separation of the egg-layers by whether or not they fly.
All animals in the list
Does it lay eggs?
Yes
No
Bullfrogs, tuna, ants,cardinals, bees,
robins, spiders, ostriches, trout,
nightcrawlers
bats, black bears, tigers,deer, whales, goats,
giraffes, chipmunks, skunks, kangaroos
Can it fly?
Yes
Cardinals, bees,
robins
No
Bullfrogs, tuna, ants, spiders,
ostriches, trout, nightcrawlers
13
Notice that asking the same question about the non-egg laying animals is not very useful,
since only bats fly. It is often the case that a question asked about one subgroup is of little or no use
for separating members of a different subgroup. When that is true, the question is normally
eliminated from the final classification scheme. and a better characteristic is suggested. Let's try
number of legs as a characteristic to split the subclass that does not lay eggs:
All animals in the list
Does it lay eggs?
Yes
No
Bullfrogs, tuna, ants,cardinals, bees,
robins, spiders, ostriches, trout,
nightcrawlers
bats, black bears, tigers,deer, whales, goats,
giraffes, chipmunks, skunks, kangaroos
Number
of legs?
Can it fly?
Yes
No
0
Cardinals, bees,
robins
Bullfrogs, tuna, ants, spiders,
ostriches, trout, nightcrawlers
whales
2
bats
4
black bears,
skunks, tigers,
deer, goats,
giraffes, chipmunks,
skunks, kangaroos
When to end the classification is a matter of choice. Usually the process is continued until
all things in a group are very similar. Most of the subgroups above still need further classification.
For instance whether the animal has gills or not would separate tuna and trout from bullfrogs, ants,
nightcrawlers, spiders, and ostriches. Likewise animals with hooves or not would further sort black
bears, chipmunks, skunks, tigers, and kangaroos from, deer, goats, and giraffes.
Most of the questions asked so far can be answered "yes" or "no" and result in two, and only
two, groups. Most classification schemes are built around binary separations of this kind.
However, questions that result in more than two groups are sometimes used, such as the "Number of
legs."
You should continue asking questions to further separate the groups of animals until you
feel that only things that are very similar remain in the same group. You may include questions that
produce more than two groups if you wish, but avoid very general questions (e.g., "What is it made
of?") that produce large numbers of categories with few members in each group.
When you finish the classification scheme, name the final groups and trace the scheme from
bottom to top and summarize the characteristics of that group. For example, the name of one group
in a final scheme would be "fish" and a summary of the characteristics of fish might be "Fish are
animals that lay eggs, don't fly, and have gills for breathing." If you cannot "name" a final group,
you probably need to take the separation further.
14
Complete the classification that I started. Demonstrate your understanding of the process by
completing two of the following exercises as assigned by your instructor:
1.
Prepare a different classification scheme for the same animals used in the previous exercise.
You may use some of the same questions, but not all. Your experience with the first scheme
should help you ask more insightful questions in the second.
2.
List twenty or more items that are in your room. Develop a classification scheme. The final
number of groups you have will depend on the items chosen, but there should be at least 5.
3.
You will be given a group of objects to classify. Develop a classification scheme that places
each item in a group by itself.
Report: Prepare a report of your work and hand it in. The report should consist of a) the completed
animal classification scheme, b) the two other classification schemes you were assigned, and c) a
summary statement.
15
Prelab Experiment 2—Separation of a Mixture
Name: _____________
Date: ___________ Lab Day and Time: __________
1. Suppose a student did this lab and got these results: the empty evaporating dish had a mass of 10.83 g
and the original sample plus evaporating dish had a mass of 23.23 g. After the sample was heated in order
to sublime the ammonium chloride, the cooled evaporating dish and contents had a mass of 20.22 g.
Remember to show your work on all parts.
a) How much did the original sample weigh?
b) What mass of ammonium chloride was sublimed?
After water “washing” was done and the remaining solid was dried, the sample plus evaporating dish
weighed 16.93 g.
c) What was the mass of the remaining silicon dioxide?
d) What was the mass of sodium chloride that was removed in the water?
e) What percent ammonium chloride was the original sample?
f) What percent silicon dioxide was the original sample?
g) What percent sodium chloride was the original sample?
2. In your experiment you are told to wait until the evaporating dish is room temperature before
you weigh it. Since mass doesn’t depend on temperature, why do we have to be sure to weigh
objects only when they are at room temperature?
16
Experiment 2—Separation of a Mixture
In this lab you will take advantage of some physical properties of some compounds to separate a
mixture of those compounds. You will also determine the percent of each component that is present.
Some physical methods that can be used to separate mixture components are described below.
Distillation: separating a liquid component from a mixture by vaporizing it and condensing the vapor. To
use this method, the liquid must have a different boiling point than the other mixture
components.
Crystallization: removing a solute from a solution by cooling it to reduce the solute solubility, or by
adding a liquid the solute is not very soluble in. To use this method, the solute to be
separated must have a lower solubility than the other components.
Extraction: removing a substance from a mixture by dissolving it (some mixture components do not
dissolve) To use this method, the solute to be separated must have a higher solubility than the
other components.
Filtration: removing a solid and liquid mixture by passing the liquid through a filter paper while retaining
the solid on the paper.
Sublimation: changing a solid directly to a gas without changing it to a liquid first. To use this method,
the substance to be separated must sublime (many solids do not sublime).
To decide which method to use to separate your mixture, we need to look at the properties of the
substances that might be in the mixture. We have looked up some properties for you (see table 3.1).
Table 2.1 Some Physical Properties of Some White Crystalline Compounds
Solubility at 25 oC
 g compound 


 100 g sol' n 
Sodium chloride (NaCl)
35
Ammonium chloride (NH4Cl) 37
0
Silicon dioxide (SiO2)
Compound
Melting Point (oC)
801
Sublimes at 340
1600
You can see that the solubility of silicon dioxide is very different than that of the other components and
the ammonium chloride is different from other components because it sublimes and they don’t.
17
Separating a mixture based on its properties is really very similar to developing a classification
scheme based on its properties (like in expt. 1)—we just have to be sure to use properties that are
practical in lab, and we keep asking questions until we separate it as much as we need. One
classification scheme we could use is:
sodium chloride
silicon dioxide
ammonium chloride
Does it sublime?
yes
no
sodium chloride
silicon dioxide
ammonium
chloride
yes
sodium chloride
Does it dissolve
well in water?
no
silicon dioxide
We can use the schematic as a guide—we could first separate the ammonium chloride by subliming
it and then take the rest of the mixture and separate the sodium chloride by dissolving it.
Procedure
Record your Unknown Number on your data sheet. Weigh an empty, clean, dry evaporating
dish and place your sample (about 10 g) into a clean, dry evaporating dish. Weigh the dish plus sample.
Be sure to record both masses to as many decimal places as possible.
Removal of Ammonium Chloride
Carry out this procedure in the hood, at least ten inches from the front of the hood.
Place a ring on a ring stand, place a wire gauze on the
ring and put a Bunsen burner under the ring. Adjust the height
of the ring so it is about 5 cm above the burner. Light the
burner and adjust the flame to gentle, but not yellow. Place the
evaporating dish on the wire gauze and heat it by moving the
burner so the flame moves around the bottom of the dish.
As the ammonium chloride sublimes, there should be a
white “smoke” given off. If there is no smoke after several
minutes of heating, skip to the removal of sodium chloride.
Otherwise, when the amount of smoke decreases, stop and stir
the contents of the evaporating dish with a glass stirring rod,
then heat it again. After doing this several times, the smoke
should completely cease to be evolved. It will probably take
about 15 minutes of heating.
After this smoke ceases, turn the burner off and allow
the evaporating dish to cool to room temperature. After you are
sure it is at room temperature, weigh the dish and contents. The
difference between this mass and the original mass is the mass
of ammonium chloride sublimed.
18
Removal of Sodium Chloride
Add about 10 mL of DI water to the evaporating dish and stir it for several minutes.
Note: You may need to use more water depending on the size of your evaporating dish. There
should be enough water so that there is at least 1 cm water covering the sand. Stop and let any
solid settle. Then, into a 250 mL beaker carefully pour off all the liquid that you can while leaving
all of the solid behind (a process called decanting). Use your stirring rod to direct the flow of the
liquid. If you don’t know how to do this, your instructor will demonstrate. Repeat this water
addition, stirring and decanting procedure three times (for a total of four times). Don’t forget to
rinse down the sides of the evaporating dish with each extraction. Insufficient “washing” of the
mixture with water by not stirring enough is a major source of error in this experiment.
If you have sufficient time, put some of the recovered solution into a small beaker and heat
it until it is almost completely dry (but not totally dry, or it will likely crack the beaker!). Is there
any evidence that you have sodium chloride?
If all of the sample dissolved in this step, you do not have any silicon dioxide. Otherwise,
do the next step.
Determining the Mass of Silicon Dioxide
After the last decantation place the evaporating dish back on the wire gauze and heat it
gently by waving the flame under it. This should drive off any water; if you heat it too quickly, it
will “spit” or spatter. Stir it often to prevent spatter and to break up any clumps. After the solid
appears dry throughout, use a hotter flame and heat it until the dish is a dull red color for ten
minutes to evaporate the last of the water. Then turn the burner off, let the evaporating dish cool to
room temperature, and determine its mass.
Determining the Mass of the Sodium Chloride
The grams of sodium chloride will be determined by calculation. Rearrange the following to find it.
Original sample mass = g of ammonium chloride + g of silicon dioxide + g of sodium chloride.
Percent by Mass Composition
The percent of a component in a mixture is determined with
Percent component 
g component
x 100
g original sample
Report:
Turn in observations, and a list of which components were present in your mixture. Turn in all data
tables, clearly labeled calculations and results and calculate the percent of each component that was
present in your mixture.
Questions
1. What components were in your mixture? Use your observations and explain why you think you
had those things and why you think you did not have the other possible components.
2. What would happen if you weighed the evaporating dish + sample while it was still warm?
3. If a student actually had sodium chloride in their mixture, what evidence of sodium chloride
would they see by evaporating the liquid that was decanted?
4. How would the percent of sodium chloride and of silicon dioxide change if you didn’t wash the
sample enough? Explain.
19
Report Sheet: Separation of a Mixture
Name ______________________
Date: ____________________
Lab Day & Time: ____________
Unknown Number: _______
Mass of ED and sample, g
____________
Mass of evaporating dish (ED), g
____________
Mass of sample, g
____________
Mass of ED and sample after 1st heating, g
(If no “white smoke” was observed skip this step)
____________
Mass of sample after 1st heating, if appropriate, g
____________
Mass of NH4Cl in sample, g
____________
Mass of ED and dry SiO2, g (after2nd heating)
____________
Mass of SiO2 in sample, g
____________
Mass of NaCl in sample, g
____________
% NH4Cl in sample
% SiO2 in sample
% NaCl in sample
____________
____________
____________
Observations:
Calculations:
Summary statement
20
Prelab for Experiment 3: Metric System and Volume of a Drop of Water
Name:
______
Lab Day & Time:_____________
In this course you will often have to prepare tables in which you will record your
experimental results, both measured and calculated values. So it is important that you learn how to
do so.
Note: There is an explanation on how to make tables on page 8, hints given below, and an example
taking you through the process on page 31. If you have trouble preparing the tables, see your
instructor before lab. They need to be done before you come to lab.
Hints for making a data table: One approach to making a table is to first carefully read through the procedure and
questions and make a list of all the measurements and calculations you will be making. If you are asked to repeat the
measurements, you need to make spaces for all the trials (better leave room for one extra, just in case). If you have to
repeat the measurements, you will also need to calculate an average value, so make a place for it at the end of the table.
Typically you should put data that is similar in one table, or data that is going to be used together into one table.
For example, you are going to use the some data from part I to answer a question about how well you have learned to
estimate mass in grams, so it makes sense to put that data together in the same table.
It is traditional to make the same kind of measurement either all rows or all columns; to decide which to make
rows, one consideration is being sure the data fits on the page well. For example, if I was going to measure the mass of
12 samples before and after I heated each sample, I’d probably make the masses the columns and the sample names
rows because I can usually make more rows on a page than I can columns (assuming the normal way to hold a paper).
1. On a separate paper, prepare appropriate tables to hold your estimates and the measured
values for all items in parts I-III of the lab procedure. Be sure to title your tables and to include
units. Since you will be answering questions based on these data (see report section of write-up),
leave appropriate spaces for your answers.
21
Experiment 3: Metric System and Volume of a Drop of Water
Introduction:
The two most common measurements you will make in lab are mass in grams and volume in
milliliters. Having some idea of the size of meters, centimeters, and millimeters will also help you
in lab and in reading your textbook. In this experiment, you will be trying to build some intuition
about the size of metric units and you will practice making measurements.
You will first make some estimates and then check to see how good these estimates are.
Some rough guides to help with estimates are: the width of one of your fingers is about 1 cm, the
distance from your nose to the tip of a finger on your outstretched arm is about 1 m, a nickel has a
mass of about 5 g and a sugar cube has a volume of 1 mL.
When making measurements in the lab, the number of decimal places you record tells the
uncertainty in the number. The rule is to record all the numbers that you are sure about and estimate
one more decimal place; in practice, this means you should record one more decimal place than is
directly indicated by the closest calibration marks on the measuring instrument. Always record the
largest number of digits that you can, given the measuring instrument you are using. For example,
if a graduated cylinder is marked every 0.1 mL, i.e., one decimal place, then you should record the
volume to two decimal places, i.e., to nearest 0.01 mL.
If (as with our balances) there is a digital display, the uncertainty is in the last digit displayed
and you should record all of the displayed digits. In this lab you are first estimating and then
measuring, and the number of digits in your estimate may be less than the number of digits in your
measurement.
In Part IV of the experiment you will explore two concepts: how to measure really small
things, and how to determine if data are reliable. The really small thing you will measure is the
volume of a drop of water. Using the class data for the volume of a drop of water will give you the
opportunity to investigate data reliability.
Lab Procedure: For Parts I-III, make the estimate for an item and then immediately do the
measurement. Then go to the next item. By finding the measured value for an item before you do
the next estimate, you’ll probably get better at the estimates as you go. Your lab instructor will
show you how to use the balances and will briefly discuss volume readings. You may do the parts
of the lab in various orders, but must do Part IIIA before doing Part IV.
I. Mass
A. Estimate the mass in grams of one of the following, then weigh it to see how close you estimate
is.
a small test tube
a 50 mL beaker
a 250 mL beaker
a scoopula
a stirring rod
a large test tube
an evaporating dish
a medicine dropper
B. Repeat step A for three more items in the list.
II. Length measurements
Use a ruler and decide which of your fingers is closest to 1 cm wide. Report which finger it is
on your report sheet. (Since this is only reported one time and not compared to other data, it is
not appropriate to put this information in a table. A sentence will be fine.)
22
A. Measurements in Centimeters
1. Estimate the length of a small test tube in centimeters, then measure it to see how close you
were.
the height of the test tube rack
the height of a 250 mL beaker
the length of the stirring rod
the length of the scoopula
2. Repeat step 1 for two more items in the table in 1.
B. Measurements in Meters
Estimate the following in meters, then measure them to see how close you were:
width of your hood in meters
depth of your hood in meters
III. Volume Measurements
A. Making Sure Your Read Volumes Correctly
On each of the center benches, there are two graduated cylinders that have water in them. For
one large and one small graduated cylinder, make a reading of the water volume in milliliters
and record it. Be sure to record it to the appropriate number of decimal places! Now, check
with your lab instructor to see what the accepted value is. Record this as the stockroom value.
For the smaller cylinder, you should be within 0.03 mL of the stockroom reading; for the larger
cylinder, you should be within 0.3 mL of the stockroom reading. If you are “off” by more than
these amounts, consult with your instructor.
B. Making and Estimating Volume Readings
1. Estimate the volume of a small test tube in milliliters, then measure its volume by filling it with
water and pouring the water into a graduated cylinder. Fill each item to the brim.
2. Repeat step 1. for the other items below:
Large test tube
The plastic bottle that is in your hood.
The vial that is in your hood.
IV. Measuring the Volume of a Drop of Water
1. Determine the volume of a drop of water by counting a large number of drops delivered from a
medicine dropper into a small graduated cylinder that is partially filled with deionized water.
Record the volume reading, to the nearest 0.01 mL, before you start and after you finish. Make
this measurement until your results seem reproducible, at least in triplicate. From this data also
calculate how many drops are in 1 mL and compare your data with the results of the class.
23
This week’s lab had two main sections—estimating and making measurements (parts I-III) and
then measuring the volume of a drop of water and evaluating data (part IV). Because the results
of parts I-III are not directly related to part IV, you should write two mini-reports instead of one
large report.
Report for Parts I-III
Of course, you need to include your tables. Before you answer Question 1, think about how
your reader will use your answer. Most likely they will look at your data and see if they agree
with your answer, so it makes sense to put the answer to this question close to the data. Answer
the question about mass close to the mass data, the question about volume close to the volume
data, etc. If a question refers to several sets of data, it should be after all the sets it uses.
Questions:
1. For length, mass and volume measurements/estimates: Discuss how reliable your estimates
of were. Were there any systematic trends in your estimates; that is, did you generally
overestimate or underestimate mass? volume? length? everything? Write out your findings
concerning this.
2. Name common objects (besides those given in the lab introduction) that have the following
sizes. Report this information in a table.
about a centimeter in length
about a meter long
about ten grams in mass
about 10 mL in volume
about 250 mL in volume
3. Use your data for the length and volume of a small test tube and discuss what percent full the
test tube should be for it to contain about one milliliter of liquid. (It also will be helpful in later
labs to know this number, so write it down in your lab manual somewhere for future reference.)
Report for part IV
Tabulate the class results for the number of drops/mL water. Calculate and report the class mean
(average).
1. Do any of the results look unreliable enough that they should be ignored? If so, list these
and discuss how far off they are. Recalculate the mean leaving out these "suspect" results.
Did the mean change enough to support your decision?
2. How close were your results to the class mean found after discarding any results in step
3. Do you think your result was reliable? Why?
4. To how many significant digits should your average and the class mean be reported?
Explain how you should decide. What keeps the class data from being more reliable?
Summary:
a. Explain how you should make sure data you collect is reliable, how do you check its reliability
and show how reliable it is (or isn’t!). Answer for a general case.
b. If you added just one drop of water to a graduated cylinder and tried to measure the volume, you
would not be able to do so with any reliability at all. The procedure in this experiment let you
get a fairly reproducible volume for a drop. Generalize what was done to tell how to measure
other small things—for example, describe a procedure to use to find the mass of a grain of rice.
b. Summary Statement
24
Hints for making a data table:
Many times this semester you will be asked to prepare data tables. Most of you have never had to do this before so
here is one approach. First, carefully read through the procedure and make a list of all the measurements and
calculations you will be making. Don’t forget average values where appropriate. For example, in part IV, you are to
determine the volume of a drop of water. From the procedure we see that we will measure the following: number of
drops, an initial volume, and a final volume. We will also need to calculate the following: net volume and volume per
drop. We are asked to repeat the measurements, at least twice more, so we will make spaces for at least 3 trials (better
leave room for 4). We will also need to calculate an average value for volume per drop and the # drops per 1mL.
One possible arrangement is shown below.
Table 1 Determination of the Volume of a Drop of Water
Final Volume, mL
Initial Volume, mL
Net Volume, mL
# of Drops
Volume/drop, mL
Average Volume/drop, mL
# Drops/1mL
Trial 1
______
______
______
______
______
Trial 2
______
______
______
______
______
Trial 3
______
______
______
______
______
Trail 4
______
______
______
______
______
_______
_______
Be sure to title your table and to include units. Note that the order in which the measurements are listed is not the same
as the order in which they are made. It is easier to subtract the initial volume from the final volume if the final volume
is listed first. Certainly, the # of drops could have as easily been listed first. This placement is just a matter of personal
choice. For tables with more than three or four items and multiply trials it is usually better to have the items listed
vertically as shown here.
25
Name__________________________
Date_________________
Partner______________________
Lab Day & Time___________________
Table ____: Determination of the Volume of a Drop of Water
Trial 1
Trial 2
Trial 3
Trail 4
Final Volume, mL
_________
_________
_________
_________
Initial Volume, mL
_________
_________
_________
_________
Net Volume, mL
_________
_________
_________
_________
# of Drops
_________
_________
_________
_________
Volume/drop, mL
_________
_________
_________
_________
Average Volume/drop, mL
_________
# Drops/1mL
_________
26
Name__________________________
Date_________________
Partner______________________
Lab Day & Time___________________
Table ___: Class Data for # Drops per mL
Student Name(s)
drops of
water/mL
27
Prelab for Experiment 4: Density and Reliability of Data
Prepare a table to hold the data you collect and values you calculate in this experiment. Be
prepared to turn in this table at the beginning of lab. Don’t forget to bring your calculator to
lab!
28
Experiment 4: Density and Reliability of Data
Introduction:
This week you will determine the density of a salt water solution. Density is the ratio of
mass
mass to volume for a sample. D 
volume
Recall that when making measurements in the lab, the number of decimal places you record
tells the uncertainty in the number. The rule is to record all the numbers that you are sure about and
estimate one more decimal place; in practice, this means you should record one more decimal place
 by the closest calibration marks on the measuring instrument. Always
than is directly indicated
record the largest number of digits that you can, given the measuring instrument you are using. For
example, if a graduated cylinder is marked every 0.1 mL, i.e., one decimal place, you should record
the volume to two decimal places, i.e., to the nearest 0.01 mL.
If (as with our balances), there is a digital display, the uncertainty is in the last digit displayed
and you should record all of the displayed digits.
Procedure:
Your instructor will assign a salt water solution whose density you are to determine. Be sure to pay
attention to significant digits when you do calculations.
A. Rinse a 100 mL graduated cylinder with a small volume of the salt water solution and discard
the rinse.
B. Put ~10 mL of the solution into the graduated cylinder and make a good volume reading.
Weigh the cylinder with this amount of solution in it. (The volume you start with should be
enough to allow you to make a good volume reading, but not much more since this would
just waste solution.) Remove the cylinder from the balance. NEVER pour anything into a
container while it is on a balance.
C. Now add 10 -15 mL of solution and make another volume reading and another weighing of
the graduated cylinder plus solution.
D. Repeat step and C at least twice so you have at least three good sets of data. You may need a
fourth trial if you have difficulty with one of your trials.
E. For each trial, calculate the net volume and the net mass (i.e., the volume and mass of each
added portion) and the density in g/mL.
F. Calculate the average density. (Before proceeding to part G, please check the reliability of
your trials as discussed in part 1 below.)
G. Record the percent salt and the average density of your solution on the class data sheet that
your instructor posts. Make sure you have a copy of all of the class data.
Report:
1. Include a table that includes your data and the results of your calculation of density for each trial.
Report an average density (g/mL) for your solution. Do any of your results seem unreliable
enough to be discarded? If so, recalculate the average density without this result. Does the new
average support your decision? Be sure to record the more reliable average on the class data
sheet and to use it in the rest of your report. Show a sample calculation of all information
leading to the density.
2. Include a table of the class data for the average densities (g/mL) of the salt solutions, grouping
solutions of the same % NaCl together.
29
3. Graph all the density data tabulated in 2, plotting density on the y-axis and %NaCl on the x-axis.
Draw a best straight line through the points using a ruler. (Do not connect the dots.) Note: Hints
for preparing graphs can be found on page 8-9 of this manual. Ask your instructor whether you
are to do the graph by hand or on the computer.
4. From your graph, decide if you think that the reliability of any point(s) is questionable.
5. Look up and tabulate the accepted densities of salt solutions in the “CRC Handbook of
Chemistry and Physics”, being sure to reference the particular handbook you used. Use these
CRC values to check the accuracy of your result. Discuss your accuracy.
6. Plot these accepted CRC values on the graph prepared above, but be sure to use a different
symbol than the one used for experimental data. DO NOT draw a line through these points. Use
the CRC data as compared to the experimental data to decide if any questionable data point
should be discarded. Explain why you did or did not discard any data points.
Questions
1. Why were you told to use 10-15 mL portions of solution for your trials, i.e. why at least, 10 but
not more that 15 mL?
2. Why were you told to make the mass and volume measurements (for the density) several times?
3. In this experiment, there were three steps in evaluating whether any data should be discarded four if another group had the same % solution you had. What were they? Summarize, in a
generic statement, those that would apply to any experiment.
Summary Statement
30
Name__________________________
Lab Day & Time___________________
Class Data for Density of Salt Water Solutions
Student Name(s)
%NaCl
31
Density (g/mL)
of NaCl solution
Experiment 5: Solubility and Conductivity
Purpose.
1.
To prepare a water solution of a compound and determine the compound's solubility.
2.
To determine electrical conductivity of solutions and use that information to classify the
solutions according to electrical conductivity.
3.
To determine whether the compounds are soluble ionic, molecular polar, or insoluble.
Part I: Preparing a solution
Water is an excellent solvent. Many materials dissolve in water, but some are more soluble
than others. In this activity you will determine the approximate solubility of several compounds.
WARNING! There are several precautions that you should follow to get useful results:
1.
Do not handle any of the solids with your hands, get them on your skin, breath dust
from them, or get them into your mouth. Some irritate skin, and some are toxic. All
chemicals should be handled with care.
2.
Clean all glassware thoroughly. Wash with a brush and tap water. Rinse three times with tap
water and then with a small amount of distilled water. If your solution is contaminated, the
results for Part II of the experiment will be ruined for everyone!
3.
Make all measurements as precisely as possible.
You will be given a 1.0 g sample of one of the compounds listed in Table 1. You will see
how much water it takes to dissolve it.
Part I Procedure:
1.
Carefully measure 1 mL of distilled water and add it to the sample in the vial.
2.
Replace the lid and shake vigorously to dissolve the solid. Be patient here; this may take
several minutes. If all of the solid dissolves, skip to step 6 of the procedure. If it does not all
dissolve, continue with the next step.
3.
Measure out an additional 2 mL of distilled water and add it to the solution.
4.
Replace the lid and shake to see if 1.0 g of your solid will dissolve in 3 mL of water. If it
does, skip to step 6.
5.
If the solid did not dissolve in 3 mL of water, measure out an additional 2 mL of distilled
water and add it to the solution. Replace the lid and shake to see if 1.0 g of your solid will
dissolve in 5 mL of water. If it does, skip to step 6.
5.
If the solid did not dissolve in 5 mL of water, continue to add 5 mL portions (aliquots) of
distilled water to the solution, to try to dissolve your solid. The vials will hold 20 ml of
solution. As soon as your solid dissolves skip to step 6. If your solid has not dissolved upon
the addition of 20 mL of water, carefully pour the solution into a clean Erlenmeyer flask.
Add 10 mL of water to the vial, rinse the vial, add the rinse water to the flask, stopper and
shake. Continue adding measured amounts of water until all of your solid dissolves. Record
the volume of water required to dissolve your sample in the appropriate place on the
chalkboard and on your data sheet.
6.
After your solid is dissolved, pour your solution into a clean 100 mL graduated cylinder.
Rinse your vial with small portions of distilled water (be sure to pour the rinse waters into
the cylinder), and add additional water so that the total volume of solution is 100.0 mL.
Pour the solution into a clean Erlenmeyer flask, mix, label the solution, and place it on the
supply table for use during Part II of the experiment. Clean your vial and return it to your
instructor. Clean your cylinder so someone else can use it.
32
Calculations: Copy from the chalkboard the volume data for Table 1. Each milliliter of water
weighs one gram, and each sample of solute weighed one gram. From this information calculate the
solubility of each compound in grams of solute per 100 grams of solution and fill in Table 1.
Convert the solubility you calculated for your compound from parts of solute in 100 parts of
solution (pph or %) to parts of solute in one million parts of solution (ppm). Be sure to include a
sample of each calculation in your laboratory report. What must you do to any quantity
expressed as a percent in order to express the quantity in parts per million? What must you do to a
quantity expressed in ppm to convert it to a percent?
Part II: Electrical Conductivity
Contrary to popular opinion, pure water is a very poor conductor of electricity. However,
most water that we use is not pure. It has a variety of things dissolved in it to form dilute solutions
similar to the one you just made. Many of these solutions are excellent conductors of electricity.
Some are poor conductors, and a few do not seem to conduct electricity at all. In this part of the
experiment, you will determine how well the solutions prepared in lab today conduct electricity.
We have provided a few solutions in addition to the ones you made. In most cases, these
solutions were prepared for you because the pure substance used to make the solution is dangerous
to work with.
Part II Procedure: You will use a conductivity tester consisting of a box with a row of 10 tiny lights
on it, an electrical cord to plug into a 110 v electrical outlet, and another electrical cord with two
metal wires attached at the end. When the two metal wires (usually referred to as the probe) are
connected by something that connects electricity, the lights on the box come on. The better the
conductor, the more lights will come on.
Do the following for each solution.
1.
Add solution to a clean, small test tube until it is about 1/3 full.
2.
Place the probe in the test tube.
3.
Count the number of lights that come on.
4.
Record your data in the second column of Table 2.
5.
Rinse the probe with distilled water between each use.
Classifying the Solutions: Solutions that conduct electricity are called electrolytes. If they are good
conductors, they are called strong electrolytes and if they are poor conductors, they are called
weak electrolytes. Solutions that do not conduct electricity at all are called nonelectrolytes.
The apparatus used in this experiment uses LEDs (lights) to indicate how "good" a
conductor a solution is. The more lights that light up the better the conductor, i.e. the stronger the
electrolyte. Nonelectrolytes light no lights.
Use these terms as categories and construct a scheme to classify the solutions that you
tested. If your data allow, add other categories such as "very strong"," moderately strong," "very
weak," and "moderately weak" to your scheme. Use scratch paper to draft a classification scheme.
When you are satisfied with it, attach a neat copy to your Report Sheet
33
Part III: Determining compound type, i.e. is it soluble ionic, molecular polar, or insoluble?
We use both solubility and conductivity data in determining compound type. Water is the solvent
we are using, and we know that molecular polar and some ionic compounds will dissolve in water.
If a compound is insoluble in water, no further conclusion as to compound type can be drawn. How
can we distinguish between ionic and polar? Previously you were told that pure water is a very
poor conductor of electricity. This is because there must be ions in a solution in order for it to
conduct electricity. We know that ionic compounds dissociate into their component ions when
dissolved in water, but polar compounds do not. So now we know how to distinguish soluble ionic
from molecular polar compounds using conductivity.
Before leaving the lab, retrieve your Erlenmeyer flask, discard any unused solution in the
sink, and return the cleaned flask to your drawer. Be sure to clean the stopper you used
before you return it to the stopper drawer.
34
Report Sheet: Solubility and Conductivity
Name________________________
Date_______
Partner_____________________
Lab Day & Time__________
Name of compound you used __________________
Table 1
Compound
mL of Water Required
to dissolve 1.0 g
Sucrose
NaNO3
Na2SO4
NaCl
Citric Acid
NH4C2H3O2
NH4Cl
CaCl2
Ca(NO3)2
Urea
Al(C2H3O2)3
Ethanol
KCl
KNO3
Hexane
Naphthalene
35
Solubility, w/w %
(pph)
Report Sheet: Solubility & Conductivity, cont.
Name________________________
Table 2
Compound
Number of
lights
Electrolyte Category
Sucrose
NaNO3
Na2SO4
NaCl
Citric acid
NH4C2H3O2
NH4Cl
CaCl2
Ca(NO3)2
Urea
Al(C2H3O2)3
Ethanol
KCl
KNO3
Hexane
Naphthalene
HCl
Glucose
Ethylene glycol
HC2H3O2
NH3
36
Soluble Ionic,
Molecular Polar, or
Insoluble
Report Sheet: Solubility & Conductivity, cont.
Name________________________
Name of compound you used __________________
Formula _________________
Solubility of compound you used expressed as a w/w % (or pph)____________
Solubility of compound you used expressed as ppm____________
Calculations, including sample calculations for solubility:
Classification key (on back)
Summary Statement:
37
Prelab for Experiment 6: Making of Zinc Chloride
Name___________________________
Date_________Lab Day & Time_______________
1. Make a data sheet to hold all the data for the whole lab--one that will be easy to read and will
neatly pair the data needed for the final graph. Assume that there will be a result from each student,
and assume that at least one student will make a calculation or weighing error so that you will need
enough data so you can make a decision to reject or recalculate his data if necessary. I suggest you
include: starting g of Zn, g of Zn recovered, g of Zn reacted, g of zinc chloride produced, g chloride
in zinc chloride, moles Zn reacted, moles of chloride in product, and the mole ratio of chloride to
zinc in the product.
2. Also make a small table to hold your own data for this experiment that will contain all of your
data for the exercise. You should be able to just copy numbers from this table into the class data
sheet, but note that there will be data in your data table that won’t appear in the class data sheet.
[Be sure to do 1 and 2 on separate paper so they can be given back to you to use during lab]
3. A 0.950 g piece of tin was added to a solution of HCl. The Sn completely reacted and 2.092 g
of product, a tin chloride, were recovered. SHOW YOUR WORK! (Watch your sig. figs.)
a) How many grams of tin were in the product?
b) How many moles of tin were in the product?
c) How many grams of chloride were in the product?
d) How many moles of chloride were in the product?
e) What is the experimental chloride to tin mole ratio?
f) What is the formula for the particular tin chloride that was formed?
4. In your experiment you are told to wait until the evaporating dish is room temperature before
you weigh it. Since mass doesn’t depend on temperature, why do we have to be sure to weigh
objects only when they are at room temperature?
38
Experiment 6: Making Zinc Chloride
Introduction:
This experiment has two functions. You will use the data you and your colleagues collect to
determine the formula for zinc chloride and to explore the concept of limiting reactant.
You will be assigned a piece of zinc whose mass is within a specific range. You will then
add exactly 10.0 mL of hydrochloric acid to it and allow it to react for a week, which is enough time
for completion.
You can find how much zinc reacted since either: a) it reacted entirely, or b) an amount
reacted that is equal to the initial weight minus the weight left at the end.
The weight of the zinc chloride product formed can be found by evaporating the product
solution to dryness. Both any excess hydrochloric acid and water will evaporate readily. The zinc
chloride product has a much higher boiling point so it will not evaporate if the mixture is not heated
too much.
You can then find the mass of chloride in the product by subtracting the weight of zinc that
reacted from the weight of the total product.
g chloride = g zinc chloride produced – g Zinc reacted
The empirical formula for zinc chloride can be found by finding the moles of zinc and of chloride in
the product, and finding the mole ratio as in the prelab.
You will prepare a graph of mass of zinc chloride formed versus initial mass of zinc.
Procedure:
WEEK 1: Weigh the assigned zinc sample and place it into a clean Erlenmeyer flask.
Make sure the zinc lays flat on the bottom of the flask. Then measure exactly 10.00 mL of 4 M
HCl and add it to the flask. Set the flask in the back of your hood for the duration of the lab.
Before you leave lab, place the flask in your drawer. Remember that you will need a clean, dry
evaporating dish next week, so wash one this week. Each student will need one; extras are in a
drawer in the center bench. Record the weight of zinc you used below.
Mass of Zn = ____________g
Record your observations for this week’s part of the experiment here.
WEEK 2: Weigh the clean, dry evaporating dish you prepared last week. Pour the liquid
contents of the flask into the evaporating dish. Use a water wash bottle to use 2-3 mL of water to
rinse the sides of the beaker, and add this liquid to the evaporating dish. Repeat the washes twice in
order to transfer all of the liquid from the reaction mixture to the evaporating dish.
If any zinc remained unreacted, use forceps to remove the unreacted Zn. Rinse its surface
over the evaporating dish with a few drops of water. Thoroughly dry the piece of Zn, and record its
mass.
THE EVAPORATION THAT FOLLOWS MUST BE DONE IN THE HOOD. During
the heating keep the sash pulled down to about a height of 20 cm. Use a medium-sized burner flame
to heat the contents to evaporate the water and any unreacted HCl. The evaporating dish should be
supported with a ring-wire gauze combination on a ring stand. Heat gently to avoid splattering
39
liquid or over-heating the zinc chloride remaining. Suggestion: Zinc chloride melts at a relatively
low temperature and will go off as a thick, white vapor if heated to excess. As you evaporate the
solution, you will first see a white crust forming on top and around the edge of the solution. As
water continues to evaporate, all of the material should become a crusty white solid that will begin
to melt around the edges almost immediately. Stop heating when the zinc chloride begins to melt or
you will lose some of it.
Allow the evaporating dish to cool to room temperature, then weigh the dish plus product.
Calculate the results as suggested in the introduction and then enter your data in the data sheet
selected for the class data. Do not leave the lab until all the data has been entered into the sheet and
you have a copy for your own use.
Report:
Include in your report your observations for the experiment, the two data tables, and your
calculations laid out clearly with units that show how you got your results.
Then you should make a good graph of the class results, plotting weight of zinc chloride
formed vertically versus weight of zinc added to the reaction mixture (the initial weight of zinc).
Your instructor will tell you whether you may use the computer or if you need to do this graph on a
regular sheet of graph paper. There are instructions for how to make a graph on pages 8-9 of this
manual. This graph should look like it consists of two straight line segments. You should draw
these two straight lines (do not connect-the-dots).
(1) Explain why the graph shape says that for awhile as more zinc is added, more zinc chloride
was made, but then no matter how much zinc was added the same amount of zinc chloride was
formed.
(2) Discuss how the graph shape tells us what the limiting reagent is for each reaction? What is
special about the point where the two line segments intersect?
(3) What was the limiting reagent for your reaction?
Report the average experimental mole ratio for the class data. Be sure to discuss any class
results you feel are “off”, and to calculate the average class ratio both with and without the result(s)
you feel are off. Were you justified in discarding those data you felt were off?
What should the empirical formula for zinc chloride be based on class results? On your results?
(These two formulae may not be the same.) If your formula is not the same as the theoretical one,
discuss what experimental errors could have occurred to cause your specific deviation.
Summary Statement.
40
Experiment 7: Types of Chemical Reactions
Introduction:
This is an exercise in careful observation of materials and the changes they undergo. You
should do three things in this experiment: carefully describing the starting materials and the changes
they undergo, decide what type of chemical change has occurred: precipitation, acid-base reaction,
redox reaction. If the reaction is a redox reaction, you should also decide what got oxidized and
what got reduced. Your third goal is to write the chemical equation for the reaction.
In truth, it is not always easy to decide what type of change has occurred, but the starting
materials give you a good hint as well as examining the typical reactions of acids.
Precipitation Reactions. Of course, a precipitate forms during a precipitation reaction. You
should expect in this type of reaction that two solutions are combined and a solid forms. The solid
may actually fall to the bottom of the test tube or it may just make the mixture look cloudy. You
should predict the identity of the solid based on the solubility rules.
Acid-Base Reactions. These are harder to figure out unless you are told that the starting materials
are an acid and base. During an acid-base reaction, an H+ is transferred from the acid to the base,
but there is often little to observe—usually the reactants and products are both colorless solutions.
This type of reaction produces heat, but just mixing a concentrated acid solution with water also
tends to produce heat, so heat production is not a good way to tell if an acid-base reaction has
occurred.
It is easier to tell if an acid-base reaction occurs by using pH paper. If a small amount of a
solution is placed on pH paper, the color will indicate the pH of the solution. We have not studied
pH yet, but it is enough to know that the pH gets lower as a solution becomes more acidic and
higher as it gets more basic.
As an example, suppose we start with a solution of NaOH and test its pH and find the pH is
10. Now, suppose we add a few drops of HCl and mix the liquid well and then retest the pH and
find it is 8.0. Since the pH is lower, some of the base (NaOH) is gone—it reacted with the added
HCl. Since the name of HCl(aq) is hydrochloric acid and since NaOH is a base, it is reasonable to
predict that a H+ was transferred during this reaction. The reaction equation is then
HCl (aq) + NaOH (aq)  H2O (l) + NaCl (aq).
Another type of acid-base reaction can be recognized simply from the reactants. An acid
will react with a compound containing either carbonate ion or hydrogen carbonate (bicarbonate) ion
to form gaseous carbon dioxide, water, and another compound. For example, suppose colorless
hydrochloric acid solution and colorless sodium carbonate solution are mixed. The product is a
colorless solution and colorless gas bubbles. Knowing that an acid and a compound with carbonate
form carbon dioxide and water, reasonable products are CO2 and H2O. So far, what we have is
HCl (aq) + Na2CO3 (aq)  H2O (l) + CO2 (g) + ??
We must also account for the Na+ and Cl-, so NaCl (aq) seems reasonable. Upon adding NaCl as
the other product and balancing the equation, we get
2 HCl (aq) + Na2CO3 (aq)  H2O (l) + CO2 (g) + 2 NaCl
41
Redox Reactions. The following are some common types of redox reactions.
(a) If two elements combine to make a compound. For example, if shiny, silver-colored sodium
and yellow chlorine gas react and form a white solid. Because we started with two elements and
ended up with something that has different properties than either element, it is reasonable to assume
that the elements reacted and formed a compound. Because one of the elements was a metal and
one was a nonmetal, the product should be ionic; we would expect a sodium ion and a chloride ion
in the product. Thus, the reaction equation should be the following:
2 Na (s) + Cl2 (g)  2 NaCl (s)
(b) If element one (uncharged) reacts with a compound to form a different element
(uncharged) and a compound that contains the cation of element one. If the uncharged reactant
element is a metal, then an ion of it typically replaces the first element written in the formula of the
compound, and then the first element written in the reactant compound forms an uncharged element
as a product.
For example, suppose a piece of reddish-gold copper metal is placed into a colorless silver
nitrate solution. After thirty minutes, a gray solid can be observed on top of copper-colored metal
and the liquid has a blue tint and is still clear. The gray color of the solid on top of copper-colored
solid indicates a new, different solid formed on top of some of the copper that was originally
present (which is leftover copper).The blue color indicates that there is something else in the liquid
than what was originally there; the fact that the liquid is still clear means the new stuff is dissolved
in solution (since solutions are clear). We have Cu and AgNO3 as reactants, and it would be good
to realize that since the AgNO3 is ionic, we really have Ag+ and NO3- ions at the beginning, along
with the Cu (s). But what are possible products? It is helpful to know that when copper(II) ions are
dissolved in water, a blue solution forms. Let us now regard this as a puzzle—if Cu2+ formed, they
surely formed when some of the Cu metal lost electrons. Something else then had to gain those
electrons (electrons don’t just float around in solution). We must ask ourselves, “Which is a better
candidate to gain electrons, Ag+ and NO3- ions? Since a gray solid is another product, it seems
reasonable to guess that the Ag+ might have gained electrons to form Ag. We now have something
that lost electrons and something that gained electrons, so we don’t need the nitrate ion to change.
We could actually write two different equations, depending on whether we write the “molecular”
equation or the net ionic equation:
molecular: Cu(s) + 2 AgNO3(aq)  Cu(NO3)2 (aq) + 2 Ag (s) or
net ionic:
Cu(s) + 2 Ag+ (aq)  Cu2+ + 2 Ag (s)
Another example of this type of redox reaction is when The H+ of an acid reacts with an
uncharged metal atom to form the metal ion and hydrogen gas (H2). For example, if a piece of
solid, silvery zinc is placed in colorless hydrochloric acid solution, colorless gas bubbles form, the
solution remains colorless, and the solid gradually gets smaller. So far, we can write
HCl(aq) + Zn (s)  H2 (g) + ? . There are two facts that help us write the other products—the
H+ of the acid is gaining electrons from somewhere, and the Zn(s) is getting smaller. This would
happen if the Zn atoms gave electrons to the H+ and formed aqueous Zn2+. If both Zn2+ and Clare in the solution, then the other product would be ZnCl2 (aq) in the molecular equation.
HCl(aq) + Zn (s)  H2 (g) + ZnCl2 (aq) (“molecular”) or
H+ (aq) + Zn (s)  H2 (g) + Zn2+ (aq)
(net ionic)
42
(c) If a compound decomposes to form an element or to form an element and another
compound.
For example, if we heat black mercury(II) oxide, it forms a silver liquid and a colorless gas. Since
one substance forms more than one substance, this is a decomposition. If we remember that
mercury is a silver liquid, it seems reasonable to suppose that HgO  Hg(l) is part of the reaction
equation. Since a colorless gas forms, then 2 HgO (s)  Hg (l) + O2 (g) is the reaction equation.
Procedure:
You should note that only part of the sample might undergo change and some of the starting
material may be leftover—if so, decide if a change has occurred based on whether any of the
sample changed.
HAZARDS ! !
1. MANY OF THESE PROCESSES GENERATE VAPORS THAT MAY BE TOXIC SO
ALL WORK MUST BE DONE IN YOUR HOOD. KEEP THE SASH CLOSED TO HALF-HEIGHT
WHENEVER POSSIBLE.
2. MANY OF THESE MATERIALS ARE FLAMMABLE, SO USE A FLAME ONLY WHEN
IT IS NEEDED, EXTINGUISHING BURNERS WHENEVER POSSIBLE. Keep a watch glass handy
for extinguishing flames in open containers.
3. THE MOST LIKELY HAZARD IN THIS EXERCISE IS BURNS. HOT TUBES AND
EQUIPMENT WILL NOT LOOK HOT. ALLOW AT LEAST TEN MINUTES AFTER HEATING
BEFORE TOUCHING AN OBJECT. Hot objects that are being cooled should be placed on a wire
gauze on the metal of the steam bath. Hot test tubes should be placed in a small beaker on a wire
gauze.
Individual instructions are provided for each part. Whenever a test tube is heated or hot, use
the wire test tube clamp to hold it. Your instructor will show you how to adjust and light a Bunsen
burner before you begin.
On your report sheets,
(1) write the formula of each starting material above its name. (Both name and formula of the
starting materials are given on the reagent bottles.) If a reactant is heated, it might react with the
oxygen in the air—if so, oxygen is listed at the start, but this does not mean it always reacts.
(2) report your observations for each part including descriptions of starting materials and products,
what you observe happening during the procedure and what it looks like after the procedure. In this
and other labs, always mix solutions together well before doing your final description.
(3) For each part, clearly state what type of reaction occurred. If a redox reaction occurred, tell if
it was a decomposition or not and also tell what got oxidized and what got reduced.
(4) Write the reaction equation; your instructor will tell you whether to write the molecular or net
ionic form of the equation.
43
Report Sheet-Types of Reactions
Name__________________________
Date_________________
Partner______________________
Lab Day & Time___________________
A. Using tongs, hold a piece of magnesium in the blue flame of a Bunsen burner just until it ignites.
Hold the burning magnesium over a wire gauze as a precaution in case the burning magnesium
drops. Observe any change.
Do not stare directly at the bright light (look to one side of it). Looking directly at the light can
cause permanent eye damage.
magnesium and oxygen
Type of reaction:
Reaction equation:
B. Place one pea-sized piece of calcium carbonate in a small test tube and place 1 mL (about 1 cm
deep) of 3 M hydrochloric acid in another small test tube. Mix the two and observe any change.
Hydrochloric acid is corrosive and causes severe burns. If it gets on your skin or in your eyes, you
must wash it off with large amounts of water.
calcium carbonate and hydrochloric acid
Type of reaction:
Reaction equation:
44
Report Sheet-Types of Reactions (cont.)
Name________________________
C. Place one matchhead size piece of calcium in a small, dry test tube. Use your test tube clamp to
hold this test tube. Place about 3 mL of water (about 3 cm deep) in another small test tube. Add the
water to test tube containing the calcium.
calcium +
water
Type of reaction:
Reaction equation:
D. Place around 1 mL of 0.1 M copper(II) sulfate in one small test tube and 1 mL of 3 M sodium
hydroxide in another small test tube. Mix the two solutions and observe any change.
Sodium hydroxide is corrosive and can cause severe burns. If it gets on your skin or in your eyes,
you must wash it off with large amounts of water.
copper(II) sulfate +
sodium hydroxide
Type of reaction:
Reaction equation:
45
Report Sheet-Types of Reactions (cont.)
Name________________________
E . Cut a 10 cm length of copper wire and shape it into a helix by wrapping it around a stirring rod.
The helix should be compressed so that it is only about 8 mm high. Place it into a small test tube.
The copper should fit loosely in the test tube, i.e. it should move about freely. Heat the test tube
vigorously over a flame using a Bunsen burner until the tube begins to glow and then allow it to
cool. After the product has cooled, transfer it from the test tube to an evaporating dish. Examine
the product carefully and use a scoopula to scrape at the outside to see if it is uniform throughout.
Copper + Oxygen
Type of reaction:
Reaction equation:
F. Cut a piece of Mg ribbon about 3 mm long and place into a small test tube. Place 1 mL (about 1
cm deep) of 3 M hydrochloric acid in another small test tube. Mix the two and observe any
change.
Hydrochloric acid is corrosive and causes severe burns. If it gets on your skin or in your eyes, you
must wash it off with large amounts of water.
hydrochloric acid and magnesium
Type of reaction:
Reaction equation:
46
Report Sheet-Types of Reactions (cont.)
Name________________________
G. Place about 1 mL (1 cm height in test tube) of 3 M ammonia in one test tube. Test its pH, as
directed by your instructor. Add a few drops of 3 M hydrochloric acid, mix the liquid well and test
the pH again. Repeat the acid addition and pH test.
hydrochloric acid and ammonia
Type of reaction:
Reaction equation:
Summary statement.
RE: Next week’s lab
Ask your lab instructor whether you need to write molecular and total ionic equations and net ionic
equations next week or if you will just be doing net ionic equations.
47
Prelab for Experiment 8: Writing Chemical Equations
Name ______________________ Lab Day & Time __________ Date ________
On this page write a list of the names and formulae of all the compounds you will be using in this
experiment. Most compounds will be used more than once, but you only need to list them once.
Keep a copy for your use during lab. Bring your text book +/or a list of solubility rules to lab.
48
Experiment 8 - Writing Chemical Equations
In this experiment you will observe aqueous solutions of compounds, mix the solutions,
observe the mixture for any sign of a chemical reaction, and write the chemical equations for the
reaction. Be sure to bring your textbook to lab since you might need to use the solubility rules to
help you decide the formulas for some of the products. Remember that in chemical equations we
use subscripts to indicate the state of a compound, i.e. is it a solid-(s), liquid-(l), gas-(g), or an
aqueous solution-(aq). Since ions must be in solution we do not have to include their state.
If you need to write molecular, total ionic and net ionic equations, read examples 1 and 2. If you
need to write only net ionic equations, read examples 3 and 4
Example 1:
Iron(III) chloride(aq)
clear, yellow sol'n
+
Sodium sulfide(aq)
clear, colorless sol'n
________
black solid
>
After you have recorded your observations, write the “molecular” formulas for the reactants. Then
separate each compound into its appropriate ions. Don’t worry about the number of each ion, only
what the ions are. See below.
FeCl3(aq)
+
________
Na2S(aq)
Fe3+ + Cl- (NOT Cl3- !!!!)
>
+ Na+ + S2-
Now that we know the ions we can write correct molecular formulas for the products.
FeCl3(aq)
+
________
Na2S(aq)
> Fe2S3
+ NaCl
From solubility rules, we know that sodium salts are soluble, and that chlorides, with the exception
of PbCl2, Hg2Cl2, and AgCl, are soluble. So the NaCl must still be dissolved, so we label it (aq).
Therefore, the precipitate must be formed from the Fe3+ and S2-, i.e., it is iron(III) sulfide. Its
formula would be Fe2S3. The solubility rules support this conclusion. We would label this insoluble
solid as (s) . After inserting the states of the products, we also balance the “molecular” equation.
2 FeCl3(aq)
+
________
3 Na2S(aq)
> Fe2S3(s)
black
+ 6 NaCl(aq)
We know that only soluble compounds in aqueous solution will break into ions, i.e., those labeled
solids, liquids, and gases stay the same. So we can complete the right-hand side of the ionic
equation we started above. We also now include the numbers of each ion and compound in the
equation
The balanced ionic equation would be
2 Fe3+ + 6 Cl- + 6 Na+ + 3 S2-
________
>
Fe2S3(s)
black
+
6 Cl- + 6 Na+
We cancel the "bystander" (spectator) ions, i.e., the species that are exactly the same on both sides.
In this case the Na+ and Cl-. This leaves the following net ionic equation.
2Fe3+
+
3S2-
________
> Fe2S3(s)
black
49
Example 2:
Nitric acid(aq)
Clear, colorless soln
HNO3(aq)
+
+ Sodium hydrogen carbonate(aq)
Clear, colorless soln
________
NaHCO3(aq)
________
>
colorless gas
>
We know that when we mix an acid with a carbonate, or bicarbonate, we get carbon dioxide gas and
water. Also, in this case, no precipitate was observed since both nitrates and sodium salts are
soluble. Therefore, the balanced molecular equation is:
HNO3(aq) +
NaHCO3(aq
________
> CO2(g) + H2O(l) + NaNO3(aq)
colorless
and the balanced ionic equation will be:
H+ + NO3- +
Na+ + HCO31-
________
> CO2(g) + H2O(l) + Na+ + NO3-
Cancel "bystander" (spectator) ions as before to get the net ionic equation.
H+ + HCO31-
________
> CO2(g) + H2O(l)
colorless
50
Example 3—If solid formed and you do not need to write molecular or (total) ionic equations
(1) Record your initial observations under the name of each solute.
Write the description of the observed product(s) after the arrow. Do not forget to note the color (or
colorless nature) of the liquid in with any solid. If you cannot tell whether a liquid is colored in
addition to the solid being colored, centrifuge the mixture.
Iron(III) chloride
clear, yellow sol'n
+
________
Sodium sulfide
clear, colorless sol'n
>
black solid and colorless solution
(2) Write the formulas of any ions present in the solutions. If it is an ionic compound and it
dissolved, then its ions are separated.
Fe3+
Na+
Cl- (NOT Cl3- !!!!)
S2-
(3) Now use the ion formulas you wrote and look at the combinations that might result if the cation
and anion “switched partners”. Use the solubility rules to see if these “maybe” compound(s) are
soluble or insoluble.
Fe3+
Cl -
Na+
S2-
In this example, check to see if iron (III) sulfide is soluble or
insoluble, and check to see if sodium chloride is soluble or
insoluble.
We find that the sodium chloride is soluble (it follows the
sodium rule) and that the iron(III) sufide is not soluble (it
follows the sulfide rule).
The insoluble compound is the solid. The ions that would have gone into a soluble compound
didn’t react—they remain unchanged and we leave them out of the net ionic equation. So, if a solid
formed, only the ions that go into the insoluble compound actually changed their surroundings—so
they are all we put in as reactants in the net ionic equation. Since they started out dissolved, we put
(aq) after those ion formulas. In this example, the product is the solid formed from Fe3+ and S2- .
We must remember to make the charges add to zero, to show that the state is solid (s) and to balance
the equation. The net ionic equation is:
2 Fe3+ (aq) + 3 S2- (aq) → Fe2S3 (s)
This explanation is longer than what you would write on the report sheet. On the report you would
write the following:
Iron(III) chloride
clear, yellow sol'n
Fe3+
Cl -
Na+
S2-
+
________
Sodium sulfide
clear, colorless sol'n
2 Fe3+ (aq) + 3 S2- aq) → Fe2S3 (s)
51
>
black solid and colorless solution
Example 4—If a solid didn’t form, but the reaction is between an acid and a base.
To do the lab equations for acid reactions, you need to know that
(i) if we mix an an H+ provider (an acid) with a sample containing carbonate, or bicarbonate,
the products are carbon dioxide gas and water.
(ii) if we mix an acid with a compound that contains hydroxide ion, the OH- takes on an H+
to form water.
(iii) hydrochloric acid is a strong acid, and when place in water it forms H+ and Cl-.
Steps to follow:
(1) Record your initial observations under the name of each solute.
Write the description of the observed product(s) after the arrow. Do not forget to note the color (or
colorless nature) of the liquid.
+ Sodium hydrogen carbonate(aq) ________>
Nitric acid(aq)
Clear, colorless soln
Clear, colorless soln
colorless gas and
clear, colorless liquid
(2) Split the acid into H+ and whatever is left after the acid loses one H+.
Also write the formulas of any ions present in the solutions. If it is an ionic compound and it
dissolved, then its ions are separated.
For this example if we start with HNO3 and remove one H+, then NO3- is left.
The sodium hydrogen carbonate contains sodium ion and hydrogen carbonate ion.
So, the starting solutions actually contain these:
H+
NO3Na+ HCO31Using rule (i) above, we know that the H+ and HCO31- react and form carbon dioxide and water.
The other ions do not react. In the net ionic equation, we write what is going to react and then what
is produced. Mark the states (aq, g, etc.) and balance the equation.
So the net ionic equation for the reaction of nitric acid and sodium hydrogen carbonate is
H+ (aq) +
HCO31- (aq) → CO2 (aq) + H2O (l)
Again, the explanation is long, but on your report sheet, you would write
+ Sodium hydrogen carbonate(aq) ________>
Nitric acid(aq)
Clear, colorless soln
H+
NO3-
H+ (aq) +
Na+
Clear, colorless soln
HCO31-
HCO31- (aq)
→ CO2 (aq) + H2O (l)
52
colorless gas and
clear, colorless liquid
Procedure: Put 10 drops of the first solution in a small test tube; observe the color and clarity. Put
10 drops of the second solution in another small test tube; observe the color and clarity. Add the
contents of one test tube to the other, mix, and observe. A centrifuge will be available for use in
cases where you are unsure if a precipitate was formed, or what color the ppt is. Record your
observations concisely on the data sheet as indicated in the examples. Write balanced molecular,
ionic, and net ionic equations as indicated by your instructor.
Caution: SALTS OF HEAVY METALS SUCH AS LEAD, BARIUM, AND NICKEL ARE
TOXIC AND CAN BE ADSORBED THROUGH THE SKIN. CLEAN UP ANY SPILLS AND
WASH ANY OF THESE FROM YOUR SKIN IMMEDIATELY. DISPOSE OF THESE IN THE
APPROPRIATE WASTE CONTAINER.
HYDROCHLORIC ACID AND SODIUM HYDROXIDE ARE CORROSIVE. EXERCISE
EXTREME CAUTION TO AVOID CONTACT WITH THE EYE.
53
Name_________________________ Date ________
Partner _______________________
1.
copper(II) sulfate + sodium hydroxide
2.
barium chloride + sodium sulfate
3.
sodium phosphate
4.
copper(II) sulfate + barium chloride
+ potassium chloride
54
Lab Day/Time _________
Name_________________________ Date ________
Partner _______________________
5.
potassium carbonate + copper(II) sulfate
6.
hydrochloric acid + sodium hydrogen carbonate
7.
sodium sulfate + lead(II) nitrate
8.
copper(II) sulfate + sodium carbonate
55
Lab Day/Time _________
Name_________________________ Date ________
Partner _______________________
9.
copper(II) sulfate + sodium chloride
10.
potassium carbonate
11.
hydrochloric acid + ammonium chloride
12.
hydrochloric acid + calcium carbonate (s)
+ cobalt(II) chloride
56
Lab Day/Time _________
Name_________________________ Date ________
Partner _______________________
13.
sulfuric acid + sodium hydrogen carbonate
14.
nickel(II) sulfate + sodium hydroxide
15.
sodium carbonate
16.
hydrochloric acid + sodium hydroxide
+ nickel(II) sulfate
57
Lab Day/Time _________
Prelab for Experiment 9: Solution Stoichiometry / Acid-Base Titration
Name:_________________________ Date:________ Lab Day & Time:_____________
1. You will need to prepare a data sheet for this week’s experiment. It will need one set of
spaces to list the name, type, appearance, and nominal concentration of the vinegar you choose and
the molarity of the sodium hydroxide. You also need to provide spaces for up to 4 sets of titration
data – 4 trials. Each of these sets will include a final buret reading, an initial buret reading, the net
volume of titrant used, the volume of vinegar used, moles of NaOH added, moles of acetic acid
present, g of acetic acid, % of acetic acid, and molarity of acetic acid present. You will also need
places for average % and average molarity of acetic acid in the vinegar.
2. You must solve the following set of problems which represent the sequence of
calculations in this experiment. (Refer to Supplements 29 & 31.) Be sure to report your answers to
the correct number of significant figures. SHOW YOUR WORK!
a. If 42.28 mL of 0.1178 M NaOH are used in a titration, how many moles of NaOH were used?
b. If the NaOH above was used to titrate a sample of a solution of sulfuric acid, H2SO4 , how many
moles of sulfuric acid are in the sample?
The reaction is; 2NaOH + H2SO4  Na2SO4 + 2H2O
c: How many grams of H2SO4 are in the sample?
d. If the volume of the sulfuric acid solution titrated in this experiment was 10.00 mL, what is the
% sulfuric acid in this solution? Hint: This is a w/v, i.e. weight/volume, percent.
e. What was the molarity of the original solution of H2SO4 ?
58
Experiment 9: Solution Stoichiometry /Acid-Base Titration
Introduction:
Titrations are the most common way to determine accurately how much of a material is
present in a solution. Since you are determining the amount of material, not the identity of the
material, this is an example of a quantitative analysis. In this process, a buret is used. It is a long
thin graduated cylinder with a stopcock (valve) at the bottom.
Using a buret to make a volume reading
The left sketch below is of a buret on a stand.
The way volumes are read with a buret is different than when a graduated cylinder is used.
The major difference is the volume is “read down” instead of up. To show you what that means,
look at the right sketch above. The volume would be read as 15.74 mL. Note that if a buret has
calibration marks every 0.1 mL , then volumes are read to two decimal places.
A buret is also used to measure the net volume added rather than the volume contained in the buret.
A volume reading is made at the start, the buret is used, and then another volume reading is made.
The difference in the volume readings is the volume added. In some ways volume readings are the
same as with graduated cylinders—you should make the reading at the bottom of the meniscus and
you should have the meniscus at eye level.
The other piece of equipment you will use in this lab is a pipet. Your instructor will help you
with this in prelab.
Titrations
The basic process of a titration depends on using a buret to add a solution of known
concentration to a solution of a sample that is being analyzed. Just enough titrant is added to react
exactly with the material of interest. Acid-base titrations are some of the best known examples of
titrations since many indicators exist. An indicator should change color when just enough base has
been added to react exactly with an unknown acid in the solution—this is what tells us to stop
adding the base. When the indicator changes color, we say the endpoint has been reached.
Being able to calculate an accurate result from titration data first depends on the careful
measurements of the volume of titrant added from the buret and of the volume of solution of the
material of interest at the start. The volume and molarity of the titrant are combined to calculate
how many moles of titrant have been added at the endpoint. If the chemical equation for the
reaction is known, then this allows us to calculate how many moles of the material of interest are
present in the original solution. Then the moles of the material of interest and the volume of its
solution (not the volume of the titrant!) are combined to determine the w/v% and/or molarity of the
material of interest.
59
In this specific exercise, the amount of acetic acid, HC2H3O2 , in a vinegar sample will be
found by titration with a sodium hydroxide solution. The reaction is:
NaOH (aq) + HC2H3O2 (aq)
 H2O (l) + NaC2H3O2 (aq)
HAZARD: SODIUM HYDROXIDE IS EXTREMELY CORROSIVE, THOUGH ITS
ACTION MAY BE DELAYED. IT CAN CAUSE INSTANT EYE DAMAGE. Wear eye
protection at all times and immediately wash off spills with large amounts of water.
Procedure:
Record the type of vinegar (white, cider, etc.), the brand of vinegar, and the percent acidity
from the bottle label.
Your instructor will demonstrate for you how to use a pipet and how to read and use a buret.
You should get a buret, rinse it with distilled water and rinse and fill it with the sodium hydroxide
whose molarity you should record. Don’t fill the buret to any number exactly, but just to some
volume near the top that you will read to the nearest 0.01 mL. Also record on your report sheet
the molarity of the sodium hydroxide from the label on the carboy. Pipet a 5.00 mL sample of
the vinegar of your choice into an Erlenmeyer flask. Use your squirt bottle of distilled water to
rinse down the sides of the flask so that you now have 10 - 15 mL of solution in the flask. The
actual amount of water is unimportant since you know the volume of vinegar added. Add 3-4 drops
of phenolphthalein indicator and slowly add titrant from the buret with vigorous, continuous
swirling. As you get near the endpoint, it will take longer and longer for the pink indicator color to
dissipate. Finally, at the endpoint, after vigorous swirling an extremely faint pink color will persist
for at least 25 seconds. Read the buret to the nearest 0.01 mL. Then titrate a second sample. This
should be repeated until successive triplicate trial net volumes agree with a relative range less than
0.3%.
Relative range  100% * (Highest result  Lowest Result) / Middle Result
Calculations: For each trial you should calculate moles NaOH, moles of acetic acid, grams of
acetic acid, molarity of acetic acid, and % acetic acid in your vinegar. This % should be a w/v, or
weight/volume, percent, i.e. the % acetic acid = (g acetic acid/ mL vinegar) x 100. You should
also calculate an average % acetic acid. Finally, you should also calculate an average molarity of
acetic acid.
Report:
The report is simply a table with all your results in it, including the average % acetic acid
and average molarity of acetic acid in vinegar. If you feel that any of your results are unreliable,
leave them out of your average, but be sure to justify this action. Show a sample calculation for
each step of the calculation.
Your instructor may ask you to do a cost comparison between your brand of vinegar and
another.
You should conclude by stating how good your results seem to be based on evidence you
have and whether this seems to be a reliable method of analysis. You should consider your
duplicate trials and concentration information on the bottle label.
Summary statement.
60
Experiment 10: Nomenclature
Before lab, study the nomenclature rules and learn the list of polyatomic ion names and formulas
assigned in class. This lab will be an opportunity to practice using the nomenclature rules so that
you will do well on the exam! The diagram below may help you with memorizing the number of
oxygen atoms in some polyatomic ions.
Charges of most polyatomic ions with oxygen: To figure out the charge instead of memorizing it,
first look on the perioidic table and find symbol of the non-oxygen atom in the polyatomic ion.
Take the last digit of its column number and subtract (2) (number of oxygen). This will be the
charge on the polyatomic ion.
REMEMBER THE “ELBOW”
B
C
N
O
F
Si
P
S
Cl
As
Se
Br
Te
I
At
“Elbow” consists of B, C, N, O, F, Cl, Br, I, & At. All other elements are “outside
the elbow.”
“Elbow”
Polyatomic ions:
Polyatomic ions:
-ate = 3 oxygen
-ite = 2 oxygen
“Outside Elbow”
Polyatomic ions:
Polyatomic ions:
-ate = 4 oxygen
-ite = 3 oxygen
Note that the –ite ion has one less O than the –ate ion.
Also, the acid that has no hydro prefix and ends in –ic acid has the same number of O as the –ate ion.
The acid that ends in –ous acid has the same number of O as the –ite ion (one less than the –ic acid)
The acid that starts with hydro and ends in –ic has no O.
61
Prelab for Experiment 11: Lewis Structures and the Shapes of Molecules
Name:
______
Lab Day & Time:_____________
Do Lewis Structures for 2 molecules listed in each of the 3 sections in the experiment: binary
compounds, one central atom compounds, and more-than-one central atom compounds.
62
Experiment 11: Lewis Structures and the Shapes of Molecules
Introduction: Chemical formulas indicate what elements are in a compound and how many atoms of
each element are in one molecule of the compound. For example, ethanol has the formula C2H6O;
which tells us that this molecule comes in “packets” of 2 carbons, 6 hydrogens, and 1 oxygen. But
chemical formulas do not indicate how the atoms are connected or how many electrons are between
each pair of atoms. We draw model pictures called Lewis structures to provide such information.
Imagine having some Lego blocks or Tinker Toys. If you have played with these, then you
realize how they can be used to build many different structures – some of the structures look
pleasing and “make sense,” while others might use all of the pieces but really don’t look like
anything except a jumbled mess. Think of atoms of each element in this same sense; we can
imagine connecting the atoms like the toy parts. There are many silly ways to put them together;
however, as you might expect, scientists have figured out that certain “structures” make chemical
sense. And understanding chemistry relies heavily on understanding how the pieces--the atoms--go
together.
Chemists have found that how the valence electrons are spread around the structure is of
utmost importance. Valence electrons are those in the outermost shell, and you need to learn how to
get that number.
Determining the number of valence electrons: We get the valence electron count for some
atoms by simple looking at the last digit of the group number in the periodic chart. Group 1 and 2
elements have one and two valence electrons, respectively. For atoms of the other representative
elements, the number of valence electrons equals the last digit of the group number, or
(equivalently), the group number minus ten. Here are a few examples. Hydrogen, in group 1, has 1
valence electron; Mg, in group 2, has 2 valence electrons; C in group 14, has 4 valence electron (1410); N and P both have 5 valence electrons (group 15); the halogens (F, Cl, Br, and I) in group 17
all have 7 valence electrons.
Lewis structures are drawn using two parts – the chemical symbol abbreviation of the atom
and the valence electrons spread around them. The chemical symbol represents the nucleus of any
atom plus all electrons except valence electrons. Lines (for pairs of electrons) and dots (for single
electrons) are placed around the chemical symbols to represent the valence electrons. Let’s do a few
examples.
63
Now, if we have more than one atom shown in a Lewis structure, then we will have to show
what we think are the “bonds” holding these atoms together. We most commonly use Lewis
structures for covalently bonded molecules, but we can represent ions as well. In the calcium and
oxygen atom symbols shown before, we are representing atoms and their valence electrons. In
molecules we show bonding and non-bonding electrons.
Atoms tend to react to get a full valence shell of electrons. Hydrogen (H) will have a
complete valence shell with 2 electrons (1 pair) and most other atoms have 8 electrons (4 pairs) to
get a complete valence shell. Our job is to figure out how to put the atoms and valence electrons
together to give each atom a full shell.
Drawing a framework—which atoms are connected to which?
There are a few basic rules for drawing Lewis structures later in the lab. The first rule
involves drawing a framework.. Drawing a framework involves deciding which atom(s) are in the
center and which go on the perimeter (bonded to one of the central atoms). Several things affect
this. Let’s look at some examples and some general rules that will help.
(1) Many carbon-containing formulas written with the same element symbol in several
places in the formula are written this way to indicate some of the framework. In this type of
formula, the symbols that follow a central atom are bonded to it. For example, writing the
structural framework for CH2CH2 involves starting with a C and bonding 2 H's to it and then
another C with its own 2 H's.
H H
H
C
C
H
CH2CH2 could have been written as C2H4 to tell you that there were 2 C and 4 H, but C2H4 does
not indicate structure at all. Note that the above structure only shows the connections between
atoms in the compound; it is not a completed Lewis structure.
(2) If you are not given a formula that tells something about the structure, you will have to
decide which is the central atom(s) and which are bonded to it. (a) Often the most symmetrical
arrangement of atoms is correct. (b) Neither hydrogen nor a metal atom will ever be a central atom
for Lewis structures we do in 101, and halogens will be not be central except in anions with only a
halogen and oxygen (such as ClO4-). (c) The more pairs an atom is willing to share, the more likely
it is to be central. Nonmetal atoms farther from F on the periodic table are generally more willing to
share than those closer to F. (d) Carbon atoms are often central, and may often be bonded to each
other as well. (e) Nitrogen, oxygen and sulfur may be central atoms, particularly in molecules that
have only carbon, hydrogen or halogens atoms.
In this exercise, the formula of all the molecules that have only one central atom are written
with the central atom as the the first non-hydrogen atom in the formula.
In drawing the framework, it may also help to remember the following:
H shares one pair of electrons and has zero nonbonding pairs
C tends to share 4 pairs of electrons
N tends to share 3 pairs of electrons and have one pair of nonbonding electrons
O and S tend to share 2 pairs of electrons and have 2 pairs of nonbonding electrons
Halogens tend to share 1 pair of electrons and have 3 nonbonding pairs.
64
Steps for Drawing Lewis Structures
1. Connect the central atoms to each other using a line to represent a bond of two electrons.
Then connect the rest of the atoms around the central atoms, again using a line to represent a bond
of two electrons. Often a more symmetrical arrangement is the correct one.
2. Add up the total number of valence electrons contributed to the Lewis structure. Find the
number for each atom as discussed in paragraph 3 above. If the formula is of an ion, then a.) add
the number of electrons equal to the negative charge or b.) subtract the number of electrons equal
to the positive charge. Now divide the number of valence electrons by two to get the total number of
electron pairs that must be in the Lewis structure.
3. Count the number of electron pairs you used in part 1. Subtract that number from the
number of electron pairs you counted in part 2. If you used all the electron pairs, the structure is
done. Usually at this point you will have electrons left over, and will need to do step 4.
4. For any atom not yet having a full valence shell (2 electrons or one pair for H and 8
electrons or 4 pairs for other atoms), use the extra electron pairs you determined in step 3 to finish
the Lewis structure as follows:
(a) Add pairs of electrons (as lines or double dots) to perimeter atoms first to complete that
atom's valence shell of electrons. Spread them around between various perimeter atoms if there are
several on the perimeter. After the perimeter atoms get the full valence shell, finish with electrons
on the central atom if it needs some.
DO NOT LET THE TOTAL NUMBER OF ELECTRON PAIRS GO OVER THE TOTAL
YOU CALCULATED IN STEP 2.
(b) Occasionally you will run out of electrons before all atoms have completed valence
shells. If this happens, then you will have to modify the Lewis structure by making multiple bonds.
For every pair of electrons you fall short in part 5, move a non-bonding pair of electron from one
atom to make a bonding pair of electrons between an outer atom and a central atom. This will mean
that some of the atoms will have double or triple bonds between atoms.
5. Never, NEVER, NEVER, NEVER, never have more than eight electrons about any period
(row) one or period two element. Never.
6. Go back and check your work. (a) Did you get the correct central atom and bond the
other atoms around it? (b) Does your Lewis structure use exactly the number of electron pairs
allowed? (c) Is each atom's valence shell filled?
Geometry of Groups of Electrons.
One of the uses of Lewis structures is to be able to predict the shapes of molecules. It is very
simple to predict the shapes and then carry that further in predicting properties of molecules. For
now we will be concerned only with shapes or geometries. A simple set of rules and drawings
follows to help with predicting and illustrating molecular shapes:
1. Draw a proper Lewis structure.
2. Determine the number of electron groups surrounding any central atom. Any atom counts as
one group, no matter how many bonds are involved. So a single bonded atom counts as one
group, and a double bonded or triple bonded atom also counts as one group. Also every pair of
non-bonded electrons also counts as a group. So one non-bonded pair counts as one group, two
non-bonded pairs counts as two groups, etc.
3.The total of bonded atoms plus the non-bonded pairs determines the number of groups. These
groups try to move as far apart as possible. The reasoning for them trying to move apart is that
the electrons all have the same energies and that they all are negatively charged [like charges
65
repel]. This limits the shapes or geometries that are available, as shown below, where each bond
is shown as one line regardless of whether it is a single, double, or triple bond. A line with a pair
of electrons at its end is only showing the direction the electrons “point”, and the line itself does
not represent a pair of electrons in this case.
Number of
electron groups
about central atom
Electron
Geometry
around
Central Atom
2
linear
3
trigonal
planar
4
tetrahedral
Sketches
Molecular Geometry
The molecular geometry describes the way the atoms are arranged in space; it differs from an
electron geometry only if there are lone pairs on the central atom. There are a number molecular
geometries, but all of these can be obtained by first drawing the arrangement of the electron
groups as above (showing the correct 3-D angles), and then simply erasing the lone pairs.
When doing this, it will be easier to interpret the shape if you put lone pairs out of the plane of
the paper (with the wedges). For atoms with more than one central atom, it would be best to put
bonds between central atoms in the plane of the paper if possible, and leave any wedges for
perimeter atoms.
Let’s do an PH3 as an example. First we would draw the Lewis structure.
H P
H
H
We see that there are three bonds and one unbonded pair, so there are four groups of atoms around
the P. This means the geometry of the electron groups is shown by the figure on the left below.
To get the molecular geometry, we simply erase the lone pair and its pointing line and thus
obtain the figure on the right
66
MOLECULAR GEOMETRIES
Once you draw the 3-D structure of the whole molecule and erase any lone pairs, you can find the
name of the molecular geometry from the table below. If your’s doesn’t seem to match, make a model
of it and turn it around to see if it matches. If it still doesn’t match, ask your instructor for help.
Geometry name
linear
Sketch
.
O
O
O
O
.
trigonal
planar
.
O
O
angular
or bent
.
O
.
O
.
O
O
O
.
tetrahedral
.
O
O
O
O
O
.
trigonal
pyramidal
O
O
O
.
.
O
angular
or bent
O
.
O
67
Specific Angles Between Groups of Electrons
These geometries have all have specific bond angles associated with the groups. Think of
the central atom as being the locus of an angle defined by lines drawn towards any two of the
o
groups on that central atom. An example showing the 180 angle for a linear set of three atoms is
shown below:
180o
O
O
O
Likewise here is a model showing the angles for a trigonal planar group of atoms:
O
120o 120o
O
O
120o
O
o
The 109.5 angle of the tetrahedral is a little more difficult to show on paper because, unlike the
previous two structures which can have all of the groups on the plane of the page, the tetrahedral is
a three dimensional figure. A model might look something like this:
O
109.5o 109.5o
109.5o
O
O
O
109.5o
O
Four of the six bond angles have been marked to give you some feel for this structure.
Procedure:
1. You are to draw good Lewis structures for each molecule or ion.
2. Then you are to predict the geometry of all electron groups about any central atom. There may be
more that one atom you will need to describe. Make certain to label each one distinctly.
3. Then you are to predict the geometry of just the atoms about any and all central atom(s). Be
certain to include names.
4. Finally you are to give a good 3-D sketch for the molecule or ion.
68
Report Sheet: Lewis Structures And VSEPR Theory
Name_________________________________
Date__________________
Lab Day & Time_______________
List the number of valance electrons for the following elements
C ___
O ___
H ___
N ___
Cl ___
P ___
I ___
S ___
F ___
Br ___
Section 1:
Draw the Lewis Structure for the following Diatomic Molecules
Molecular
formula
Lewis
Structure
I2
O2
N2
HBr
CN-
69
Report Sheet: Lewis Structures And VSEPR Theory
Name______________________
Section 2:
Draw the Lewis Structure and sketch the model. Also, tell the geometry for the central atom.
Lewis
Geometry of
Structure
Sketch
each central atom
CS2
PBr3
H2S
+
H3O
COCl2
70
Report Sheet: Lewis Structures And VSEPR Theory
Name _____________________
Lewis
Structure
Sketch
CHIF2
SO2
SO3
N3
-
SO3
SO4
2-
2-
71
Geometry of
each central atom
Report Sheet: Lewis Structures And VSEPR Theory
Name _____________________
Section 3:
All of the molecules below have more than one atom which should be described as central. Be
certain to give the geometry of all which are central.
Lewis
Geometry of
Structure
Sketch of complete molecule
each central atom
H2SO4
NH2SH
CF2CBr2
CHCH
CH3OH
72
Report Sheet: Lewis Structures And VSEPR Theory
Name _____________________
Lewis
Structure
Sketch of complete molecule
CH3CO-2
(only C's are central)
H2O2
C3H8
CH3CH2OH
CH3OCH3
Write a summary statement.
73
Geometry of
each central atom
Pre-lab for Experiment 12: PREPARATION AND PROPERTIES OF SO2 AND NO2
Name:
______
Lab Day & Time:_____________
1. One of the questions you will be asked to answer in today's lab relates to the determination of
how "clean" the air is, i.e. air quality. In order to answer this question you need to keep in mind
what chemists means when they say they analyze a sample qualitatively or quantitatively.
So:
a) Give an example of a qualitative analysis we have done this semester. Explain what makes this a
qualitative analysis.
b) How would a quantitative analysis differ from this? What would you be trying to figure out
about a sample when you did a quantitative analysis?
2. Give the name or formula for the following compounds:
a) N2O
b) CO2
c) N3O5
d) Sulfur trioxide
e) PCl5
f) NaNO2
g) Carbon monoxide
h) CaSO3
74
Experiment 12: PREPARATION AND PROPERTIES OF SO2 AND NO2
Name: __________________________
Partner: _________________________
Lab Day & Time: _______
Date: ______
Introduction: Nitrogen and oxygen are the main atmospheric gases. Other atmospheric gases such as
Ar, CO2, CH4, and a variety of nitrogen oxides and sulfur oxides are also present, but in much
smaller concentrations, varying by location. In this experiment you will prepare two common
atmospheric pollutants, sulfur dioxide and nitrogen dioxide. You will then examine properties of
these gases.
Purpose:
1) To prepare SO2 and NO2.
2) To examine properties of these gases.
CAUTION: Both of the liquids used in this experiment, 6M sulfuric acid (H2SO4) and 6%
hydrogen peroxide (H2O2), are corrosive materials and should be handled cautiously. SO2
and NO2 are toxic gases. Work in the fume hood and do not allow the gases to escape into the
lab.
Procedure: Follow the procedure as outlined on the following pages.
Report: Complete the attached data sheet. Attach your Summary Statement and answers to the
following questions to the data sheet and hand it in.
Questions (USE YOUR EXPERIMENTAL RESULTS TO SUPPORT YOUR ANSWERS)
1.
2.
3.
4.
5.
Considering your observations when sulfur dioxide was bubbled through water, do you think
that a coal fired power plant that emits fairly large quantities of sulfur dioxide contributes to
acid rain?
Nitrogen oxides emissions are primarily from automobiles. Considering your observations
when nitrogen oxides were bubbled through water, do you think that driving automobiles
contributes to acid rain?
Why was it important that you work in the fume hood when preparing the gases in lab
today?
Did our experiment today contribute to the problem of acid rain? If so, do you think that our
contribution was large or small? Explain your reasoning.
After observing the effects of the gases on the Blepharisma culture, how might blepharisma
be used to determine air quality? Explain for both a qualitative determination and a
quantitative determination of air quality.
75
Sulfuric
acid (6 M)
Sodium
sulfite
(2 g)
The reaction that takes place in the bag:
2H+ + SO32- -> H2O + SO2(g)
I. PREPARATION OF SULFUR DIOXIDE
1.
2.
Place about 2 grams (about 1/2 teaspoonful) of sodium sulfite, Na2SO3 in the bottom corner
of a 1-pint heavy duty Ziploc bag.
Fill a Beral pipette with 6 molar sulfuric acid (H2SO4).
Caution: Sulfuric acid is a corrosive material. Be careful not to spill any on yourself, your
clothing, or the bench top. If you do get any on your skin, wash it off quickly with plenty of
water. If any is spilled on the bench, wipe it up immediately with a wet sponge and rinse
several times with water.
3.
Place the Beral pipette with the sulfuric acid in the plastic bag with the Na2SO3. Smooth out
the bag so it contains a minimum amount of air, and then seal the bag. (Take care not to
press against the pipette.)
KEEP THE BAG IN THE HOOD FROM NOW UNTIL IT IS CLEANED OUT.
4.
5.
Hold the sealed plastic bag as shown in the diagram and slowly squeeze the Beral pipette so
that the acid drops onto the Na2SO3. Keep the bag sealed and in the hood. Record your
observations on the data sheet.
You should now have a sealed plastic bag partially filled with sulfur dioxide, SO2.
II. PROPERTIES OF SULFUR DIOXIDE
1.
2.
Use a Beral pipette full of gas to perform each test on the SO2 you generated in this
experiment.
To get a pipette full of SO2 gas, squeeze the bulb of a dry Beral pipette to expel the air
inside. Keep squeezing the bulb and slowly push the tip of the pipette against the zip seal at
one corner of the plastic bag containing the SO2 gas. With a bit of practice you will be able
to just push the pipette tip so that the seal opens around it. Push the pipette tip into the bag
and release the bulb so that the gas enters the pipette and withdraw a sample of gas. As the
tip leaves the bag, reseal the bag along the "zip strip." Take care not to contaminate the
tip of the pipette by touching the solid or the liquid within the bag. DO NOT FILL
THE PIPET WITH SO2 UNTIL JUST BEFORE YOU USE IT.
76
A. Blepharisma Test
1.
2.
3.
4.
Obtain a microscope slide, clean it with deionized water, and take it to your instructor who
will dispense the Blepharisma culture.
Place the slide on the stage of a microscope and adjust the focus so that the Blepharisma are
clearly visible. Blepharisma are pink, oval, one-cell organisms found in pond water. (There
will probably be 3 or more types of microorganisms visible.)
Observe the nature and movement of Blepharisma. Record what you see on the data sheet in
the space labeled Observations of Blepharisma in the row for the gas, Air.
Fill a Beral pipette with SO2 and then gently squeeze a puff of gas just above the water (do
not bubble gas through the water) containing the Blepharisma. Observe the nature and
movement of the Blepharisma. Record your observations on the data sheet.
B. pH Determination
1.
2.
3.
4.
Fill a cell of a well plate about ½ full of deionized water.
Dip a clean stirring rod into the water and transfer a drop to a piece of pH paper. Compare
the color produced with the color/pH code provided to determine the pH of the water. This
is the control pH, i.e. the pH of water saturated with air, for comparison with the other gases.
Record this result on your data sheet.
Fill a Beral pipette with SO2.
Place the tip of the pipette filled with SO2 into the well containing distilled water. Slowly
bubble the gas through the solution by gently squeezing the pipette. Immediately determine
the pH of this solution as you did above. Record this result on your data sheet.
III. PREPARATION OF NITROGEN DIOXIDE
This gas will be generated in steps. First you will generate, in separate plastic bags, nitric oxide
(NO) and oxygen. These two gases will then be reacted to form nitrogen dioxide (NO2).
CAUTION: Work with this gas in the fume hood. Do not open the bag containing the gas
outside of the hood.
The inner, oxygen generating bag should have pull tabs on the edges of the bag above the zip strip
to help make opening the small bag easier. With a scissors, cut the small bag as shown in the
illustration.
77
A. Preparation of the Oxygen Generating Bag - Inner Bag
1.
Place about 1 g (1/4 teaspoonful) of iron(III) chloride (FeCl3) in the bottom corner of the
small, heavy duty Ziploc bag with the pull tabs.
2.
Fill a Beral pipette with 6% hydrogen peroxide (H2O2).
3.
Place the Beral pipette containing the H2O2 in the small plastic bag containing the FeCl3.
This is the oxygen generating bag.
4.
Smooth out the bag so it contains a minimum amount of air and then seal the bag. (Take care
not to press against the pipettes.) You will place this bag inside the larger bag after it is
prepared (described next).
B. The Nitric Oxide Generating Bag - Outer Bag
1.
Weigh about 1.5 g (about 1/2 teaspoonful) of potassium nitrite (KNO2) and place it in the
bottom corner of the large heavy duty Ziploc bag. (Note: Sodium nitrite (NaNO2) might
be used instead of KNO2. Why would this be an acceptable substitution?)
2.
Fill a Beral pipette with 6 molar sulfuric acid.
3.
Place the Beral pipette with 6 molar sulfuric acid in the plastic bag with the KNO2.
4.
Place the smaller bag inside the larger bag.
5.
Smooth out the bag so it contains a minimum amount of air and then seal the bag. (Take care
not to press against the pipette.) This is the nitric oxide generating bag.
6.
Since most nitrogen oxides are toxic and irritating gases, be sure to keep the plastic bag in
the fume hood.
7.
Generate the oxygen in the small, inner bag by very slowly dropping the H2O2 onto the
FeCl3. (This reaction produces heat as well as oxygen gas. If the reaction takes place
quickly, the heat from the reaction will melt a hole in the plastic bag.) Record your
observations on the data sheet.
8.
Without opening either bag, generate NO in the outer bag by slowly squeezing the Beral
pipette so that the 6 molar sulfuric acid slowly drops onto the potassium nitrite. When the
reaction is complete this ZiplocTM bag should be partly filled with NO (and possibly some
NO2 ). NO is colorless; NO2 is rusty brown. Record your observations on the data sheet.
9.
Without opening the larger bag, open the smaller bag and observe and record what happens.
10.
The larger bag now contains a mixture of nitrogen oxides to be used in the tests below.
11.
The bags should be kept in the fume hood since nitrogen oxides slowly diffuse through the
plastic bag.
78
Reactions that take place inside the bag:
Step 7:
2 H2O2 ___> 2 H2O + O2
Step 8:
6 KNO2 + 3 H2SO4 ___> 4 NO + 2 HNO3 + 2 H2O + K2SO4
Step 9:
2 NO(g) + O2(g) ___> 2 NO2(g)
IV. PROPERTIES OF NITROGEN OXIDES
1.
Using the same technique used to fill a Beral pipette with SO2 gas, fill a dry Beral pipette
with NO2 gas. Use a Beral pipette full of gas to perform each test. But remember, don't fill
the pipet until just before you use it.
A. Blepharisma Test
1.
2.
Obtain another sample of the Blepharisma culture from your instructor.
Using the same technique as before, observe the Blepharisma before and after squeezing a
puff of NO2 gas over the water containing the Blepharisma. Record your observations.
B. pH Determination
1.
2.
3.
Fill a cell of a well plate about ½ full of deionized water. Try to use the same volume as in
the previous pH test.
Place the tip of a NO2 filled pipette into the cell and slowly bubble the gas through the
solution by gently squeezing the pipette.
Immediately determine the pH of the gaseous solution as before. Record your results in the
data table.
CLEAN UP
1.
2.
3.
4.
5.
Open the bags containing the nitrogen oxides and the sulfur dioxide in the back of the fume
hood.
Allow all of the gases to be vented into the hood. Squeeze the pipettes several times to flush
gases from them.
Using the water supply in the hood, fill the bags with tap water. Carefully pour the water
into the cup sink at the back of the hood.
Still working in the hood, rinse the bags and pipettes with tap water and then discard them in
the trash.
The wellplates should be washed with tap water, rinsed 3 times with small amounts of
deionized water, and allowed to dry.
79
Report Form: Preparation & Properties
of SO2 and NO2
Date
Name ________________________
Partner _______________________
Lab Day & Time _______________
1. Observations regarding the generation of sulfur dioxide (SO2).
2. Observations regarding the generation of nitrogen oxides (NO & NO2).
3. Record the pH and the response of Blepharisma when the indicated gases are added to water or
puffed over a Blepharisma culture.
Gas
pH
Changes Observed in Blepharisma
Air
SO2
NO/NO2
80
Prelab for Exp. 13: Acid-Base Classification I
Name:
Date:_________
Lab Day & Time:___________
After reading the write-up for this experiment and reviewing the procedure used to construct
a classification scheme in Experiment 1, answer the following questions. Hand in your answers at
the beginning of the laboratory period.
1.
Assume that you have solutions numbered 1-10 that were tested as described in this
experiment and you got the following results:
Blue Litmus turned red when solutions 1, 2, 4, and 5 were placed on the paper. No other
solutions changed the color of blue litmus.
Congo Red turned blue when solutions 4 and 5 were placed on the paper. No other solutions
changed the color of congo red.
Alizarin Yellow-R turned from yellow to red when it was added to solutions 3, 6, and 8.
No other solutions caused this color change.
Use these data to prepare a classification scheme which separates the ten solutions into
groups that have the same properties (refer back to Exp.1 for review of classification
schemes)?
2.
A classification scheme that is based on properties that are difficult to observe will lead to
classification errors. Last year a student obtained the same results when the compounds
used in this experiment were tested with red litmus paper and when they were tested with
phenolphthalein. When tests with red litmus were done, he and his partner argued about
several results before deciding whether there had been a color change. They had no
disagreements about the results of the tests with phenolphthalein. Since both tests provide
the same information, either set of results will serve the same purpose in preparing a
classification scheme, but using both sets will not result in a greater separation of the
compounds tested. Which set of results would you advise the Students to use? Why?
81
Experiments 13-14: Acid-Base Classification I & II
Introduction: This activity will require two laboratory periods First, you will see how different
acids and bases react with a variety of other compounds and use the data to construct a classification
scheme. In the second laboratory period you will use your classification scheme developed during
the previous week to classify a variety of compounds.
Compounds that release hydrogen ions (H+) are acids, and compounds that accept hydrogen
ions are bases. Many bases contain hydroxide ions (OH-) or react with water to form hydroxide
ions, and it is these hydroxide ions that accept the hydrogen ion. Stronger acids and bases release
more hydrogen and hydroxide ions, respectively. If fewer ions are released, they are described as
weaker. These chemical facts are the basis for constructing the classification scheme in this
experiment. In lecture you may learn to classify a strong acid and base by whether 100 percent of
the molecules provide hydrogen or hydroxide ions. This is different than saying stronger acids and
bases, which simply provide lots of hydrogen or hydroxide ions; the –er makes a difference!
Part I: Preparing a Classification Scheme for Acids and Bases: Compounds can be classified
based their response during a set of tests. Certain characteristics tend to occur together, and over
time scientists have assigned names to compounds with particular sets of characteristics. This is
what you will do in week one-you will find which characteristics tend to go together. You will test
the solutions using three test papers and two test solutions.
1) Red Litmus Paper: Red litmus paper is prepared by soaking paper in a slightly acidic
litmus solution. Some compounds will change the color of litmus from red to blue.
2) Blue Litmus Paper: This is prepared by soaking paper in a slightly basic litmus solution.
Some compounds will change the color of litmus from blue to red.
3) Congo Red Paper: Similar to litmus paper, this test paper is prepared by soaking paper in a
solution of congo red. Some compounds will change the color of congo red to blue.
Procedure for Test Papers: Place each test papers on a clean surface such as a clean sheet of
notebook paper. Dip a clean stirring rod into the solution you wish to test, and touch the rod to the
test paper. Observe carefully to see if there is a color change. Wet paper looks different from dry
paper, so test with plain water first. That will help you distinguish the color change that you are
looking for from the change in appearance due to wettingYou should test several solutions on a
single strip of test paper. In these tests you should not worry about whether the change is moderate
or drastic—if one blue paper turns pink and another dark pink, just record both as pink.
4) Phenolphthalein (fee-null-tha-lean): This compound is an acid-base indicator like litmus
and congo red. It changes from colorless to pink.
5) Alizarin Yellow-R (uh-liz-uh-wren): Another indicator, it changes color from pale yellow
to pale red (or orange).
Procedure for Indicator Solutions: Place about 1 mL of the solution to be tested in a clean test tube
and add 1-2 drops of the indicator solution. Look for a color change. If the color solution becomes
paler, this is due to dilution and should not be interpreted as a color change.
In addition to causing several compounds to change color, acids and bases have other
properties that can be used to classify them. Some of the compounds you will test will react with
carbonates and bicarbonates to produce CO2 gas. Perform the following test on each solutions:
Bicarbonate: Place 1 mL of the solution to be tested in a test tube and add several drops of 1M
NaHCO3 solution. (What are you looking for?) This reaction is fast, so watch closely as you add
the solution. If you don't see a reaction, mix and then observe again, just in case.
82
1.
2.
Use each test outlined above to test each of the numbered solutions, recording the result in
Table 1.
Use the data in Table 1 to prepare a classification scheme that will separate the solutions
into groups. (See Experiment 1.) When you finish, the compounds in each group should
produce the same results in each test that you performed. If all of your tests were done
carefully, every one of the original 15 solutions should be in one of the final groups. Be
sure to hand in your classification scheme before you leave lab during Week 1.
Reminders on classification schemes: once solutions are split into two categories, don’t recombine
them again. If a question doesn’t split a group (that is, the answer for all solutions in a subgroup is
yes or the answer for all solutions is no), then don’t use that question.
Next week you will classify various household products. Pick up a clean vial and label
before you leave lab today so you can bring in a sample of a commercial product you would
like to test. Be sure to record product name and brand name on the label.
83
Report Sheet: Acid-Base Classification I
Name________________________
Date_________
Partner______________________
Lab Day & Time_____________
Use each test outlined in the procedure for Part I to test each of the numbered solutions, recording
the result in Table 1. (The table will be easier to interpret if you draw a line in each cell that does
not produce a change.)
TABLE 1
Red Litmus
Blue Litmus
Congo Red
#1
#2
#3
#4
#5
#6
#7
#8
#9
# 10
84
Phenolphthalein
Alizarin
Yellow R
Bicarbonate
Prelab for Experiment 14: Acid-Base Classification II
Name:__________________________
Time:_______________
Date:_________
Lab Day &
The classification scheme shown below is based on experimental data collected in this
experiment last year. The names given to the six classes that can be identified with this scheme are
shown in parentheses under the letters represented the 15 solutions used to develop the
classification scheme. Use the information in the scheme to answer the questions below the
scheme.
|
Solutions A-R
|
Blue Litmus Turn Red?
|
YES
B,C,E,N,Q
|
Bromocresol Purple Turn Yellow?
|___________|
|
YES
NO
B,C,E,Q
N
|
(Very Weak Acid)
Metacresol Purple Turn Yellow?
|
|
|
YES
NO
B,C,E
Q
(Strong Acid)
(Weak Acid)
|
NO
A,D,F,G,H,J,K,M,P,R
|
Red Litmus Turn Blue?
|
|
|
NO
YES
A,H,J,M
D,F,G,K,P,R
(Neutral)
|
Red in Alizarin Yellow R?
|
|
|
NO
YES
D
F,G,K,P,R
(Weak Base) (Strong Base)
Questions:
1.
If you were using this classification scheme to determine whether the juice from rhubarb is a
strong acid, weak acid, very weak acid, neutral, a weak base, or a strong base, which test
should you do first? Why?
2.
If the juice from rhubarb turns blue litmus red, would it be necessary to test it with Alizarin
Yellow R in order to classify it? Why?
3.
If juice from rhubarb is a very weak acid, what is the minimum number of tests you would
need to run before you could classify it?
4.
Give the name or formula for the following acids and bases:
a) Sulfuric acid
b) NaH2PO4
c) NaOH
d) Potassium hydrogen carbonate
e) Sodium hydrogen phosphate
f) Ammonium hydroxide
g) Hydrochloric acid
h) NaHSO4
i) NH3
j) Sodium phosphate
85
Part II (Week 2): Using Your Acid-Base Classification Scheme: Once a classification scheme
has been constructed, it can be used to classify things other than those used to construct the scheme.
Last week you constructed a scheme for classifying compounds, but you did not have information
that could be used to name the groups of compounds that were obtained using your scheme. At the
beginning of lab today, your instructor will tell you the names of the compounds tested last week.
With that knowledge and information about acids and bases found in your textbook, the class
should be able to label the end groups of the classifications scheme as stronger acids, weaker acids,
neutral, weaker bases and stronger bases.
1.
2.
You will participate in a group discussion led by your instructor and arrive at a final
classification scheme and appropriate class names for acids and bases.
After names have been assigned to each subclass in the classification scheme, use it to
classify these household materials:
1. shampoo
3. hand soap
5. The Works drain cleaner
7. apple juice
9. household cleaner
2. automatic dishwasher soap
4. wine
6. Drano
8. soft drink
10. glass cleaner
11. bleach
12. your sample
Part II Procedure: You will use the same tests that you used last week, but you need not perform all
tests on each product. Use the minimum number of tests required to classify each item; a maximum
of 3 tests should be needed. If more than 3 are needed, speak with your instructor about your
choices. [Before going on, study your classification scheme and decide a) which test you should do
first, b) which test you will do second if the first test is positive, c) which test you will do second if
the first test is negative, and so forth. Be sure to write down in Table 2 the test you are going to do
before you do it.]
3.
After you decide how you will go about testing the materials, fill in Table 2.
SAFETY NOTE: Be careful not to mix the bleach with any other product. Flush all bleach
solutions down a hood drain separately.
When you test any item that contains bleach, look at the color of any indicator paper or solution as
soon as you add it. Sometimes the bleach reacts with the indicator and most color is lost, but this is
usually slow, so you can still see if there is a color change if you look right away.
86
Report Sheet: Acid-Base Classification (cont.)
Name________________________
Date_________ Lab Day & Time_____________
Partner______________________
Use the classification scheme that you developed last week to test and classify each household
product listed in Table 2. Write in the brand name of the specific shampoo, hand soap, etc. that you
test.
TABLE 2
Product
Name of Test
Test Results
Shampoo
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
The Works
drain cleaner
1.
2.
3.
1.
2.
3.
Drano
1.
2.
3.
1.
2.
3.
Apple Juice
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
1.
2.
3.
______________
Dishwasher Soap
______________
Hand Soap
______________
Wine
______________
______________
Bleach
______________
Glass cleaner
______________
Soft Drink
______________
Household Cleaner
______________
______________
______________
Classification
Since it was not necessary to use all of the tests that you ran last week in order to classify the
household products, you did not run every test on each product. However, using the results from
last week and this week, you should be able to predict what would happen if you did do each test.
In Table 3 place an X in each cell representing a test that you ran on each household product.
(The results for those tests are recorded in Table 2.) In the remaining cells, predict what would
happen if the test were run.
87
TABLE 3
Product
Classification
Red Litmus
Blue
Litmus
Congo
Red
Shampoo
Dishwasher Soap
Hand Soap
Wine
The Works
drain cleaner
Drano
Apple
Juice
Bleach
Glass Cleaner
Soft Drink
Household
Cleaner
_________
Summary Statement
88
Phenolphthalein
Alizarin
Yellow R
Bicarbonate
Prelab Experiment 15: pH and Buffers
Name__________________________
Date_________
Lab Day & Time__________
The exercise today is designed to give you experience in measuring pH, working with buffer
solutions and calculating [H+] from the pH.
1. If the scale below represents solutions of various pH values, indicate which solutions are
considered basic, neutral or acid.
0---1---2---3---4---5---6---7---8---9---10---11---12---13---14
2. a) Define buffer in terms of function, i.e., what does it do?
b) Define buffer in terms of its components, i.e., what is it made out of?
3. Is a solution of HCO3- and CO32- a buffer solution?
4a. Write an equation for the reaction of H+ with a solution containing both HCO3- and CO32- .
4b. Write an equation for the reaction of OH- with a solution containing both HCO3- and CO32- .
5. Is a solution of NaHCO3 and Na2CO3 a buffer solution? Explain using chemical equations.
6. Give the name or formula for the following:
a) Sodium bicarbonate
c) NaC2H3O2
b) HC2H3O
d) Sodium carbonate
89
Experiment 15: pH and Buffers
Name:________________________________
Partner:______________________________
Lab Day & Time:_____________
We are going to test 3 acid/conjugate base pairs to see which are buffers.
Fill the labeled vials just to the bottom of the label with the appropriate solutions.
Measure the pH as instructed. Add 10 drops of 0.1 M NaOH to the acid/base pairs; measure the
pH again.. Record all values in the table below
Table 1: pH Values
Solution
Initial pH
Initial [H+]
pH after 10
drops NaOH
[H+] after 10
drops NaOH
Distilled H2O
Acetic acid/
Sodium Acetate
H2PO4- /
HPO42HCl/NaCl
Which of the acid/base pairs are buffers? Use your experimental results to support your answer.
Use theory to explain how the pairs that are buffers work. To aid in this explanation, write
chemical reaction equations (net ionic) for each buffer solution with H+ or OH- as reactant to show
how the pairs that are buffers work when acid and when base are added. Note: equations alone
aren’t sufficient; you need to discuss how the buffers work.
Summary Statement.
90
Check-Out
Obtain your check-out sheet from your instructor. Make sure you have all the equipment listed and
that it is all clean and in working order. Make sure your drawer and hood are clean. Your
instructor will check your drawer, and assign you a lab-cleaning chore. Failure to check out
properly will result in a ‘0” lab grade being figured into your average.
91
GLOSSARY FOR CHEMISTRY 101
In this glossary, the following conventions are followed:
The term being defined is presented in bold type.
If a word is used in the definition that is defined elsewhere in the glossary, it is indicated in italics.
When an supplemental description of the term is shown, it is shown in small letters. These descriptions are not a part of the
definition but often help explain a common context in which the term is used.
Abbreviated electron configuration. An electron configuration in which the nearest previous
noble gas in [ ] is followed by the electron configuration for the remaining electrons.
Absolute temperature scale. The temperature scale whose zero is the temperature at which the
volume of an Ideal Gas would be zero and whose scale divisions are the same size as the Celsius
degree.
This temperature scale is more commonly called the Kelvin scale.
Accuracy. The numerical difference between the experimental value for some quantity and the true
or generally accepted value for that quantity.
Accuracy is inversely related to error. The larger the error is, the lower(or worse) is the accuracy.
Acid. A substance that can act as a proton donor.
The definition given is the Bronsted-Lowry definition. The Arrhenius definition of an acid as a substance that forms H +in
water and the Lewis definition of an acid as an electron pair acceptor will not be used extensively in this course.
Acid-Base Reaction. A proton transfer reaction.
Acid Dissociation Constant. Same as acidity constant.
Acid Ionization Constant. Same as acidity constant.
Acidity Constant. The equilibrium constant for the proton transfer from an acid to water.
Alternative definition:
[ H3O ][ A  ]
HA + H2O  H3O+ + A  Ka = Keq 
[ HA ]
Actinide. One of the elements 89-103 (Ac  Lr).
These elements and/or their ions have partially filled 5f energy levels.
Activated Complex. A pseudo-molecule that is formed as reactants go to the products in a
mechanistic step.
This must not be confused with the term intermediate because an activated complex is not a chemical species and exists only
in the process of passing from reactant to product in the step It has a lifetime of only one vibration and can decompose either
into the reactants or products of that step.
Activation Energy.
A quantity that is determined empirically from the dependence of the rate of reaction on the
temperature. Note this is only a description since it does not define how this quantity is measured. Activation energy is
commonly interpreted as the energy change in passing from the reactants to the activated complex of the rate-determining step
of the mechanism.
Actual Yield (of a substance). The amount of a single substance isolated from a chemical reaction.
92
Alkali. An obsolete term for a substance that produces a basic solution when dissolved in water.
Alkaline. An obsolete term for a basic solution, or substance.
Alkali Metal. One of the elements from Column 1 (IA) in the Periodic Table.
These are the elements Li, Na, K, Rb, Cs, and Fr.
Alkaline Earth Metal. One of the elements from Column 2 (IIA) of the Periodic Table.
These are the elements Be, Mg, Ca, Sr, Ba, and Ra.
Allotropes. Forms of an element that have different properties from each other.
The word "form" is ambiguous. Allotropes have different molecular formulas or different crystalline structures and are
different substances even though they have the same elemental composition.
Alloy. A mixture in which a metal is the solvent.
An alloy usually contains two or more metals and it usually appears homogeneous when viewed by eye. It may contain nonmetals as solutes and may be either homogeneous or heterogeneous when viewed with a microscope. It rarely occurs that an
alloy is a compound of more than 1 metal.
Alpha () Particle. A particle consisting of 2 protons and 2 neutrons, that can also be notated as a
4He nucleus..
Alpha () Ray. A stream of  particles.
Amorphous solid. A solid without a regularly repeating three-dimensional structure.
Amorphous
solids are commonly mixtures though they may be pure substances. They do not have a regularly repeating three-dimensional
structure and may also be called non-crystalline solids.
Amphoteric. A material that can act as either an acid or a base in aqueous solution.
The acid/base behavior of an amphoteric material depends on what other substances are present in the solution.
Anion. An ion that has a negative charge.
Angstrom Å. A unit of wavelength; 1 Å = 10-10.
Aqueous. In water.
Atmosphere.
1. A unit of pressure defined as: 1 atm = 101,325 pa.
2. The gaseous surroundings of the earth.
3. The gaseous surroundings of an experiment.
Atmospheric Pressure. The pressure measured with a barometer.
This experimental number is interpreted as the pressure due to the mass of air directly above the barometer from the top of the
atmosphere down to the barometer.
Atom. The smallest unit that, when replicated, will generate a sample of an element.
Atom is frequently used more loosely than this to indicate the nucleus and the kernel electrons of a atomic particle, regardless
of the number and state of its valence electrons.
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Atomic Mass. The mass of an atom in atomic mass units.
Atomic Mass Unit. 1 amu is exactly 1/12 of the mass of a 12C atom.
Atomic Number. The number of protons in the nucleus of an atom.
Atomic Orbital. A space, centered on an atom, that may be occupied by an electron.
This term is also used to describe a map of this space.
Atomic Radius. In an element, this is one-half of the distance between nuclei. In a compound, the
sum of the atomic radii for bonded atoms must equal the distance between their nuclei.
The space between bonded atoms is divided between the two atomic radii by a combination of experiment and theory to find
the atomic radius of each atom.
Atomic Size. Usually the same as atomic radius.
Atomization energy. The amount of energy absorbed when one mole of an element in its normal
state is converted to one mole of gaseous atoms.
Aufbau Principle. The theory that atoms have hydrogen-like atomic orbitals to hold electrons.
The electrons are placed in these orbitals, one at a time starting at the lowest energy orbital, until
all the electrons in the atom are present. This allows the electron configuration of the atom to be
predicted.
The ordering of orbitals by energy is standard, and each orbital can be occupied by a maximum of two electrons. The results
of this procedure are not always correct but the failures mainly indicate an inaccuracy in the standard order of orbital energy
ordering for the atom in question.
Avogadro's Number. The number of particles in exactly one mole of particles.
Balancing. The process of changing the coefficients in a chemical skeleton so that the number of
symbols indicated for each element is the same on each side of the arrow.
This is the process by which a chemical skeleton is converted into a chemical equation.
Barometer. Any instrument that measures the atmospheric pressure.
Base. A species that can serve as a proton acceptor.
This is the Bronsted definition as contrasted to the Arrhenius definition that states a base is a source of hydroxide ion, or the
Lewis definition that defines bases as electron pair acceptors.
Basic. A solution with pH > 7 or a substance that produces such a solution when dissolved in
water.
Basic Ionization (Dissociation) Constant. Same as basicity constant.
Basicity Constant. The value of the equilibrium constant for the acid-base reaction of a substance
with water as the reference acid.
[ BH ][ OH  ]
B + H2O  BH+ + OH- Kb = Keq =
[ B]
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Beta () Particle. An electron emitted from a nucleus in a nuclear reaction.
Beta () Ray. A stream of beta particles from a radioactive sample.Bimolecular Process. A
mechanistic step in which two species collide and change to one or two new species.
Binary Compound. A pure substance containing only atoms of two different elements.
Boiling Point. The temperature at which the vapor pressure of a liquid is equal to the external
pressure on the liquid.
This is observationally determined as the temperature at which bubbles of the vapor of the liquid can form anywhere in the
liquid. This point clearly depends on the external pressure.
Bond. Same as chemical bond.
Bond Energy. The amount of energy absorbed when one mole of identical bonds are broken
between two atoms in a mole of gaseous molecules to form 2 moles of neutral fragments..
X-Y(g)  X(g) + Y(g) Bond Energy = H orxn
The bond energy is stated to be for the two bonded atoms. Even though the rest of the molecule may be involved, this can be
approximately found as a tabulated value since it is a good approximation that the bond energy depends only on what atoms
are bonded and their bond order.
Bond Length. The distance between the nuclei of two bonded atoms.
Bond length is an experimental quantity. This value can be approximated for singly bonded atoms as the sum of the radii of
the two bonded atoms.
Bond Order. The number of electron pairs shared between two atoms.
Bronsted Acid or Base. See acid or base.
Buffer Solution. A solution containing both a weak acid and a weak base.
A buffer solution resists change in pH because any addition of strong acid or base is offset by reaction with the weak base or
weak acid in the solution. Buffers (buffered solutions) are most commonly prepared with a mixture of a weak conjugate acidbase pair and their pH is determined by the ratio of the concentrations of these two species.
Calorie. 1 calorie = 4.184 Joule
1 food Calorie is called a Large calorie and indicated with a capitol C. 1 Calorie = 1000 calorie
Calorimetry. The measurement of enthalpy changes for an experimental system.
Catalyst. A material that increases the rate of reaction but it does appear in the equation for the
reaction because it is not chemically changed at the end of the reaction..
A catalyst usually changes the entire mechanism of a reaction and is involved in at least two steps in which the catalyst is
chemically changed and then regenerated.
Catenation. The ability of atoms of an element to form bonds with each other so that stable pure
substances can contain chains of these atoms.
Cation. An ion that has a positive charge.
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Celsius. The temperature scale that has a 0 at the freezing point of water and has divisions
(degrees) set by choosing the boiling point of water as 100oC.
The Kelvin scale is related to the Celsius scale by having the same size degree, but a different method of establishing a 0 is
used.
Charge. The property an object has because of a surplus (-) or deficiency (+) of electrons
compared to the number of protons. The SI unit of charge is the Coulomb (C).
The charge of one electron is -1.6 x 10-19 C. Alternatively, the charge of one mole of electrons is -9.649x104 C.
Chemical. As a noun, the same as material or substance in common speech. As an adjective, it
describes a process in which a substance is changed to a different substance.
To a chemist, the term chemical is simply another way of saying "substance".
Chemical Bond. The attractive forces that act between atoms that are attached in compounds.
Chemical Composition. The ratio of the amounts of each element to the amount of compound.
Chemical composition may be expressed in a number of ways, and the above definition is just one possible method.
Chemical Equation. An accurate description of a chemical change. The formulas of the reacting
substances are written to the left of an arrow separated by " + " signs and the formulas of the
substances produced are written on the right of the arrow, separated by " + " signs. Each formula is
preceded by a numerical coefficient which states the mole ratio in which that substance reacts or is
formed..
Note that since atoms are conserved, no description of a chemical reaction is accurate unless it is balanced. If no coefficient is
shown, then the coefficient is "1" .
Chemical Equilibrium. A state of a chemical system in which, after some reaction has occurred,
no further changes in concentration can be observed.
Note that if a system simply does not seem to change with time, there is no guarantee that an equilibrium state exists since the
reaction may be extremely slow.
Chemical Formula. A statement of the composition of a substance in which the elemental
symbols for the elements present are written with their mole ratios in the compound written as
subscripts.
The simplest kind of formula is called an empirical formula. More informative chemical formulas are molecular formulas and
structural formulas. The subscripts can also be interpreted as the number of moles of element in 1 mole of formula units.
Chemical Kinetics. The study of chemical reactions in order to elucidate their mechanisms.
The more common use of the term by non-chemists is that chemical kinetics is the study of the rates of reaction.
Chemical Property. Any property of a material whose measurement depends on the chemical
composition of the materials used to make the measurement .
This is a half of the classification scheme for the properties of materials, chemical or physical properties. Chemical properties
are determined experimentally by observing whether or not the material is changed into a material of different composition by
the presence of a test reagent. Physical changes are changes that result in no change in composition and are independent of the
materials used in the test procedure.
Chemical Reaction. Any process in which at least one substance is changed to at least one
different substance.
Chemistry. The study of the properties of materials and the nature of their reactions with each
other.
Closed System. A system that allows the exchange of energy but not mass with the surroundings.
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Coefficients. The numbers in chemical equations that state the mole ratios of the substances whose
formulas follow them in the equation.
Colligative Properties. Properties of solutions that depend on the number of solute particles in the
solution but do not depend on the identity of the solute particles.
Complex ion. A species found in solution or a solid, that consists of a central metal ion
surrounded by neutral molecules or anions bonded in a definite geometry to the metal ion.
This definition is limited because a complex ion may contain more than one central metal ion. The molecules or ions bonded
to the central metal ion are called ligands. Sometimes, when the only ligand attached to the metal ion is water, the aggregate
is not called a complex ion but simply an aqueous metal ion.
Compound. A substance composed of two or more elements chemically bonded in a fixed ratio.
This is a weak definition since it requires a knowledge of what the rather ambiguous term "chemically bonded" means. At a
microscopic level, a compound is more easily defined as a containing a basic chemical unit that is composed of more than 1
element and that is replicated many times to make a sample of the compound.
Compressibility. The response of the volume of a sample to external pressure. The larger the
response of V / P is, the larger the compressibility.
Concentrated. A qualitative description of a solution that implies the concentration of solute is
relatively large.
Concentration. The ratio of the amount of solute to the amount of solution or solvent.
Condensation. The process in which a gas is converted to its liquid or solid.
Condensed phase. A solid or liquid.
Conjugate Acid-Base Pair. Two chemical species which differ by a single H+ .
Coordination Number. The number of neighboring atoms (or ions) surrounding the atom or ion
being considered.
Coulombic. Any action or property of particles that results from their charges.
Covalent Bond. A chemical bond in which the attraction between the attached atoms is due to the
coulombic attraction between shared electrons and the nuclei of those atoms.
This type of bond is commonly stated as being due to the sharing of electrons.
Covalent Compounds. Compounds in which the atoms are bonded to each other only by covalent
bonds.
Critical Pressure. The vapor pressure at the critical temperature.
Critical Temperature. The highest temperature at which there is a distinction between liquid and
gas.
Above this temperature, only a single fluid phase can exist that is called a supercritical fluid can exist.
Crystal Structure. The geometrical arrangement of the repeating units in the crystalline solid, and
the geometrical arrangement of atoms within these units.
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Crystalline Solid. A solid that consists of a single repeating unit that is replicated on a regular
three dimensional pattern called a lattice.
Many chemists use the term "solid" to mean a crystalline solid even though they may not be explicit about this usage.
Crystallization. The process in which a solute is converted to its pure solid form in contact with
the solution from which the solid is formed.
Dalton's Law of Partial Pressures. The total pressure of a mixture of gases is the sum of the
partial pressures of the component gases in the mixture.
Decomposition. A process in which one substance is converted into several simpler substances.
Simpler substances are substances with smaller formula units than the original substance.
Degenerate orbitals. Atomic or molecular orbitals whose electrons have the same energy.
Density. The ratio of mass to volume for a material.
Alternatively, the equation: density = mass/volume, is a definition of density. Some people prefer to describe density as a
property whose measured value indicates the "compactness of the mass".
Diatomic molecule. A molecule that consists of two atoms.
Diffusion. The movement of microscopic particles from regions of higher concentration to regions
of lower concentration as a consequence of their kinetic energy.
All atoms, molecules, and ions have kinetic energy. Since they are moving, they will spread out into any space available to
them, i.e., they will diffuse until their concentration is uniform in the volume to which they are confined..
Dilute. A qualitative term that is used to describe solutions containing a relatively small
concentration of solute.
Dilution. The process of making a less concentrated solution from a more concentrated solution.
Dimer. A molecule made from two identical molecules.
Dipole moment. The product of the charge and the distance between the charges.
Dipole-dipole attraction. The intermolecular attraction between molecules that have separation of
positive and negative charge.
Directly proportional. A relationship that follows the equation y = mx.
Dispersion forces. Same as London Dispersion Forces.
Disproportionation reaction. A chemical reaction in which a single species is converted into at
least two other species, each of which contains the part of the original species that was of interest.
This term is most commonly used in redox reactions in which an atom in the reactant is converted into two different products
in which that atom has two oxidation numbers, one higher and one lower than the original reactant.
Electrolysis. A process in which a non spontaneous chemical reaction occurs when an external
electrical source is applied to the chemical system.
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Electrolyte. A material that produces an aqueous solution that conducts electricity better than pure
water.
Electron. A subatomic particle that has a -1 atomic charge and a mass of .005 amu.
In an atom, the electrons are external to the nucleus and randomly scattered in its space. In SI units, the charge of the electron
is -1.60 x 10-19 C and its mass is 9.1 x 10-31 kg.
Electron affinity.
Either: 1. The amount of energy absorbed when one mole of electrons is absorbed by one mole of
gaseous atoms.
or, more general:
2. The standard enthalpy change for the following process:
A(g) + e  A-(g) H orxn = Electron Affinity of A
"A" may be an ion, molecule, or atom.
Electron configuration. A listing of the type of occupied atomic orbitals in an atom with the
number of electrons in the orbital type shown as a superscript.
The standard orbital designations are used and they are grouped ignoring the orientation possibilities. The orbitals are usually
stated in order of increasing energy. See also abbreviated electron configuration and valence electron configuration.
Electron density. The ratio of electron charge to volume at a point in space. An alternate interpretation
of this term is the probability that an electron will found at a particular point in space.
Electronegativity. The ability of a bonded atom to attract shared electrons toward itself in a
covalent bond.
There is no general agreement on how this term is measured. The Mulliken definition is commonly used to explain the
general experimental sense of electronegativity. This definition is:
Electronegativity = k * (Electron Affinity + Ionization Energy)
Electrostatic. Same as Coulombic.
Element. A substance containing only atoms with the same atomic number.
Elementary process. A simple collision of two molecules, atoms or ions resulting in a chemical
change or a decomposition of a single molecule into two fragments in a mechanism.
This is usually called a "step" for simplicity.
Emission spectrum. The spectrum of light emitted by a sample.
Normally emission spectra are measured on samples that have been excited using thermal or electrical energy.
Empirical. An adjective that means derived from an experiment.
Empirical Formula. A chemical formula in which the smallest subscript is
" 1 " or in which the subscripts are the smallest possible non-zero integers.
Endothermic process. A process in which energy is absorbed from the surroundings.
Energy. The sum of the kinetic energy and the potential energy for a system.
This is really not a definition, but it as more useful than the generally used ""capacity to do work" which has little meaning to
most people.
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Energy level. One of the allowed energy states of an atomic or molecular system. These are
quantized and discreet.
Atomic and molecular electronic energy levels are closely associated with the orbitals that have these energies so the terms are
frequently interchanged.
Enthalpy change (H). The amount of energy absorbed by a system during a change that occurs at
constant pressure.
We use this term as synonymous with the energy change for a system. "Enthalpy" is difficult to define since, being an energy,
it is impossible to define its zero value.
Enthalpy of fusion (Hfus) The enthalpy change when one mole of a solid pure substance is
converted to a liquid.
Enthalpy of reaction (Hrxn). The amount of energy absorbed in a chemical change when a
reaction occurs at constant temperature and pressure.
Enthalpy of solution. The enthalpy change when 1 mole of solute is dissolved to make a solution
of a stated concentration.
Enthalpy of vaporization (Hvap). The enthalpy change when one mole of liquid substance is
converted to gas.
Entropy (S). A thermodynamic quantity that states the amount of geometrical order in a system.
For a perfectly ordered system, its value is zero. Its numerical value is a maximum for a completely
disordered system.
This is only a description, not a definition (since it does not state how entropy is calculated or measured experimentally).
Enzyme. A biological catalyst.
Invariably, the largest portion of an enzyme molecule is protein.
Equilibrium. A state in which there are no further observable changes with time.
Equilibrium constant. A number equal to the ratio of the equilibrium concentrations of products
to the equilibrium concentrations of reactants, each species being raised to the power of its
coefficient.
Alternatively, the equilibrium constant is defined for a generic reaction as below.
For the reaction : aA + bB  cC + dD,
Keq = [C]c[D]d / [A]a[B]b
The square brackets [ ] mean the concentration of the enclosed species at equilibrium. The equilibrium constant for a reaction
can be found a number of ways and is most commonly found from tabulated experimental data.
Error. The numerical difference between the measured value for a quantity and the "true" or the
accepted value for that quantity.
Error is most commonly stated as relative error which is the ratio of the error to the "true" or accepted value for that quantity.
Almost always, this relative error is multiplied by 100% so that becomes a percent error, i.e., %error = 100% * (xptl value true value)/true value.
Evaporation. The process of converting a liquid to a gas.
Note that evaporation can occur at any temperature and occurs at the surface of the liquid, whereas boiling can only occur at a
specific temperature (see boiling point) and can occur any place in the liquid, hence bubbles.
Excess reagent (reagent in excess). A reactant in a chemical reaction that would not be consumed
entirely if the expected reaction were complete.
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Excited state. A state for an atomic or molecular system that is of higher energy than the ground
state.
Exothermic process. A process in which energy is released from a system to the surroundings.
Extensive property. A property that depends on how much material is being considered.
Family. A group of elements whose symbols are in the same vertical column on the Periodic
Table.
First Law of Thermodynamics. A statement equivalent to the Law of Conservation of Energy,
that energy is neither created nor destroyed.
The more useful statement that the sum of the work done on a system and the heat gained by a system is the energy change of
the system is the form of the First Law used in chemical thermodynamics.
First-order reaction. A reaction whose rate is directly proportional to the concentration of one of
the reactants.
Fluid. A material that can flow and change its shape to fit the shape of its container.
Gases, liquids and supercritical fluids are fluids. While solids that are not very rigid like glass and many unreinforced plastics
flow very slowly, they are rarely considered fluids.
Fluorescence. The rapid emission of light ( a photon) by an atom or molecule after it has absorbed
a photon of a different energy.
The emitted photon is always of lower energy than the absorbed photon.
Formal charge. The difference between the number of valence electrons for an isolated neutral
atom and the number of electrons assigned to that atom in a Lewis structure, assuming that
electrons used to form covalent bonds are evenly divided between the atoms that are bonded..
Formal charge may also be defined for an atom by the equation form of the above definition.
Formal charge = # of Valence e for neutral atom - # non-bonded e - 1/2 e in bonds.
Formation. The process in a which a compound is formed. If no starting materials are stated, the
starting materials are assumed to be the constituent elements of the compound.
Formula unit. An atomic grouping that has the same composition in atoms as stated in its
chemical formula.
Formula mass (weight). The mass (weight) of 1 mole of formula units.
Free Energy. See Gibbs Free Energy change.
Frequency. The number of wave peaks or troughs that pass a point per unit time.
Fusion. The conversion of a solid to a liquid.
Gamma () Ray. A stream of high energy photons from a radioactive source.
Gas. A material that entirely fills its container.
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Gas Constant (R). The proportionality constant in the Ideal Gas Law,
PV = nRT
P is pressure, V is volume, n is the number of moles of gas and T is the Kelvin Temperature. The units of R depend on the
units of P and V.
Gibbs Free Energy Change (G). This is the maximum amount of work that can be obtained
from a process that occurs at constant temperature and pressure.
G can be calculated by the equation:
G = H - TS
Ground State. The lowest energy state available to a system.
Half-life. The elapsed time at which the concentration of a reactant has decreased to one half of its
original concentration.
In radioactive systems this is the time required for the number of radioactive nuclei to be reduced to half of the original
number present.
Halogen. One of the elements whose symbol is in Periodic Table Column 17 (VIIA) elements,
F, Cl, Br, I, At.
Hard Water. Water that contains a large concentration of metal ions with charges greater than +1.
The most common ions that make water hard are Ca2+ and Mg2+ . Concentrations higher than around 50 ppm tend to react
with normal soaps resulting in a scummy precipitate.
Heat. The transfer of energy between two systems at different temperatures.
Heat capacity. The amount of heat required to raise the temperature of a sample by one Celsius
degree.
Heterogeneous. Not uniform.
Heterogeneous mixture. A mixture in which different regions can be distinguished.
The presence of boundaries indicates a mixture is heterogeneous, and the scattering of light from boundaries between regions
is commonly used as a test for the presence of a heterogeneous mixture. The classification of all mixtures into heterogeneous
and homogenous depends, to some extent, on the method of observation.
Homogeneous mixture. A mixture which is uniform throughout.
See heterogeneous mixture above for the other main class of mixtures. The term solution is a synonym for homogeneous
mixture.
Hund's Rule. The empirical rule that in a system where different arrangements of electrons in the
orbitals of the same energy are possible, the electron configuration with the least number of paired
electrons is of lowest energy.
Hybrid orbitals. Atomic orbitals that are made by combining hydrogen-like orbitals on a central
atom.
Hybrid orbitals (hybrids) are grouped in sets of equivalently shaped orbitals with as many members as the number of
hydrogen-like orbitals that make them up. They are not individually symmetrical around the nucleus, but are directed toward
atoms that will be bonded. They are used as a model to be consistent with the observed geometry of molecules.
Hybridization. The process of combining hydrogen-like orbitals on an atom to make a new set of
hybrid orbitals.
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Hydrates. Compounds that contain a number of units recognizable as water molecules.
Hydration. The process in which a gaseous ion is converted into its normal form in aqueous
solution, i.e.,

or B(maq )
A (ng) or B(mg )  A (naq
)
This term is also used, especially in organic chemistry, to describe a reaction in which the net effect of a process is to add one
or more water molecules to another molecule.
Hydrogen bond. An intermolecular force that is used to describe the unexpectedly large
intermolecular attractions between molecules that contain H attached to F, N, or O.
Hydrolysis. The chemical reaction of a substance or a species with water to form at least two new
species..
This term is commonly used when an anion or cation reacts with water to change its pH.
Hydrophilic. A material that is easily dissolved or dispersed in water.
This is half of the hydrophilic/hydrophobic classification of materials.
Hydrophobic. A material that is not dissolved or dispersed by water.
This hydrophilic/hydrophobic classification is commonly used in the biological sciences and is interpreted as classifying
biological materials by whether their largest intermolecular forces are of the dipolar, hydrogen-bonding, and ionic types or of
the London dispersion force type.
Hygroscopic. A material that can absorb water vapor from the atmosphere.
Hypothesis. An explanation for one or more observations.
An hypothesis is most easily understood to be an answer to the question "Why. . .?"
Ideal gas. Any material that is accurately described by the Ideal Gas Law .
If a substance behaves as an ideal gas, it can be viewed microscopically as a simple set of point masses with an average kinetic
energy determined by the temperature of the system.
Ideal Gas Law. PV = nRT, where P is the pressure the gas exerts, V is the volume of the
container, n is the number of moles of the gas, and T is the absolute temperature of the gas. R is a
constant independent of the identity of the gas, but dependent on the units of P and V.
Ideal solution. A solution whose components obey Raoult's Law.
Induced dipole. The separation of positive and negative charge in a neutral atom or molecule
caused by a neighboring ion or dipolar molecule.
Inhibitor. A material that decreases the rate of reaction.
This term may be used in the sense of a "negative catalyst" that slows a reaction without being chemically changed, but it is
more general than that.
Inorganic compound. Any compound that does not contain carbon or carbonates or carbonate
derivative.
This is one of the two exclusive classes of compounds, inorganic/organic. The dividing line is not perfectly clear, since some
compounds containing carbon bonded only to oxygen or other carbon atoms may be classified as either.
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Insoluble. Having a solubility less than 0.01 M.
Insoluble is a qualitative term that is arbitrarily defined above. Sometimes the term "sparingly soluble" is used as a synonym.
See soluble for further comments.
Insulator. A material that has a high resistance to the flow of electricity or heat.
Intensive property. A characteristic of a material that does not depend on the amount of sample.
This is half of the classification of properties as either intensive or extensive. Any pair of extensive properties can be made
intensive by taking their ratio.
Intermediate. A molecule that occurs in a mechanism but is not present in the overall chemical
equation for the process.
Intermolecular forces. The forces that exist among separated microscopic particles.
This term is defined above rather generally. A more common definition is: the forces that exist between separated atoms or
molecules.
Intramolecular forces. The forces that exist between the particles in a formula unit.
This term is often defined more specifically as the "attractive forces between the atoms in a molecule.
Inversely proportional. A relationship that follows the equation y = m(1/x).
Ion. A electrically charged atom or molecule.
Ionic bond. The electrostatic attraction between oppositely charged ions that holds them together
in compounds.
Ionic compound. Any compound that is viewed as containing ions.
Any compound that contains both a metal and a non-metal will be viewed as an ionic compound.
Ionic radius. The radius of an ion.
This is found in tables which represent average values in a range of compounds. In fact, an accurate ionic radius is an
experimental value found from a crystal structure determination.
Ionization energy (first). The amount of energy absorbed when one mole of gaseous atoms loses a
mole of electrons to become a mole of gaseous positive ions.
Alternately, Ionization energy is the standard enthalpy of reaction for the process:
A(g)  A (g ) + e , H orxn = Ionization Energy
This is a definition of first ionization energy (sometimes called ionization potential). In fact, similar definitions an be made
for second, third, etc. ionization energies. In each case, it is the energy absorbed when a mole of particles loses a mole of
electrons to produce a mole of gaseous particles a single unit more positively charged than the original gaseous particles.
Isoelectronic. Having the same number of electrons and thus the same ground state electron
configurations.
Sometimes this adjective is used to describe atoms or ions that have the same valence electron configurations disregarding
principal quantum number.
Isolated system. A system that exchanges neither energy nor mass with its surroundings.Isotopes.
Atoms or nuclei of the same element containing different numbers of neutrons.
Joule. The metric unit of energy, 1 J = 1 newton * m
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It requires around 1 J of work to lift a glass of water 40 cm. When heating, 1 J seems much smaller since heating that same
glass of water by 1oC requires around 1000 J.
Kelvin temperature scale. see Absolute temperature scale.
Kernel. The nucleus of an atom and its electrons that are not valence electrons.
Kinetic energy. K.E. = 1/2 * m * v2
Kinetic energy is the energy a particle has because of its motion. It is one of the two kinds of energy a particle can have,
potential energy being the other kind. The kinetic energy of a system is usually found as the sum of the kinetic energies of its
individual particles.
Lanthanide. One of the elements with atomic number 57  71, La  Lu.
These elements or their ions have partially filled 4f energy sublevels.
Lattice energy. The standard enthalpy of reaction for the following process:
MaXb (s)  M(mg ) + b X (xg)
L.E. = H orxn
Lattice points. The positions of the points that make the repetitive lattice that defines a crystalline
solid.
This term is sometimes used to describe the positions of atoms or groups within the repeating unit that is used to form the
lattice.
Law. A simple verbal or mathematical statement of the relation that is always observed between
phenomena.
Law of conservation of atoms. The number of atoms of every element present in a chemical
process is not changed during the process.
This is a modification of the Law of conservation of mass for chemical systems. In practice, it is used in the process of
balancing skeletons to make chemical equations by making sure the number of symbols for each element present are the same
in number on both sides of the arrow.
Law of conservation of energy. The total amount of energy in the universe is fixed.
This is more commonly stated that the total of the energy changes for a system and its surroundings is zero. This is not true if
mass is converted to energy as it may be in nuclear reactions.
Law of conservation of mass. The total amount of mass in the universe is fixed.
This is modified to the Law of conservation of atoms in chemical reactions. See the Law of conservation of energy for the
exception.
Le Chatelier's Principle. Any change in a system intensive variable (stress) for a system at
equilibrium, results in a shift of the position of the equilibrium so that the effect of the change is
decreased.
Change in intensive variables (stress) means either a change in the concentrations or pressures of either reactants or products,
or in the laboratory conditions of temperature or pressure.
Lewis acid. A species that can accept an electron pair to form a covalent bond.
Lewis base. A species that can donate an electron pair to form a covalent bond.
The Lewis definitions of acids and bases are used especially in organic chemistry and in non-aqueous systems.
Lewis dot symbols. The representation of an atom or a monatomic ion that uses the atomic symbol
with a dot to represent each of its valence electron pairs or a line for each pair of valence electrons.
Sometimes a single line is drawn to represent a pair of electrons.
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Lewis structure. A representation of the bonding in a molecule using Lewis dot symbols. Shared
electron pairs between two atoms are shown as lines and the number of these lines is taken as the
number of bonds between those two atoms.
Lewis structures are also called Lewis formulas or electron dot structures. These must satisfy the set of rules. It is possible
that what looks like a Lewis structure is not complete because it makes peripheral atoms appear non-equivalent. In this case,
what is shown is a resonance hybrid in which case the Lewis structure for the molecules is the complete set of resonance
hybrids that make non-equivalent peripheral atoms appear equivalent.
Ligand. A molecule or ion bonded to a central metal ion in a complex ion.
Ligands are invariably Lewis bases and the resulting bonding can be considered as the result of the ligand acting as a Lewis
base donating the metal ion that is a Lewis acid.
Limiting reagent. The reactant that would be entirely consumed were the reaction to proceed to
completion according to the chemical equation written to represent it.
There may be more than one limiting reagent (or limiting reactant) if an exact stoichiometric ratio of reagents is mixed. Other
reactants are said to be excess reagents.
Liquid. A material whose volume is fixed but whose shape is determined by the shape of its
container.
Liter. The volume of a cube that is 1 dm on a side, i.e., 1 L = 1 dm3 .
Convenient equivalence equations are: 1000 L = 1 m3 ; 1 cm3 = 1 mL.
London dispersion forces. The intermolecular attractions between all molecules that result from
the spontaneous formation of a temporary dipole in one molecule inducing a corresponding dipole
in neighboring molecules.
These are also called London forces or dispersion forces. Occasionally they are called Van der Waals forces though they are
only one of this class of forces. They are the result of the spontaneous polarization of a molecule as a result of the random
distributions of electrons. These forces occur in all molecules and are fairly characterized as having a magnitude directly
dependent on the number of electrons the molecule contains.
Lone pairs. Electron pairs in a Lewis structure that are not used for bonding.
These are also called non-bonding pairs.
Macroscopic. A property of a system that can be observed directly.
This word is commonly used as an adjective to describe anything that is on a laboratory size scale, can be observed visually,
or can be measured using normal laboratory equipment.
Mass. The amount of matter in an object as measured by its resistance to motion.
Mass is frequently confused with weight by chemists, and we will frequently use the two terms interchangeably though they
have different meanings. See the description of weight for this distinction.
Mass number. The total number of protons and neutrons in a nucleus.
Material. A more modern term for any sample of matter.
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Matter. Anything that has mass or occupies space.
The two choices are simply two alternative methods to test whether the object is matter. Either test is sufficient by itself.
Mechanism. A detailed sequence of steps that show the progress of a chemical reaction in passing
from reactants to products.
Melting point. The temperature at which solid and liquid forms of a pure substance can both exist
in equilibrium.
This is the same as the freezing point for a pure substance. When this term is applied to a mixture it is ambiguous. Instead,
the term melting range should be used.
Melting range. The range of temperatures between which solution and its solid can both exist at
equilibrium.
This term is used experimentally to mean the temperature at which the first sign of liquid is formed to the temperature at
which the mixture forms a homogeneous liquid.
Metals. 1. The elements whose symbols are to the left of the diagonal zig-zag line on the Periodic
Table, including those elements at the bottom of the table.
2. Experimentally, metals are those materials that have a shiny metallic sheen and are good
electrical and thermal conductors.
Metalloids. Elements with properties intermediate between those of metals and non-metals.
This class of elements is found near the diagonal line of the periodic table and whether an element belongs to this class may be
depend on the allotrope being considered. Normally, B, Si, Ge, As, Sb, Te, Po, and At are placed in this class, but allotropes
of C, P, Se, and Sn might be included.
Metallurgy. The science AND technology of metals and their ores.
Microscopic. At the scale of individual atoms, ions, and molecules.
This is one of the two scales used in chemistry, the macroscopic and microscopic. The microscopic view is a conceptual point
of view that is useful for visualization of materials and their changes but it is not normally directly observable (except possibly
as inferred from very specialized instrumentation). This is a very different definition than that a biologist would use.
Mineral. A naturally occurring substance.
Normally, this term is used in the wider sense of one of a range of naturally occurring substances which have a limited range
of possible chemical compositions.
Miscible. Two substances are miscible if they are soluble in each other at any ratio.
Mixture. A combination of more than one substance.
The main classification of mixtures is into homogeneous and heterogeneous.
Molar heat of fusion. The same as enthalpy of fusion.
Molar heat of vaporization. The same as enthalpy of vaporization.
Molar concentration. The same as Molarity.
Molar mass. The mass of 1 mole of formula units.
Molar mass and formula mass are used interchangeably. Some people even use formula weight the same way as molar mass.
Molar solubility. The molarity of a saturated solution of a substance.
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Molarity. The concentration of a solution expressed as the ratio of moles of solute per liter of
solution.
M = moles solute / liter solution
Mole. The counting unit that is the same size as the number of atoms in exactly 12 g of 12C .
Mole fraction. The ratio of the number of moles of a substance in a system to the total number of
moles of all substances present in the system.
Mole fraction of A, XA = nA / nA + nB + nC + . . .
Molecular formula. A chemical formula in which the subscripts are the actual numbers of atoms
whose symbol precedes the subscript in a single molecule of the substance.
Molecular mass. The mass of a molecule of a substance in amu.
This term is sometimes used as being identical to formula weight but this is incorrect. Numerically, disregarding units, the
molar mass is equal to the formula weight of a molecular substance.
Molecular orbital. An orbital extending over more than one nucleus.
Molecularity. The number of species on the left of the arrow in a mechanistic step.
Molecule. Two or more atoms chemically combined into a unit that is further from neighboring
molecules than the bonded atoms in the molecule are from each other.
Monatomic ion. An ion that contains only one nucleus.
Monomer. A single species that can be chemically combined with others like it to form a larger
molecule containing more than one of these species.
Multiple bonds. Covalent bonds that contain more than a single pair of electrons.
Net ionic equation. A chemical equation that shows the ions reacting and produced in a reaction,
with no spectator ions.
Neutral. 1. An object having no charge.
2. A solution having a pH of 7.
Neutron. A particle that has no electrical charge and a mass of 1. amu.
Neutrons in an atom are found in the nucleus of an atom and have a mass of 1.67x10 -24 g.
Newton. The SI unit of force, 1 kg * m/ s2.
The force (weight) exerted by a 1 kg mass is around 10 N on the earth's surface.
Noble gas. One of the elements whose symbols are in column 18 (VIIIA) of the Periodic Table,
He, Ne, Ar, Kr, Xe, or Rn.
Nonbonding electrons. Valence electrons that are not used to form covalent bonds in a molecule.
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Nonelectrolyte. A substance that does not produce electrical conductivity in its aqueous solution.
Nonmetal. An element whose symbol is to the right and above the diagonal zig-zag line on the
Periodic Table.
These elements are not good conductors of electricity or heat.
Nonpolar molecule. A molecule that does not have a dipole moment.
This term is applied to any molecule with a zero or a very small dipole moment.
Nonspontaneous. Any process that cannot occur without any outside influence.
Chemists call a chemical reaction that produces essentially no products compared to the amount of reactants a nonspontaneous
reaction.
Nonvolatile. A material that has no measurable vapor pressure.
Normal boiling point. The boiling point of a substance at 1 atm external pressure.
When the term boiling point is used without specifying pressure, the normal boiling point is usually meant.
Normal state. The state of a substance that is of lowest energy at the stated conditions.
This is used in thermodynamics to refer to the form of a substance that is normally observed at 298K, 1 atm, and which has a
G of = 0.
Nuclear binding energy. The energy absorbed when a nucleus is separated into its individual
protons and neutrons.
Nuclear chain reaction. A spontaneous branching sequence of nuclear fission reactions.
In these branching sequences, a single fission event initiates more than one other fission event so the reaction accelerates as it
proceeds.
Nuclear fission. A nuclear reaction in which a heavy nucleus divides to form smaller nuclei.
Usually, two unequal smaller fragments and some neutrons are the products of a nuclear fission. The heavy nucleus has a
mass number greater than 200 in a natural spontaneous fission.
Nuclear fusion. The process in which two nuclei combine to form a larger nucleus.
Nuclear reaction. Any process in which a nucleus changes.
Nucleus. The central massive part of the atom containing its neutrons and protons.
The nucleus has essentially no volume and contains almost all of the atomic mass.
Nuclide. Any atom or nucleus with a specific mass and proton number.
Octet rule. The Lewis structure rule that atoms C -> Ne will be surrounded by four pairs of
electrons in stable compounds, and that the heavier non-metals prefer to be surrounded by a
minimum of four pairs of electrons.
Orbital. Either an atomic orbital or a molecular orbital.
Reminder: An orbital is simply the space occupied by an electron or pair; alternatively, an orbital is a map of that space.
Order. The exponent of a concentration in a rate law.
While the order of a reaction may be occasionally be used to describe the sum of all the exponents in a rate law, the more
common usage is to speak of the order with respect to some reagent. It is this latter common usage that is covered by the
definition above.
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Ore. The material in a natural deposit from which a substance of economic value can be obtained.
Organic chemistry. The field of chemistry that specializes in the chemistry of compounds
containing carbon.
See inorganic chemistry for a minor refinement of this definition. The specialized study of the chemistry of living organisms
is a different field called biochemistry, or it may be called physiology, etc. While these fields are not mutually exclusive,
organic chemistry tends to exclude the materials that are encountered under the term biochemistry.
Organic compounds. Compounds that contain carbon.
Osmosis. The movement of species through a membrane permeable to them, from the side of
higher concentration to the side of lower concentration of that species.
Osmotic pressure. The pressure required on the low concentration side of an osmotic pressure
experiment to just stop osmosis.
Oxidation number. A number assigned to an atom according to the oxidation number rules.
This is the charge on a monatomic ion and can be understood, in a molecule or molecular ion, as the charge an atom would
have if all the electrons in bonds were entirely transferred to the more electronegative atom in the bond.
Oxidation state. Same as oxidation number.
Oxidizing agent. A species that can accept electrons.
Oxyacid. A compound containing hydrogen bonded to oxygen bonded to a central atom that is
possibly bonded to other oxygen atoms.
Paramagnetic. A material that is attracted to a magnet.
Paramagnetic substances contain at least one unpaired electron, and these unpaired electrons act independently of each other.
Partial pressure. The pressure exerted by one component in a mixture of gases.
Pascal. A pressure of 1 N / m2 .
The size of a pascal can be better understood by giving the definition of an atmosphere:
1 atm = 101,325 pa
Pauli exclusion principle. No two electrons can have an identical set of quantum numbers.
If two electrons were to have the same quantum numbers, then they would be the same electron since electrons have no other
distinguishing properties.
Percent. Percent = 100% * Quantity of interest/ Total Quantity
Percent composition. Percent X = 100% * mass X / mass sample.
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Percent yield. Percent yield = 100% * actual amount product isolated / theoretical yield.
Period. An entire horizontal row of the Periodic Table.
Periodic Table. An arrangement of the symbols for the elements into tabular order according to
increasing atomic number, so that elements of similar chemical properties are arranged in vertical
columns.
pH. pH = -log[H+]
This is also written as : pH = -log[H3O+] since H+ and H3O+ are both ways to write the hydrated form of a proton.
Phase. A homogeneous region in a heterogeneous system.
The term phase has also been used to describe a pure substance as s, l, or g because of its usefulness in describing a melting or
boiling mixture of a pure substance; however, the term state is more widely used in this context.
Phase change. The process of changing a substance from one phase to another.
Phase diagram. A diagram that shows the conditions at which the phases of a material exist.
For a pure substance, the conditions normally plotted are external pressure vertically and temperature horizontally. The lines
separating regions on these diagrams are the equilibrium lines where the two adjoining phases can coexist.
Photon. A particle of light.
Physical property. Any property that can be measured without changing a substance into another
substance and that can be measured independent of the materials used in the measurement.
This is one half of the classification scheme of properties as chemical/physical.
Plasma. A state of matter containing gaseous ions.
Polar molecule. A molecule that has a dipole moment.
Polarizability. The ability of a species to have its electron distribution distorted from the ground
state distribution by electric fields.
Polyatomic ion/molecule. An ion/molecule that contains more than two atoms.
Polymer. A material whose molecules consist of large numbers of identical simple units
(monomers) joined together by covalent bonds.
Position of equilibrium. The conditions and concentrations or partial pressures of all species in
an equilibrium system.
Potential energy. The energy a particle has because of its position.
This is one half of the classification of particle energies by position (potential) and by velocity (kinetic). The changes
occurring in chemical systems because of bond breaking and bond formation are changes in potential energy.
Precipitate. An insoluble solid that is formed when solutions are mixed.
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Precision. The numerical agreement between a set of measurements of the same quantity.
"High" precision means the deviation between measurements is very small.
Pressure. The force per unit area. Pressure = Force / Area.
Product. A substance formed from a chemical reaction.
In a chemical equation, each of the substances to the right of the arrow is a product.
Property. An observable characteristic.
Proton. A nuclear particle with mass around 1 amu and a single positive atomic charge.
The charge of a proton is +1.60x10-19 C and a mass of 1.67 x10-24 g in metric units.
Qualitative. A non-numerical property of a system.
Quantitative. A numerical property of a system.
Quantized system. A microscopic system whose possible energy states are at discrete and
separated energies.
The adjective quantized is used to describe any system with a limited number of discrete possible energies.
Quantum. The smallest quantity of energy a system can absorb or emit.
Quantum numbers. One or more integral or half-integral numbers that are used to describe the
state of a microscopic system. These numbers are the result of a solution of the quantum
mechanical equation that describes the system.
Radiation. The emission or direct transmission of energy through space.
Radical. Any atom or molecule containing an unpaired electron.
Radioactivity. The emission of particles and/or photons from nuclear reactions.
Raoult's Law. The partial pressure of a solution component is the mole fraction of that component
times the vapor pressure of that pure component. In symbols,
PA = XA * PAo
PA is the vapor pressure of component A in the mixture, XA is the mole fraction of A in the
solution, and PAo is the vapor pressure of pure component A.
Raoult's law is an idealization and only works well in solutions where the intermolecular forces of solutes and solvent are of
the same type and strength.
Rare earths. Another name for lanthanides.
Rate of reaction. The slope of the line on a graph of product concentration as a function of time,
Rate = dc/dt
The rate is commonly approximated as : Rate =  c /  t, where  c is the change in product concentration over the time
interval  t. The expression for rate is somewhat different if stoichiometric coefficients are not all 1, and the sign is changed if
reactant concentrations are followed.
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Rate constant. The proportionality constant in rate laws between the rate of reaction and the
concentrations of reactants raised to their orders.
Rate = k * Aa * Bb
As defined above, the rate constant should be called the specific rate constant since it is the rate when the concentrations of
reactants are 1 M . However, in the generalized definition of rate law below, rate constants will be any numbers occurring in
the function that are not dependent on concentrations of reactants or catalyst.
Rate law. The functional dependence between the rate of reaction and the concentrations of
reactants, catalysts and any other species that effect that rate.
Rate-determining step. The slowest step in the mechanism of a reaction.
In many situations, depending on concentrations, more than one step may limit the rate of reaction, and this makes the ratedetermining step viewpoint a limiting viewpoint.
Reactants. The starting substances that are put into a chemical reaction.
The species on the left of the arrow in the chemical equation that describes that reaction.
Reaction. Same as chemical reaction.
Reaction mechanism. The sequence of elementary steps that leads from the reactants in a reaction
to the products.
Reaction quotient (Q). A number equal to the ratio of the product concentrations raised to their
coefficient powers divided by the reactant concentrations raised to their coefficient powers,
( C) c ( D )d
Q=
;
( A )a ( B ) b
Q is the reaction quotient and ( ) are the quantities of materials present at any time in the generic
equation:
aA + bB  cC + dD
Q has the same value as Keq if the system is at equilibrium; however, Q is defined for any mixture of reactants and products.
Reaction rate. Same as rate of reaction.
Redox reaction. An electron transfer reaction.
Electron transfer reactions are also called oxidation-reduction reactions.
Reducing agent. A species that can donate electrons in a redox reaction.
Representative element. One of the elements whose symbols is found in the Periodic Table
columns 1, 2, and
13  18 ( IA  VIIIA).
These are the elements that are not transition elements and which have only s and p electrons in their valence electron
configurations.
Resonance. The use of two or more Lewis-type formulas to represent molecules.
These formulas are called resonance structures and are used when the formula drawn fits the Lewis rules adequately but
makes peripheral atoms non-equivalent. See Lewis structures for further comments.
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Resonance structures. Two or more Lewis-type formulas that satisfy the Lewis rules but cannot
be correct because they make peripheral atoms with identical nuclei non-equivalent.
These structures with non-equivalent peripheral atoms are not correct individually because they do not agree with
experimental structure determinations.
Reverse osmosis. The water purification process in which impure water under high pressure is
placed against a membrane permeable to water only. This results in a lower concentration of
contaminants on the other side of the membrane.
Reversible reaction. A reaction that can occur in either the forward or reverse direction compared
to that shown in the chemical equation for that reaction.
All chemical reactions are reversible, but some are spontaneously so complete or incomplete that reversing the reaction to the
extent normally possible still has no appreciable effect on concentrations. These reactions are often called irreversible though,
strictly speaking, they are no different than any other chemical reaction.
Salt. Any compound that contains a metal and a nonmetal.
Commonly, salts with hydroxide or oxide anions are classified separately as bases and excluded from the class of salts.
Saturated solution. A solution that has a high enough concentration of solute so that can exist in
equilibrium with pure solute.
Science. The study of anything using scientific method.
Scientific method. The logical cycle of observation, hypothesis, hypothesis testing and hypothesis
reformulation combined with the generalization of hypotheses in laws or theories.
Second Law of Thermodynamics. The generalization that the entropy of the universe increases in
any spontaneous process.
This is just one of a large number of alternative statements of this idea.
Semiconductivity. The property of material that it is a very poor conductor of electricity at low
temperature but has a conductivity that increases exponentially with temperature.
Semimetal. Same as metalloid.
Semipermeable membrane. A thin sample of a material that allows the flow of one or more
species through it.
Significant figures. The digits in a numerical measurement or calculation that are known with
assurance plus one last digit that has a limited uncertainty.
Skeleton. A list of reactants in a chemical reaction written on the left of an arrow separated by "+"
signs followed by a list of the products on the right of the arrow separated by "+" signs.
A skeleton contains an accurate listing of reactants and products but it does not contain coefficients.
Solid. Any material that is rigid enough to maintain its shape.
In chemistry, the term solid is more commonly used as a short form of the term "crystalline solid".
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Soluble. A substance that has a solubility greater than 0.01 M .
Basically, this term is qualitative and the dividing line between soluble and insoluble substances is arbitrary depending on who
defines the term. It may also be applied to materials if a mass solubility is used to set the soluble/insoluble dividing line.
Solubility. The concentration of a saturated solution of a specific solute.
Solubility product. The equilibrium constant for the process of dissolving a sparingly soluble salt
in water to form its separated aqueous ions.
Solute. One of the substance(s) present in smaller quantity in a solution.
Solution. Same as a homogeneous mixture.
Most people (chemists included) think of liquid solutions when the word solution is used by itself, but solid and gaseous
solutions are equally possible.
Solvent. The major component in a solution.
Specific heat. The amount of energy absorbed when 1 g of a material is heated enough to raise its
temperature by 1 K (1oC).
Spectator ions. Ions that are present in an aqueous solution but are not affected by the chemical
reaction being considered.
These are the ions that are omitted in net ionic equations since they do not change chemically.
Spectrum. 1. A graph of intensity of light emitted or absorbed vs. the wavelength or frequency of
light studied.
2. A direct presentation of the light from a sample dispersed by wavelength or frequency.
3. The appearance of white light as indicated under def. 2.
Spontaneous. A description of any change that can occur with no change in conditions and without
the influence of any force.
A more chemical definition is that a chemical reaction that can produce a large amount of products compared to the amounts
of reactants.
Stable. At the lowest potential energy accessible to the system.
This adjective is used in many other ways, and it pays to ask how it is being used. It may mean also:
1. Not able to react rapidly under normal conditions.
2. Not chemically reactive with the normal components of the atmosphere.
3. Not as reactive chemically as some other system being compared.
Other descriptions of similar import may also be intended.
Standard enthalpy of formation ( H of ). The standard enthalpy of reaction when one mole of a
compound or species is formed from its elements in their stable forms at the standard conditions.
Standard enthalpy of reaction ( Ho
). The enthalpy change occurring when one mole of
rxn
product is formed under the standard conditions of 1 atm pressure, 1 M concentration at
298.15 K.
Other temperatures may be specified.
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Standard temperature and pressure (STP). 0oC, 1 atm.
State function. A property of the system that is determined by the state of the system.
This is to be contrasted to a property that depends on the history of the system, that is, on the details of how the system got to
where it is.
State of a system. The values of all the macroscopic variables of a system, composition, pressure,
temperature, volume, and phases present.
Step. One of the elementary processes in a mechanism.
Stoichiometric amounts. The exact amounts of reactants so that when a reaction occurs to
completion, according to the chemical equation written, the reactants are all exactly consumed.
Stoichiometry. The relationships between the amounts of reactants and products in chemical
reactions.
Structural formula. A chemical formula that is written in such a way as to provide details about
which specific atoms are bonded to each other.
In organic chemistry, structural formulas commonly provide enough information so that all the bonded atoms in a molecule
are mapped out.
Structural isomers. Molecules that have the same molecular formulas but different structures.
Sublimation. The process in which a solid is converted directly to a gas.
Substance.
Macroscopic definition: A homogeneous material with a sharp melting point and boiling point.
Microscopic definition: A material consisting of a single repeating formula unit.
Intuitive definition: An element or compound.
Currently, substance and pure substance mean the same thing as defined above. However, you must be wary of the older
usage, particularly before this century, when substance and material were considered as synonyms.
Supercooling. The process of cooling a liquid below its freezing point without the formation of a
solid, or of cooling a solution below its saturation temperature without solid formation.
Supersaturated solution. A solution with a solute concentration larger than that of a saturated
solution at the same temperature.
Supersaturated solutions are most commonly prepared by rapidly cooling a saturated solution or by mixing solutions which
can form a less soluble solute when they are combined..
Surface tension. The amount of energy required to increase the surface of a liquid by a unit in
area.
Surroundings. The remainder of the universe outside the system.
System. Any specific part of the universe that is being studied.
Together, system and surroundings divide the universe into two parts, what of interest (system) and the rest (surroundings).
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Technology. The use of results from science to manipulate materials.
Generally, the outcome of technology is a product that can be sold.
Temperature.
Microscopic definition: The average kinetic energy per particle in the system.
Experimental definition: The reading of a device called a thermometer that has some property that
changes as heat is added to it.
Typically, volume, length, voltage, resistance, or infrared emission spectra are properties used in thermometers.
Theoretical yield. The amount of product predicted to form by a chemical equation if the limiting
reagent were to react completely.
Theory. A general principle that explains many observations with their resultant Laws.
Thermal energy. The energy a system has because of the total kinetic energy of its constituent
particles.
Thermochemistry. The study of the energy changes occurring in chemical reactions.
Third law of thermodynamics. The entropy of a perfect crystalline substance is zero at 0 K.
Transition metals. The elements whose symbols are in columns 3  12 (IIIB  IIB) of the
Periodic Table.
These elements have atoms or ions that have partially filled d orbitals. This term is sometimes used to denote all the elements
that are not representative elements.
Triple point. The temperature and pressure at which solid, liquid and gaseous phases of a
substance can be in equilibrium with each other.
On a phase diagram, this is the point where the l-s, l-g, and s-g equilibrium lines all intersect.
Unimolecular step. A process in which only one reacting species appears on the left of the arrow.
Unit cell. The repeating unit in a crystalline solid.
Valence. When used as a noun, the term is archaic and ambiguous. However see valence
electrons.
Chemists have abandoned the term valence because it had become too vague in its definition, meaning either oxidation
number, number of bonds an atom could form, number of bonds an atom had in a particular compound, coordination number,
or charge.
Valence electron configuration. An electron configuration that lists the orbital occupation of the
valence electrons only.
Valence electrons. The outer electrons of an atom. These electrons are those of occupying the
highest principle quantum number energy levels plus any electrons in partially filled lower quantum
number sublevels.
Valence shell electron pair repulsion (VSEPR) model. The model that predicts the shape of
molecules by assuming electron pairs on a central atom act as vectors pointed as far apart as
possible due to electron pair repulsions.
117
Van der Waals forces. The intermolecular forces that act between neutral molecules due to
dipolar attractions and London dispersion forces.
Van der Waals radius. The sum of the Van der Waals radii for two non-bonded atoms that are in
contact is the distance between them.
Vapor pressure. The partial pressure of the gaseous form of a substance in equilibrium with its
liquid phase.
Vaporization. The process of converting a liquid to a gas.
Viscosity. The resistance of a fluid to flow.
Volatile. Has a measurable vapor pressure.
The general usage of this term is that it means evaporates rapidly. This is, of course, a consequence of the fact that a volatile
liquid must have a reasonably large vapor pressure. This is clearly a very qualitative term.
Wave. A repetitive phenomenon.
Wavelength. The distance between successive waves.
Weight. The force an object exerts because of its interaction with a gravitational field.
X-ray diffraction. The scattering of X-rays by passing through a solid.
X-ray structure determinations, our main source of structural data on compounds, are done by the measurement and analysis
of X-ray diffraction measurements.
Yield of a reaction. Same as actual yield.
AM 8/00
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Supplementary Materials Table of Contents
Supplement 0
Supplement 1
Supplement 2
Supplement 3
Supplement 4
Supplement 5
Supplement 6
Supplement 7
Supplement 8
Supplement 9
Supplement 10
Supplement 11
Supplement 12
Supplement 13
Supplement 14
Supplement 15
Supplement 16
Supplement 17
Supplement 18
Supplement 19
Supplement 20
Supplement 21
Supplement 22
Supplement 23
Supplement 24
Supplement 25
Supplement 26
Supplement 27
Supplement 28
Supplement 29
Supplement 30
Supplement 31
Supplement 32
Supplement 33
Supplement 34
Supplement 35
Supplement 36
Scientific Method
Reading Technical Material
Classification of Matter
Solids, Liquids & Gases
Submicroscopic View of Matter
Submicroscopic Particles & Prop. of s, l, & g
Chemical Formulae
Metric System & Factor Label, part 1
Accuracy and Precision
Factor Label Problem Solving, Area and Volume
Elem. Atomic Structure & the Periodic Table
Light & the Bohr Model of the Atom
Finding # Valence e- & Lewis Symbols of Elements
Predicting Ion Charges & Naming Ions
Nomenclature of Ionic Compounds
Lewis Structures
Polarity of Molecules
Nomenclature of Covalent Compounds
Chemical Equations & Limiting Reagent
Polarity and Solubility of Substances
Osmosis
Outline of Acids and Bases
Buffers
Predicting Spontaneity of a Reaction
Risk Management
Concentration in %, Parts per Thousand, & ppm
Chemical Equilibrium
Gases - P, V, T, & Amount
Determining Molecular Geometry from Lewis Structure
Concentration Units, Molarity, and Dilutions
The Mole
Mass Percent (Composition)
Combustion Analysis
Chemical Equations
Limiting Reactants
Some Solution Chemistry
Solution Stoichiometry – Titrations
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120
121
123
126
127
131
133
134
138
139
140
142
144
145
148
151
153
156
159
162
164
165
167
168
169
170
172
174
176
179
184
186
187
189
191
192
195
Supplement 0: Scientific Method Supplement
The term “scientific method” is a name for the approach you can use for investigations. The
first step generally involves observation--seeing what has happened. You may then wonder why the
thing happened. You then use information that you know (or can find) to think of a possible
explanation, which is called a hypothesis. Generally you then test your hypothesis, often many
times.
Be careful to distinguish facts from interpretations. For example, if someone got sick after
eating at George’s restaurant, the fact is that they got sick. One interpretation is that the meal made
them sick. Other interpretations may be possible.
Your tests must fit the hypothesis and have certain characteristics or you may reach false
conclusions. Be careful that your test is not too narrow (too specific) for your hypothesis, or else
narrow your hypothesis. A more general hypothesis should be tested under many different
conditions before it is considered true. To run a test, change only one thing and see what affect that
change has. Changing only one thing at a time to test its effect is called control of variables. If you
change more than one thing, you will not know which change causes any effect that you might see.
Also, a test should be run several times. Sometimes small changes which you don’t notice will
affect results,or the result is a small change hard to see. Multiple trials help avoid this problem.
Whenever possible a control should be run too. A control run is doing the same procedure without
changing whatever you are testing. This helps make sure the effect is really due to what you are
testing rather than something else about the test.
As an example, consider the hypothesis that Cheer is a good detergent for removing stains
from clothing. The hypothesis implies that Cheer is good on any stains and any clothing, so if you
tested the Cheer only on mustard stains on cotton, your test would be too narrow for the hypothesis.
You should either try a variety of staining materials and types of cloth or say that Cheer is good for
removing mustard stains from cotton. Multiple trials would make your results more certain.
Varying the conditions would help make a more general statement about how good the Cheer really
is--is it good in all temperatures and in hard or soft water? You can’t say until you test with those
conditions. A control in this detergent test would be washing without any detergent--you should be
sure that the “stain” won’t came out even if detergent isn’t used.
There are some other terms that scientists use when doing and reporting investigations. The
terms law, theory and model are frequently used. A law is a generalization of something which has
been observed many times under many situations. An example of this is the law of conservation of
energy (energy is neither created nor destroyed). Over the course of many, many years, scientists
have not observed energy being created from nothing nor disappearing. A theory is an extremely
well-tested hypothesis. It explains many different types of observations. For example, the theory
that matter consists of tiny particles called atoms explains many of the things we see.
One other term that we should review is the word model. In science a model is a simpler
system or way of thinking about something. Though it is simpler, it has the same properties as the
ones that we are particularly interested in explaining. Thus, a model is used to examine a few
important aspects of a more complex issue or object. It helps us to zero in on the things we think of
most concern at that time and helps us to ignore other things. The items that we ignore may be
different in the model as long as they do not influence the characteristics that we are examining.
For instance, a model car is simpler than a real car, but it generally has 4 wheels. Sometimes a
model is used to suggest that something behaves in a way similar to something more familiar. For
example, using a typewriter could be used as a model for using a computer keyboard.
The usefulness of a model depends on how well suited it is to the task. A model car might
be useful in explaining to a non-english speaker what the term “car” means, but it would not be
helpful if you need to go to the movies!
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Supplement 1: Reading Technical Material
Reading a science textbook is quite different than reading a novel for pleasure. In a science
book there is more to remember from a reading session, and there are often more new words than
you would encounter in a similar length novel. The concepts themselves may be more difficult to
understand than those of a novel. Approaching the reading of science material differently than the
way you approach the reading of a novel is likely to be more productive.
There are several different approaches to reading technical material such as science
textbooks, but they tend to have many common features. The one we will use in this class is a
modification of one of these methods, and can be called PQRRST for short. The steps involved in
PQRRST are listed and described below.
PQRRST: Preview--Question--Read--Reflect--Self-recite--Test (Review)
Preview the Material--Turn through the chapter pages to find out what the chapter is about
and where it’s going. Examine the title, introductory paragraphs, headings, graphic aids, terms in
italics or bold-face type, and the chapter summary. Read through any assigned homework
questions, noting unfamiliar terms. Doing all this will help you to decide what you need to know
from the chapter. As you do these things, ask yourself “What do I already know about this?”.
Connecting new knowledge with things that you already know will help you to remember the new
stuff better.
Question--Turn headings into questions. Look for topic sentences in paragraphs and turn
them into questions. Ask yourself “What questions might my teacher make about this material?”
All of these techniques help you to concentrate on what is important in the material.
Look for “hidden” questions, too. For example, a passage might say something such as “As
elements, hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine and iodine occur in molecules
that contain two of the same type of atoms per molecule.” A question that might occur at that point
is “How do other elements occur?”
Read--Do your in-depth reading, working on one section intently before starting the next.
Match what you are reading to the questions that you wrote earlier. Look carefully at new terms
and examine the graphic aids carefully; see if their meanings are clear to you. If you have reread
the material and know the definitions of all the terms and still don’t feel sure about the material,
place a mark in the margin of the book so that you can ask your instructor about it. If you are
having problems with a specific term, graphic aid, etc., then ask about it.
Reflect--Say to yourself as you finish a section “What are the main points?, “How does this
relate to things I learned before?”, “What are some everyday examples of what I just read about?”,
“Does this make sense?”
Self-Recite. Every few paragraphs look away and “play back” what you have just read by
summarizing it in your own words. Write the summary down and say it aloud to yourself. Many
students remember information better if they write it down and read it aloud to themselves.
If you find that you have difficulty summarizing the material in your own words, you may need to
go back and look at that material again. Look then for any terms whose definitions might be fuzzy
in your mind. If you find some, look those terms up in the glossary. Again, if you still don’t
understand or feel unsure after you have made an honest effort at the material, ask your instructor
for help.
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Test Yourself. Recall the broad organization of the chapter and decide how the sections
you’ve just read fit into the rest of the material. See if you answered your questions and if you have
any new ones. Run through the major points and subpoints of the material. Review this
information again the next day and you’ll remember it better than if you wait until exam time!
I would not advise highlighting or underlining the material as you read it the first time.
Most people underline/highlight too many things when they do this because they haven’t yet sorted
out what is important to remember. After you summarize to yourself, then it may be useful to take
notes or go back and mark the passages that are particularly important. Because many students
remember information better if they write it down, taking notes is often preferable to highlighting.
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Supplement 2: Classifications of Matter--Elements, Compounds and Mixtures
There are several different ways that scientists classify matter. In each case, some set of
properties is used to develop the classification scheme. The scheme discussed below differentiates
pure substances from mixtures, and then distinguishes the two types of pure substances: elements
and compounds.
Macroscopic Differences between Mixtures and Pure Substances.
A mixture is a mix of two or more things, each of which retains many (but not necessarily
all) of its own properties. Perhaps most importantly, each of the components can still chemically
react the same way that it did when it was separate. The components of the mixture can be in any
percent (ratios). There is no chemical reaction when the mixing occurs--no new substance is
produced. As an example, consider sand and sugar put together. The sugar still tastes sweet and the
sand is still gritty, even when mixed together. You can mix 2 liters of sand and 1 liter of sugar or 5
liters of sand and 3 liters of sugar--any amounts of each that you want can be mixed together. There
is no chemical reaction between the sand and sugar. If pure sugar is heated to high enough
temperature, it will burn, and the sugar in a mixture of sand and sugar will also burn.
There are two types of mixtures--heterogeneous and homogeneous. As long as they are well
mixed, they are classified simply by how they look. Homogeneous mixtures look uniform to the
naked eye while heterogeneous ones do not look uniform. A water solution of sugar is a mixture of
water and sugar. It looks uniform, so we would classify it as homogeneous. A mixture of oil and
water will not look uniform and we classify it as heterogeneous. The principle importance of the
distinction between homogeneous and heterogeneous is that heterogeneous mixtures are usually
much easier to separate into components than are homogeneous ones. It should be noted that the
term homogeneous can also be applied to a pure substance—a sample of water is homogeneous, for
example.
In order to show that a sample is a mixture, you must show that there is more than one kind
of “stuff” there--that some of the sample has one property and some of it another property, all under
the same temperature and pressure conditions. This is generally done by separating the mixture into
parts that have different properties. To do this, you must find a property which is different for the
components of that particular sample. For example, if your sample happened to be one of sugar and
sand and you put water in it, the sugar would dissolve and the sand would not. Some of the sample
would dissolve, some would not, so that would show that the sample was a mixture.
A pure substance has only one kind of chemical in it. In chemistry the term “substance” is
used to mean what we would call “pure substance” in everyday life. Oddly enough, if you don’t
have sophisticated instruments, then to show that something is a substance, you carry out tests to
see if it behaves like a mixture! If it behaves like a mixture, it is not a substance. If you carry out
many, many tests on a sample and it doesn’t behave like a mixture in any of them, it is most likely a
substance. However, this really doesn’t prove that the sample is a pure substance. What if you
just haven’t found the correct property to check yet? You can never know without the necessary
instruments. Lacking definitive proof, you would have to cite what you have done and say that all
results support the sample's being a pure substance.
The term “pure” in everyday life is different than “pure” in science. In everyday life, we
sometimes say that a sample is pure if there is nothing unexpected or unwanted in it. Thus, in
everyday life we might refer to “pure bleach”, which would contain water as well as the bleaching
chemical. Using the scientific meaning of the term pure, the bleach would be a mixture; in science
a sample is called pure when only if there is only one substance present in the sample.
Macroscopic differences between Elements and Compounds
There are two types of substances: elements and compounds. A compound can be made to
react to produce two or more different substances that were not originally there. Remember, all of
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the sample is the same substance in the beginning, then it chemically reacts to form two or more
new substances. (We say it has been decomposed). Elements cannot be made to react this way.
If you want to show that a substance is a compound instead of an element, you must make it
chemically react (without other chemicals present) to produce some other substances. If you pass
an electrical current through a pure water sample, two gases are formed that were not there before
the experiment. This shows that the original sample was a compound.
In the absence of sophisticated instruments, if you want to show that a substance is an
element, you must test it to see if it behaves like a compound! If you have run many tests and
cannot get the substance to decompose, it is most likely an element. However, you wouldn’t have
proven that the sample is an element; suppose you just didn’t find the correct test? In this case, you
would report all the tests and their results and say evidence supports the hypothesis that the
substance is an element. One other thing you could do is to check if the sample has the same
properties as a particular element. The more properties that are the same, the more likely the two
are the same substance. For example, if a sample has the same appearance, density, melting point
and electrical resistance as copper, it seems likely that it ‘s copper.
To summarize the macroscopic characteristics of elements, compounds and mixtures:
Matter
Does it have components with
different properties under
same conditions?
Yes
No
Mixture
Substance
Can it be decomposed?
Does it appear uniform?
Yes
No
Yes
Homogeneous
mixture
Compound
No
Element
Heterogeneous
mixture
Question: Which of the following would be pure substances? mixtures? elements? compounds?
heterogeneous? Explain how you arrived at your answer.
(a) a piece of wood
(b) a shoe (c) coffee beans (d) a piece of aluminum wire
(e) water
.
Further Classifications of Mixtures
Earlier, we discussed how to classify mixtures as homogeneous and heterogeneous. I would
like to look briefly at the two other categories of mixtures that you will sometimes encounter in
everyday life and in lab: solutions and colloids.
A solution is just a homogeneous mixture. Whichever solution component is present in the
largest volume is normally considered the solvent. The other components are called solutes.
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There can be solutions of solids dissolved in liquids, liquids dissolved in liquids, gases
dissolved in liquids and gases dissolved in gases. Strictly speaking, we can have a solid in a solid
and a gas in a solid, but we won’t frequently encounter these in CHEM 101 as we will concentrate
on solutions that have a liquid solvent.
Examples of solutions are sugar in water, alcohol in water, oxygen in water and oxygen in
nitrogen. There may certainly be several solutes involved. For example, if you put sugar and Koolaid in water, the sugar and the Kool-aid are both solutes.
How would you know if a liquid can be classified as a solution? First, it must contain more
than one component since a solution is a mixture. Second, it must be homogeneous. How can you
tell if it’s homogeneous, though? In lab we will frequently “determine” this by watching for certain
characteristics and using a little common sense. If a mixture with a liquid solvent is to be a solution,
it must appear clear--you must be able to see through it. Note that this does not mean the same
thing as colorless. A colored solution can be clear. For example, if you put blue food coloring into
water and stir, the resulting solution is blue, but it is still clear.
Frequently if you pour one colorless liquid into another one, the two don’t mix very well by
themselves. This is sometimes evident by “wavy” lines that appear in the resultant liquid. These
“lines” are not the result of a reaction, but just the fact that the two components transmit light
somewhat differently. Even if the lines don’t appear, you shouldn’t assume that your mixture is
homogeneous until you have stirred or otherwise carefully agitated the mixture.
To summarize, in lab you may assume that you have one solution in a container if the mixture is
clear and you have made sure the components are well mixed.
How are liquid samples that are not solutions classified? If they are not mixtures, they are
just pure liquids. If they are mixtures and the components will separate by themselves, then they
are heterogeneous mixtures. The other classification is colloid. In a colloid the particles are too
small to settle to the bottom of the container, but they are large enough to scatter the light that
passes through them. Thus they appear cloudy, yet they do not separate by themselves even after a
long period of time.
Solution Concentrations.
Sometimes we need to express how much solute is present in a particular amount of
solution. We may do this in a qualitative or quantitative way.
In a qualitative manner, we say that a solution is more concentrated if a particular volume
of solution contains more solute. For example, wine and beer are both solutions which contain
alcohol. One cup of wine contains more alcohol than one cup of beer, so we would say that the
wine is more concentrated in alcohol than the beer. If we want to qualitatively state that a solution
contains very little solute in a certain volume of solution, we say the solution is dilute. The beer is
more dilute in alcohol that the wine. Note that the terms concentrated and dilute imply comparisons
and communication is generally more clear if we say specifically what solutions we are comparing.
There are many ways to quantitatively state the amount of solute in a solution. For now, we
will consider only units that involve mass of solute in a volume of solution.
One of the most common ways to express the concentration of a solid in a liquid solution is
as a percent. This generally means the number of grams of solid that are dissolved in 100 mL of
solution. For example, if I say a solution is 5% salt, I mean that each 100 mL of the solution
contains 5 g of salt.
Questions:
(1) If a solution is 7 % in salt, how many grams of salt are in one liter of the solution? in 5 liters of
the solution? in 300 mL of the solution?
(2) If a solution is made by mixing 9 grams of salt with enough water to make 20 liters of solution,
what is the concentration of the solution in percent?
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Supplement 3: Solids, Liquids and Gases--Macroscopic Properties
Properties of Solids, Liquids and Gases.
You can distinguish a solid from a liquid from a gas, but you may not have explicitly thought
about which properties allow you to do so. The most obvious of these is shape. Solids have their own
shape. Also, if I pick a certain volume of a solid, its mass will be much higher than that of the same
volume of a gas; this is another way to say solids have higher densities than gases. It is very difficult to
compress most solids. If you push down on most solids, their volumes do not decrease. Instead, the
pressure is transmitted through the solids to items on the other side of the solids. An example of a solid
with these properties is a piece of copper. Each piece has its own shape, which does not change if it is
put into a different container. The density of a piece of copper is fairly high compared to that of a gas,
such as air, and pushing down on the piece of copper will not decrease its volume (though it might
change its shape).
Liquids take on the shape of their containers on all sides except the top surface. Except for the
flat top surface, the boundaries of water in a glass are determined by the glass. The density of liquids is
generally high compared to gases and they are very difficult to compress. As an example, liquid water
takes on the shape of its container, a chosen volume of water has much more mass than that volume of
air, and pushing on water does not change its volume.
Gases take on the shape of their containers on all sides. You may not have thought of gas samples
as having a shape, but consider what happens if you open a container of a smelly gas in one corner of a
room. Soon you would be able to smell it in all corners of the room, evidence that the gas has moved into
all of the room. The boundaries of the gas are then the walls of the room, so the gas sample has the shape
of the room, its container. Gases are compressible--if you apply pressure to them, their volume gets
smaller. Gases are also much less dense than solids and liquids. For example, when you squeeze on a
balloon on all sides the volume of the air in the balloon gets smaller, and the mass of the air in a balloon
is much less than that of the water in a water balloon!
To summarize:
solids
liquids
gases
own shape, all sides
shape of container, except top
shape of container
higher densities
higher densities
lower densities
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hard to compress
hard to compress
easy to compress
Supplement 4: A Submicroscopic View of Matter
Scientists believe that all matter is composed of extremely tiny particles called atoms. The
properties of a sample are determined by what types of atoms are present as well as whether and
how the atoms are bonded together. If the way the atoms are bonded together changes, the
substance present is different. We’ll now look at how scientists think some of those submicroscopic
(atomic level) properties determine macroscopic properties of mixtures, elements and compounds.
Submicroscopic Differences Between Elements and Compounds--a Theory.
Elements.
Scientists believe that all of the atoms in any one element are the same in essentially all
ways (more on this later). During chemical reactions, the atoms themselves are not changed. Each
element has a different kind of atom than the other elements. Thus a theoretical definition of an
element is a substance with atoms that are all the of the same type. We will put some further
stipulations in this definition later in the term; for now just imagine that if you could see the atoms
in any one element, they would all be alike. The atoms in anything are extremely tiny—it would
take a billion billion of them lined up to go across the period at the end of this sentence.
For now, you may imagine them looking much like extremely tiny balls or bb’s.
There are often attractions between atoms that are close to each other, so in most elements
any atom is directly bonded to other atoms. When we say that two atoms are bonded together, we
mean that they are somehow “stuck” to each other and travel together. In a few elements the atoms
are not bonded to other atoms. In a number of other elements, all of the atoms are bonded in pairs
and in most elements the atoms are bonded in larger groups. Whenever two or more atoms are
bonded together in small groups, the bonded group is called a molecule.
If we use a circle to represent an atom, then samples of elements might look like the
following. Note that we represent all the atoms in any one element as looking the same. This is
what shows that the sample which we are representing is an element.
monoatomic
element
diatomic
element
There is only one type of atom in an element sample and the type of atoms present
determines what element is there. A chemical reaction cannot change one type of atom into another
type. An element sample has only one type of atom and a chemical reaction cannot make it contain
another type of atom. Thus, if you begin with one element, you cannot make a different element or
a compound . Also, the element cannot be decomposed. (By the way, if the paired atoms of an
element are separated, they immediately pair up again, so that doesn’t count as a decomposition.)
Compounds.
Scientists believe that one type of compound is composed of molecules that contain more
than one type of atom. If we again use circles to represent atoms, a sample of a compound might
look like the following. Circles that look different represent atoms of different elements.
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an example of a compound
an example of another compound
The molecules of a compound have different properties than the separate atoms or
molecules in an element, so the macroscopic properties of the compound are different than those of
the elements used to make the compound. For example, molecules of water have different
properties than the elements oxygen or hydrogen, even though water molecules contain oxygen and
hydrogen atoms. Because the molecules of water are different than those of the elements oxygen
and hydrogen, at least some of the macroscopic properties of water are different from those of
oxygen and hydrogen. Ammonia is a smelly gas which is used for fertilizer and whose solutions are
used for cleaning. One molecule of ammonia contains three hydrogen atoms and one nitrogen
atom; however, neither nitrogen nor hydrogen smell and neither are particularly good cleaners when
dissolved in water. Plants cannot take up nitrogen as an element, so the element nitrogen, when
found in its pure state, is not a particularly good fertilizer. You can see that the ammonia has very
different properties than the hydrogen and nitrogen that were used to make it.
A note on terminology is in order here. An atom of element “X” may be called “an atom of
X” whether it is in a molecule of a compound, a molecule of an element or a separate atom. To
communicate clearly, we must distinguish where the atom is located. For example, oxygen atoms
exist in both water molecules and in oxygen molecules in elemental oxygen. We refer to the
oxygen atoms in the water molecules or oxygen atoms in oxygen molecules, not just oxygen atoms.
Similarly, we must not refer to “the oxygen in water” as this could mean either the oxygen atoms in
the water molecules or the oxygen dissolved in a water-oxygen mixture. We must be specific about
which we mean or people will misunderstand. This is important because the oxygen in the water no
longer has the properties of the pure oxygen.
To avoid having to say “molecule or separate atom”, I will introduce the term
submicroscopic particle. When we are describing something that can be true of either molecules or
separate atoms or both, we will use the term submicroscopic particle. This term does not mean
atoms within molecules.
Decomposing a compound.
When a compound is decomposed, it can form two (or more) substances because the
molecules can be pulled apart and combined in more than one way during the reaction. If the
compound represented above was decomposed, it might form substances whose submicroscopic
particles look like the following:
can decompose to form
a compound
a mixture of the same
atoms bonded differently
Notice that the total number of atoms of each type is the same as in the original compound-they are just connected in different ways. This difference in how they are bonded together makes
the sample have different properties than it did at the start. Also, there are two types of
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submicroscopic particles present now, so the sample represented immediately above is no longer
just a compound. It is, in fact, a mixture.
Mixtures
Scientists believe that mixtures contain a mix of separated atoms, a mix of molecules or a
mix of both. An example of a mixture would be:
This is an example of one element and one compound. We do NOT say that it is three elements.
When we classify a sample as an element, it means there is only one type of atom present. If we say
a mixture contains an element, it means there are either single atoms or molecules with all the same
type of atom in them within the sample. If we mean the element within a compound, we must say
the “within” part—for example, if we mean the oxygen within water, we must say the “within
water” part or is sounds like there is pure oxygen there.
There are an infinite number of types of mixtures. We could have mixtures of three
elements, mixtures of three compounds, mixtures of one element and two compounds, etc.
A mixture of 3 elements.
A mixture of 3 compounds
A mixture of 1 element and
2 compounds
In a mixture, each type of submicroscopic particle retains its own properties, so all parts of a
mixture do not have the same properties. When the different types of submicroscopic particles can
be separated from each other, then the mixture has been separated into its components. If we were
to separate the first mixture represented above, the before separation and after separation
representations would look like the following.
mixture, before
separation
mixture components separated
from each other
Because the bonding has not changed, there was no chemical reaction.
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To summarize the submicroscopic properties of elements, compounds and mixtures:
Matter
Is there more than one type
of submicroscopic particle?
Yes
No
Mixture
Substance
Does it have only
one type of atom?
Yes
No
Element
Compound
Question: Draw a picture which represents the submicroscopic view of (a) a mixture of aluminum
and sulfur.
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Supplement 5: Submicroscopic Particles And Properties Of Solids, Liquids And Gases.
Now we will look at how scientists believe the behavior of submicroscopic particles leads to
the shape, density and compressibility properties of solids, liquids and gases. The term
“submicroscopic particles” means separated atoms and/or molecules, but not atoms within
molecules.
Gases. Scientists believe that the submicroscopic particles of gases are fairly far apart and have
nothing between them. They believe these submicroscopic particles move constantly and that those
in any given sample of gas will move faster if its temperature is increased.
a gaseous element
a gaseous compound
a gaseous mixture
Scientists believe that a submicroscopic particle in a gas moves in one direction until it
bounces off another particle or the container wall. It then keeps moving in its new direction until it
bounces off something else. In this way, the gas’s submicroscopic particles move about until they
reach all portions of their container. Because the gas’s submicroscopic particles are far apart and
have nothing between them, when a force is applied they can be pushed into a smaller volume
without applying much force. This is why it is easy to compress a gas sample. Also, because the
gas’s submicroscopic particles are far apart and have nothing between them, there is little mass in a
volume of the sample, so the gas sample has a low density.
a sample of a gas
same sample, compressed
Liquids. Scientists believe that submicroscopic particles of liquids are much closer together
and not as free to move compared to those in gases. They are not as free to move because the
submicroscopic particles in a liquid are attracted to each other in some way. They can move some,
though, and so a liquid doesn't have its own shape. Because the submicroscopic particles of liquids
are fairly close together, it is very difficult to push them closer together--so liquids are hard to
compress. Also, because the submicroscopic particles are close together, there is a large amount of
mass in a volume and the liquid is dense.
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a liquid element,
maybe mercury
a liquid compound
maybe water
a liquid mixture
Solids. Scientists believe that the attractions between the submicroscopic particles of a solid
are stronger than those between submicroscopic particles of a liquid. This causes the
submicroscopic particles to remain in essentially the same spot except for some "quivering" around
that spot. The attractions between the submicroscopic particles of a solid hold the submicroscopic
particles in place. Gravity can’t pull the particles away from the other particles close to them.
Because of this, the solid has its own shape. Like those in a liquid, the submicroscopic particles of
a solid are close together. This makes it hard to push them closer and makes the solid very hard to
compress. Also, because the submicroscopic particles are close together, there is a large amount of
mass in a volume and so the solid is dense.
A solid compound
a solid element
perhaps Cu
Question: Draw a representation for a solid mixture of two compounds.
Representing Elements & Compounds with Pictures and Use of Terminology
Visible samples of substances have some properties that single atoms/molecules don’t have.
For instance, water has a boiling point and it can evaporate. Evaporation describes a change in the
relationship between submicroscopic particles; it says that the submicroscopic particles get much
farther apart and move around a lot more. Because evaporation refers to a change in how molecules
relate to each other, it really doesn’t make sense to talk about a sample that has only one molecule
evaporating. Because of this, we need to distinguish between submicroscopic particles and
macroscopic samples. A molecule and a compound are not the same thing. They are related, but
the terms do not mean exactly the same thing. When we refer to the submicroscopic particle, we
should say the molecule of the compound, not the compound. That is,
is not a compound.
Thus if you are asked to represent a compound or an element, you should draw several of the
submicroscopic particles that define the element/compound in order to show that you know the
difference between the macroscopic and submicroscopic particles.
Similarly, an atom is not the same as an element. Draw more than one atom to show a
monoatomic element.
Remember the following, and you'll be OK: the terms atom and molecule refer to
submicroscopic particles while the terms element and compound refer to macroscopic samples,
generally large enough to see.
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Supplement 6: Chemical Formulas
Let’s consider ways we can represent molecules. We have used drawing pictures with
circles to represent atoms and molecules. For example, an ammonia molecule has 3 hydrogen
atoms bonded to a nitrogen atom, so we could represent the ammonia molecule as
. While there
is something conceptually satisfying about representing molecules this way, it would be hard to
think of enough easy-to-draw geometric shapes to represent all 109 elements. So, scientists instead
use letters to symbolize each element.
Each element’s symbol is unique and may be used to represent either a sample of that
element or one atom of that element. The first letter of the symbol is always capitalized and if there
are any letters after the first, they are always lowercase. Most of the symbols we use come from the
letters in the English version of the element’s name. Carbon, for instance, has the symbol C, and
nitrogen has the symbol N. Some element names are two letters long, for example magnesium’s
symbol is Mg and lithium’s is Li. A few of the elements have symbols derived from their Latin
names. For example, Na stands for sodium, whose Latin name is natrium. So far we have only
considered how symbols represent atoms, but we should also consider how they are used to write
formulas to represent molecules.
If C represents a carbon atom and O represents an oxygen atom, then a molecule composed
of one carbon atom and one oxygen atom can be represented as CO. Similarly, a molecule made up
of one N atom bound to one O atom has the formula NO.
The octane molecule has eight carbon atoms and eighteen hydrogen atoms. The easiest way
to do represent this is to write the symbol for each atom in the molecule and to write, as a right
subscript to that symbol, the number of atoms of that element in one molecule. For example,
octane’s formula is C8H18. Another example would be sucrose, which has 12 carbon atoms, 22
hydrogen atoms and 11 oxygen atoms. Its formula is C12H22O11.
Sometimes we want to represent several molecules of the same type. To repesent more than
one molecule, put a number in front of the formula instead of changing the subscripts. For example,
suppose that we want to represent three molecules, each composed of one N atom and one O atom.
We cannot write N3O3 for this formula since N3O3 means one molecule composed of three
nitrogens and three oxygens (maybe
). Instead, we write 3 NO. By writing the three in
front of the entire formula, we show there are three NO molecules, each separate from the other.
With circles as atoms, we could draw 3 NO as
.
Sometimes you will also see parentheses in a formula. Generally, the right parenthesis is
followed by a numerical subscript. The subscript refers to the number of whatever is symbolized
inside the parenthesis. As an example, consider the formula CH3(CH2)22CH3. Since the 22 means
22 of everything in parenthesis, the number of carbon atoms in the molecule represented by the
above formula is 1+22+1 = 24 and the number of hydrogen atoms is 3+(2)(22)+3 = 50.
Notice the difference between the number in front of a formula and the numbers written as
subscripts. The numbers in front of the formulas apply to everything in the formula. Subscripts
apply only to whatever comes immediately before the subscript or to what is in parentheses if they
are right before the subscript.
Each formula may be used to represent either a sample of that element or one atom of that
element. We must either tell from the context or else actually say whether we mean a sample or a
molecule or an atom. If we draw
, it means a single submicroscopic particle, but if we write
NH3, it may mean either a molecule of the substance or a sample of it--we must somehow say
which we mean.
133
Supplement 7: Metric System and Part 1 of Factor Label Problem Solving
The Metric System and Metric-Metric Conversions Using Factor-Label
We shall concentrate on metric units of length, volume and mass and temperature. These
are the units most commonly used in freshman chemistry labs. The metric system will make more
sense to you if you pick some everyday objects and learn their approximate metric size . These
items then serve as your reference when you are trying to visualize the size of something. For
instance, a nickel has a mass of about 5 g, a sugar cube has a volume of about 1 mL and most
people have an index fingernail close to 1 cm across (check yours!). Also, for those of you familiar
with cubic centimeters, 1 cubic centimeter (1 cc) is exactly the same as 1 milliliter.
One other metric unit you should consider is that of temperature. The most common temperature
o
unit used in the chemistry lab is the degree Celsius, abbreviated as C (which is read “degrees C”).
o
Approximate temperatures for things you commonly encounter are: water freezes at 0 C and boils
o
o
at 100C. Room temperature is about 20-22 C and human body temperature is about 37 C. A
degree Celsius is 1.8 (roughly 2) times bigger than a degree Fahrenheit.
Prefixes. The metric system uses prefixes which say to multiply the root unit by some
factor of ten. Each prefix has the same meaning for all of the root units. For example, the length
of one centimeter is 1/100 of a meter, one centigram is 1/100 of a gram, and a centiliter is 1/100 of a
liter. Prefixes, their meanings and their abbreviations are listed in your text. You should learn the
ones your instructor requires.
Metric-metric conversions. We can use the relationships indicated by the prefixes to
change between metric units properly. For example, the number of centimeters in 34 meters can be
obtained by using the factor 100, but should I multiply by 100 or divide by 100? This problem can
be overcome with common units by considering the size of those units. A meter is roughly the
distance from my nose to the tip of my outstretched arm and a centimeter is roughly the distance
across my fingernail. Because a centimeter is small compared to a meter, I would expect a distance
expressed in centimeter units to have a larger number than the same distance expressed in meter
units, so I would multiply by 100 to convert meters to centimeters: 34 m = 3400 cm.
This method of converting works well for units that you are familiar with, but there is a more
general method which will help in more cases, including many types of problems that are not just
conversions of one unit to another. The method goes by the names factor-label or dimensional
analysis.
Single Step Problems
Factor Label problem solving is used to convert a unit or set of units into others that are related to
the initial ones by ratios.
The method involves: writing down:
what we know (with units) and
the units that we want to end up with
Then finding a ratio or ratios involving those units and
then using the ratios and unit cancellation to get the quantity that we want.
Example 1: Suppose we want to know how many meters are in 6.7 kilometers. We have 6.7 km
and want to know our answer in m, so we write what we know and the units that we want to end up
with
6.7 km
= ?m
1000 m
Now, to work the problem, we need a ratio that relates km and m. A ratio that does this is
1 km
If we multiply the 6.7 km by this ratio, then the km units will divide out
134

because they are on opposite sides of the fraction bar. This will leave the units m, which is what we
wanted. Then we just do the arithmetic with the numbers, using our calculator if necessary, and
we’re done.
6.7 km
1000 m = 6700 m
1 km
(Note: if a number is written by itself instead of in a ratio, you may regard this number as over 1 if
6.7 km
you wish. That is, 6.7 km could also be written as
).
1
Example 2: Suppose 6 mL of oil have a mass of 5.4 g and we want to know how much volume is
occupied by 1293 g. The numbers that come from the same sample are the 6 mL and the 5.4 g. We
know the other sample has a mass of 1293 g and want to know its mL. We set up the number we

know and the units we want to end up with:
1293 g
= ? mL.
We can see that we would like to have a ratio that involves g and mL so the units cancel:
1293 g
mL = ? mL . Now we need to find the numbers that go into the ratio. Since
g
there are 6 mL for each 5.4 g, we write
1293 g 6 mL = 1437 mL.
5.4 g
Writing Ratios
Having seen the problems above, we might ask “How do we know what ratio to use?” Besides
involving units that we want to get rid of and introduce, in order to write two numbers as a ratio,
they must pass the following tests: they always occur together in the situation being examined and
when one is multiplied by a number the other must also be multiplied by the same number. If
both of these things are not true of two numbers, then they cannot be written as a ratio.
Let’s go back to the examples above and check if we could correctly write the numbers as ratios.
How do we know that 1 km and 1000 m can be written as a ratio? Well, if an object is 1000 m
long, it is also 1 km long, so when one quantity occurs so does the other (test 1). If some object is
5000 m long, then it must also be 5 km long, so when the 1000 m was multiplied by 5, the 1 km was
also multiplied by 5, so the second test is passed. Because both of our tests were passed, we can
write ratios of 1000 m and 1 km:
1000 m
1 km
and
1 km
1000 m
We choose the ratio to use by checking which gives the unit cancellation that we want.
 Step Problems. Sometimes we don’t know a ratio with the two units
Factor Label and Multiple
we want to change. We then need to go through an intermediate step or steps--we convert what we
know into “something” and then convert the “something” into what we want.
Suppose, for example, we need to find how many kiloliters are in 12.5 milliliters. First we write
down what we know, and the units we want:
12.5 mL
= ? kL
We don’t know a direct relationship between mL and kL, so we just look for things we do know
about the two units. From the prefixes, we know 1000 L = 1 kL and 1000 mL = 1 L.
135
We could use the relationships to write ratios if they meet the criteria for ratio writing. If we have
1000 mL then we have 1 L and if we have 2000 mL we have 2 L, so both ratio tests are passed and
1L
we can write 1000 mL
The two numbers 1000 L and 1kL also pass the tests, so we can write the ratio
1 kL
1000 L
Having written the ratios, we use them to go from mL to L and then to kL. We could diagram our
plan as mL  L  kL.
The first step is kL to L. If we multiply the number of mL by the firstratio above, mL will then
divide out, leaving L.
1L
12.5 mL
= 0.0125 L
1000 mL
Now we must convert the 0.0125 L to kL. If we multiply by the second ratio, L divide out and kL
are left.
Problems that are not Metric-Metric Conversions . You can use the factor-label method to
interconvert any units of mass, or any units of length or any units of volume. This is true even if the
unit is “mixed in” with others. Just do the conversion one unit at time, carrying along the other
units until another step.
For example, if a car travels at a speed of 90 km per hour, how many m per minute does it
90km
travel? Things that we know are:
is a valid ratio for this problem since it was given in the
1hr
problem statement. (whenever a problem says “per”, we can write a ratio). We also know 1000 m =
1 km. Checking the ratio criteria, if we have 1000 m we have 1 km (having one means we have the
other) and if something is 3000 m it is 3 km, (multiplying one meant the other must be multiplied
 1000m
the same), so we can write
. Using the same type of checks for 1 hr. and 60 min., we see we
1km
can write
1 hr
. Thus our “known” information includes
60 min



90km 1000m 1 hr
?m
and we want
1hr
1km 60 min
min
Looking at the units in the ratios that we have written, we see that we can convert km to m and we
can convert hr. to min. Our plan is then to convert the units one at a time, carrying along the other
 unit present

 km/hr  m/hr.  m/min.
until the next step:
then
136
When we become comfortable with using factor-label, we will find that we can write the steps
together on one line:
Another example:
If 3 mL of oil weigh 2.7 g, what is the mass of 1100 L of oil? From the problem we know
have: 1100 L oil
want:
? g oil
3 mL oil is 2.7 g oil
We also know 1000 mL = 1 L.
We now need some ratios. If we have 3 mL oil we have 2.7 g oil. If we took a sample with 7 times
the volume, we would expect 7 times the mass, 10 times more volume would be 10 times more
mass and so forth. Because the ratio-writing criteria are true, we can write
3 mL oil
2.7 g oil
and
.
2.7 g oil
3 mL oil

The units of these ratios show us we can take g oil and get mL oil or else take ml oil and get g oil.
Unfortunately, we have neither, so we must ask whether we can get either. Looking back at the

information
we know, we have 1100 L of the oil, and we could use that to get mL of oil. So, the
plan is 1100 L oil  ? mL oil  ? g oil
Now that we have a plan drawn out, we just need to follow it.
I wrote the calculation all in one line this time, but you could split it into steps if you wish.
137
Supplement 8: Accuracy and Precision
A measurement is accurate if it is close to the “correct” value. We often do not know the
correct value and cannot always establish the accuracy of a measurement. If we do know the
correct value then the accuracy can be expressed as the percent error:
% error 
measured value  correct value
(100)
correct value
A measurement is precise if it is reproducible--if the measurement is done over and over,
then a number close to the same value is obtained each time. Precision is used to compare
measurements.If I measure the length of a room with a meter stick several times and get values of
7.4 m, 7.3 m and 7.5 m, then that measurement is more precise than if I got values of 7.0 m, 8.0 m
and 7.5 m. This is true since in the first case the digit in the tenths place varies while in the latter
case the digit in the ones place varies.
If the hitting the bullseye is accurate, the pictures below illustrate samples that are accurate
and precise, inaccurate and precise, imprecise and inaccurate and imprecise and “accurate” (where
the average position is close to the bullseye).
138
Supplement 9: Factor Label Problem Solving, Area and Volume
Area and Volume Units that are (Length)2 and (Length)3
What do unit such as 1 m2 and 1 m3 mean? A m2 means the same area as that of a square
that is one meter on each side. A 1 m3 means the same volume as that of a cube that is one meter
on each side.
We can use the relationships that we know about between length units to develop
relationships between volume units. Let’s use the relationship between cm and m to find the
1m
relationship between cm2 and m2. Since there are 100 cm in 1m, we can write
. If we
100 cm
square the whole ratio, we have the relationship between cm2 and m2:
 1 m 2
12 m2
1 m2


 
100 cm  1002 cm2 10,000 cm2 
Note that we had to square the numbers as well as the units!
1m
, if we
100 cm
 1 m 3
13 m3
1 m3

cube the whole ratio, we’ll have what we want. 
.
 
100 cm  1003 cm3 1,000,000 cm3
3
3
In a similar way,
 we can find the relationship between cm and m . Since
We can summarize more generically by saying (length ratio)2 = area
 ratio and
(length ratio)3 = volume ratio.
 is sometimes useful to know is that 1 cm3 is exactly 1 mL.
Another volume relationship that
Example Problems with Squared, Cubed Units
1. Suppose the carpet A costs $3.50 per ft2 and carpet B costs $29 per yd2. Which is more
expensive? To compare, we need to change one of the prices to the same units as the other one.
$29
$
Suppose we choose to change the
to
. We need a ratio that relates ft2 and yd2. We know
2
yd
ft 2
1yd 2 1yd 2
1yd
there are
, so if we square the whole ratio, we’ll have   
. Now, arranging the
 3ft  9 ft 2
3ft
$29 1yd 2 $3.22



ratios so the units cancel gives us our answer: 2
. So, carpet B is cheaper.
yd 9 ft 2
ft 2
2. How many liters of water would be needed
to cover a 2 m x 100 m raised flower bed to a


depth of 1 cm?
To work this problem, we need to visualize the water as a rectangular-based solid resting on
 of such a solid is width x length x height. For such a
top of the soil in the bed. The volume
calculation to make sense, the length, width and height must all be in the same unit, so let’s change
100cm
100cm
the bed dimensions to cm. 2m
 200cm and 100m
10,000cm . Now, since the water
m
m
height is to be 1 cm, the volume is 200cm x 10,000cm x 1 cm= 2,000,000 cm3. To get to liters, it
may be easier to first remember that 1 mL is exactly 1 cm3, so 2,000,000 cm3 = 2,000,000 mL.
1L
Now, all we have
to do is change mL to L: 2,000,000mL
 2,000L
1000mL
139

Supplement 10 Elementary Atomic Structure and the Periodic Table
Various interactions of matter and energy have led scientists to propose several models for
the structure of the atom. Each of these models is useful in different situations. The general rule is
to use the simplest model that will do the job at hand. Here we will only summarize two of the
models.
One of the models stipulates the identity, general location and some properties of the basic
particles in an atom. This model proposes the following:
Atoms are composed of particles called protons, neutrons and electrons. Each atom of a
specific element has a certain number of protons in its nucleus. A proton or a neutron weighs about
2000 times as much as an electron. Thus, almost all the mass of an atom is due to the protons and
neutrons it contains. However, the protons and neutrons are packed together very tightly into a very
small volume at the center of the atom. The protons and neutrons together compose what we call an
atomic nucleus. The rest of the volume of an atom is empty space with its electrons scattered about.
If the atom was the size of a baseball field, the nucleus would be about the size of an fly. Since
most of the mass of the atom is in the nucleus, yet the nucleus occupies very little space in the atom,
the atom is mostly empty space! This may seem nonsensical since we have said that everything is
made up of atoms--how do things feel solid, for instance? Something must keep the most massive
part of the atoms (the nucleus) from coming so close together.
The protons and the electrons have electrical charges. The proton has a positive charge and
the electron has a negative charge whose size is the same as that of the proton. The neutron has zero
electrical charge. An atom has an overall zero electrical charge because it contains equal numbers
of protons and electrons and their positive and negative charges balance each other. The interaction
of the charges in one atom with those in another atom are complicated. We will discuss them when
we talk about bonding of atoms.
This model of atomic structure was extended by Bohr and various other scientists, as we
shall discuss later.
We have seen that one model of the atom proposes that each atom of a specific element has
a certain number of protons in its nucleus. How can you find out how many protons are in an atom
of a specific element? That information is summarized on the periodic table. The number of
protons in an atom of an element is given above the element's symbol on the periodic table. For
example, each and every atom of carbon has six protons in its nucleus, so a 6 is written above the C
symbol on the table. The number of protons in an atom of an element is called the atomic number
of the element. Since the number of electrons in an atom equals the number of protons, we can also
tell how many electrons are in an atom of an element by looking at the periodic table. The 6 above
the C on the periodic table indicates that scientists believe that each carbon atom contains 6 protons
and 6 electrons.
6
C
12.01
Figuring out how many neutrons are in an atom is a bit harder because each atom of an
element doesn't contain the same number of neutrons. Atoms with the same number of protons but
different numbers of neutrons are called isotopes. To find out how many neutrons an atom
contains, you must start with the atom's mass number, which is the sum of its numbers of protons
and neutrons. Using the mass number and the atomic number, you can calculate the number of
neutrons in the atom as follows:
mass number – atomic number = number of neutrons
You can tell what the mass number of an isotope is by the way the isotope is named or
symbolized. Sometimes the mass number is written following the name of the element. For
140
example, carbon-12 means carbon with a mass number of 12. The mass number may also be
written as a left superscript to the element's symbol: 12C also means carbon with a mass number of
12. Sometimes the atomic number is also written as a left subscript of the element's symbol. 126 C
also means carbon with a mass number of 12. Writing the atomic number with the symbol isn't
necessary since the atomic number is implied by the elemental symbol.
You may wonder why mass numbers are not on the periodic table. The numbers on the
periodic table apply to elements as they are found in nature, where most elements are mixtures of
isotopes. For example, chlorine samples generally are a mixture of chlorine-35 and chlorine-37,
and carbon is a mixture of carbon-12, carbon-13 and carbon-14.
The numbers below the elemental symbols on the periodic table are atomic masses, which
are average mass numbers for the elements as they are found in nature. The average mass number
of chlorine is 35.45 and the average mass number of carbon is 12.01. The atomic mass of chlorine
is not 36 because there is more chlorine-35 than chlorine-37. As an analogy, think about a group of
people. Suppose their weights are 150 lb., 150 lb., 150 lb. and 190 lb. The average mass of the
individuals in the group is 160 lb. It is not equal to the mass of any one person, nor is it halfway
between 150 and 200 lb.
141
Supplement 11: Light and the Bohr Model of the Atom
Where Are The Electrons?
The model of the atom that was discussed in a previous supplement is fairly specific about
the location of the nucleus, but you may've noticed that it doesn't say much about where the
electrons are located. The model discussed below is intended to show this.
Light
Though sunlight appears colorless, it is composed of many different colors of light, all
blended together. When sunlight is passed through a prism, the different colors of light separate
and the result is a spectrum, rather like a rainbow. This is called a continuous spectrum because
there are no gaps - the colors are side-by-side. If a color was missing for some reason, there would
be a black region in the spectrum where that color normally occurred since black is the result of no
light (like a dark room at night). The other thing that we need to know about light is that each color
corresponds to light of a certain energy. Red light has a different energy than blue light, and so
forth.
When a tube is filled with hydrogen gas and an electrical current is passed through the tube,
light is produced. It is not like sunlight, however. If this light is passed through a prism, the
resulting spectrum consists of only a few narrow lines of color, with large regions of black between
the lines. A picture of such a “line spectrum” . This appearance of line spectra rather than a
continuous spectra puzzled scientists. A Danish physicist named Niels Bohr proposed a model of
the atom to explain what was seen.
Bohr Model of the Atom
Bohr suggested that the electrons in an atom revolve around the nucleus much as the planets
orbit the sun. The orbits could have only certain radii.
Ground State
Excited State
Two dimensional pictures of the Bohr Model of the Hydrogen
atom. Note that for better visibility, the size of the electron
and the nucleus
are both exaggerated.
The orbit of the electron was determined by the energy of the electron. Since electrons
could be in only certain orbits, they could have only certain energies. An electron could move from
one orbit to another by changing its energy: if it absorbed energy it would move away from the
nucleus and if it gave off energy it would move closer to the nucleus.
Bohr said that when the electrical current was passed through the hydrogen gas, the
electrons in the atoms absorbed energy and moved to an orbit farther from the nucleus. The atom
was then in an “excited state”. Soon afterward, the electrons returned to the orbit closer to the
nucleus by releasing some energy as light. Because the orbits correspond to specific energies, the
amount of energy given off by an electron in order to move between specific orbits is always the
same, so for the same transition, the color of the light given off is the same. There are several
different orbits that the electrons can move between, so light of several colors is possible. For a
sample of many atoms, some atoms give off light of one color, some light of another color, but only
a few colors are possible.
142
Bohr went on to develop an algebraic equation which could be used to predict what energies
of light would be given off by electrons in a hydrogen atom. This equation worked very well for
hydrogen, but it couldn’t be used to predict the line spectra of other elements very well. This means
the Bohr model of the atom does not match what atoms are really like. Because of this, other
models of the atom have been developed. However, in spite of the fact that the Bohr model of the
atom is not perfect, it is often useful.
143
Supplement 12: Finding the Number of Electrons in the Outer Energy Level of an Atom
and Lewis Symbols for Atoms
It will frequently be of use to us to know the number of electrons in the outermost energy
level of an atom, which are called valence electrons To find this easily, we can look at the periodic
table. It should be helpful to refer to a periodic table as you read the following.
We will consider only those elements whose symbols are in the first two columns or the last
six columns of the periodic table. Each element with a symbol in the 1st column has 1 electron in
its outermost energy level. Each element with a symbol in the 2nd column has 2 valence electrons.
Each element with a symbol in the same column as B has 3 electrons in its outermost energy level.
Those elements whose symbols are in the column with C have 4 valence electrons, those with
symbols in the N column have 5, those whose symbols below O have 6, those with symbols in the
column with F have 7. The elements whose symbols are in the last column to the right of the table
have 8 electrons in their outermost energy level. I do not advise memorizing based on the number
at the top of the column as this differs in some charts, though the number of electrons does not!
Because elements with the same number of electrons in their outermost energy levels tend to
behave in the same way, elements whose symbols are in the same column on the chart will behave
in similar ways. For example, both sodium and potassium will react with water and the elements
whose symbols are in the last column on the periodic table are hard to get to react at all.
Lewis Symbols of Elements.
The number of electrons in the outermost energy level of an atom of an element is
sometimes symbolized with the Lewis symbol of the element. The is just the element symbol with
the same number of dots around it as the atom has electrons in the outermost energy level.
We can draw the Lewis symbol for an atom as follows. We’ll draw (model) a sodium atom. It is in
group 1 and has 1 valence electron (plus the ten others we don’t show in the picture). The Lewis
symbol for a Na atom would be any one of four as shown:
Na
Na
or
or
Na
Na
or
The Na symbol represents the sodium nucleus plus the 10 inner-shell electrons. The “dot”
represents the lone valence electron. Traditionally, for most Lewis drawings the valence electron(s)
about the atom are often drawn at the corners of an imaginary diamond about the nucleus. We
could represent something like a calcium atom as follows:
Ca
or
Ca
Ca
or
or
Ca
Ca
or
or
Ca
or
Ca
or
Ca
or
Ca
For another example, an atom of oxygen (O – 6 valence electrons) is usually represented as:
O
or
O
or
O
or
144
O
or
O
or
O
Supplement 13: Predicting Ion Charges and Naming Ions
Charges of Ions:
Based on past behavior, scientists can predict the charge on some ions. This is true of ions
formed from elements whose symbols are in the 1st, 2nd and next–to–last columns on the periodic
table as well as a few others. On the periodic table outline below, the symbols have been replaced
with ion charges for those elements.
You can see that elements whose symbols are in column 1 form +1 ions, elements whose
symbols are in column 2 form +2 ions and those elements whose symbols are in the next–to–last
column form –1 ions. Also, O and S form –2 ions while N and P form –3 ions. Al forms a +3 ion
while Zn and Cd form +2 ions. Most of these charges (except perhaps Cd and Zn) can be predicted
by realizing that the atoms gain or lose electrons to match the number of electrons in the nearest
noble gas (last column).
+1
+ 1+ 2
- 3- 2 - 1
+3
+ 1+ 2
- 3- 2- 1
+ 1+ 2
+2
-1
+ 1+ 2
+2
-1
-1
+ 1+ 2
+ 1+ 2
These ions, if they are positive, are named by simply saying the element name followed by
"ion". Thus Al3+ is an aluminum ion. The negative ions are named by changing the ending of the
element name to “ide”; thus, P3– is a phosphide ion and S2- is a sulfide ion.
Practice: use a regular periodic table as needed, and fill in the blank cells in the following. The
answers are at the end of this reading supplement.
Symbol
K+
Name
Symbol
Fe2+
Br-
N3-
Mg2+
O2-
Cu+
Ni2+
Name
Another group of ions that have a predictable charge are the polyatomic ions. These are
formed from several atoms that gain or lose electrons. These are all bound together, sort of like a
molecule except they are charged. For example, SO42– is an ion with one sulfur atom, four oxygen
atoms plus two extra electrons. We might think of it as the following:
2-
There is also one common positive polyatomic ion, ammonium ion. Know the names and
complete formula (including charges) of the polyatomic ions listed below.
145
Name
acetate
carbonate
cyanide
hydroxide
hypochlorite
ammonium
Formula
C2H3O2–1
CO3–2
Name
nitrate
Formula
NO3–1
NO2–1
CN–1
OH–1
phosphate
sulfate
PO4–3
SO4–2
ClO–1
NH4+1
sulfite
SO3–2
nitrite
TO MEMORIZE THE ABOVE, IT HELPS TO REMEMBER THE “ELBOW”
B C N O
Si P S
As Se
Te
F
Cl
Br
I
At
“Elbow” consists of B, C, N, O, F, Cl, Br, I, & At. All other elements are “outside
the elbow.”
“On the Elbow”
Polyatomic ions:
-ate = 3 oxygen
“Outside Elbow”
Polyatomic ions:
-ate = 4 oxygen
Ions ending in –ate always have one more oxygen atom than those ending in –ite.
Acids with oxygen that end in –ic acid have the same number of oxygen atoms as the –ate ion.
Acids with oxygen that end in –ous acid have one less oxygen atom than the –ate ion.
146
Adding a H+ to any of the above results in the same ion name except with hydrogen in front
of it. For example, adding H+ to CO32– gives us HCO3–1, which is hydrogen carbonate and adding
H+ to SO32– gives us HSO3–1, which is hydrogen sulfite. Notice the charge is always one higher
after one H+ is added. Because phosphate has a –3 charge, we can add either one or two H+ ions
to it and still have an ion. Adding one gives us HPO42–, which is called (mono)hydrogen
phosphate; adding two H+ ions gives H2PO4–1, which is called dihydrogen phosphate. You should
also know the names of anions with hydrogen ion added to the ions above that have a –2 or –3
charge.
Metal ions not in groups 1 or 2 or Al or Cd or Zn can form ions of different charges,
depending on the circumstances in which the ion is made. For example, copper can form either a
+1 or +2 ion: Cu+ or Cu2+. To indicate which ion is there, we write the charge in Roman numerals
in parentheses after the element name. Thus Cu+ is copper(I) ion and Cu2+ is copper(II) ion; the
names are read as "copper one" and "copper two", respectively.
Key for Practice: use the periodic table as needed, and fill in the blank cells in the following:
Symbol
K+
Name
potassium ion
Symbol
Fe2+
Name
iron(II) ion
Br-
bromide
N3-
nitride
Mg2+
magnesium ion
O2-
oxide
Cu+
copper(I) ion
Ni2+
nickel(II) ion
147
Supplement 14: Nomenclature of Ionic Compounds
Formulas of ionic compounds from their names:
To write a formula of an ionic compound when you are given its name, simply do the following.
1. Write the complete formulas of the ions present, being sure to include the charge. An
ending of “ide” indicates only one type of atom in the anion unless it is hydroxide and cyanide.
2. Find the number of each type of ion that makes the total positive and negative charge add
to zero. This may be done by algebra, trial and error, or “criss-cross”. This gives the relative
number of each type of ion.
“Criss-cross” is the method of using the sizes of the charges on each ion for the subscript of
the other ion.
3. Write the symbol for the cation first followed by a subscript to indicate how many of the
cations are needed; now write the symbol of the anion followed by a subscript to indicate how many
anions are needed. If there is more than one of a polyatomic ion, its formula is put in parentheses.
4. Make the subscripts of the ions smallest whole numbers while retaining the ratio found in
step 2. Note that step 4 is done only with ionic compounds and not with covalent compounds.
5. Remove any charges indicated on the ions as compound formulas don't show charges.
Example 1: What is the formula for potassium sulfide? Potassium ion is +1 since K is in the first
column on the periodic table. So, the ion in this compound is K+1. The sulfide ion is the ion of S.
The position of S on the periodic table indicates that an atom of it forms a -2 ion, so it is S-2.
1. If I do a “criss-cross” with K+1 S-2, then I will have a subscript of K+1
S-2
2 for K and a subscript of 1 for S
2
1.
3. The subscripts 1 and 2 are in smallest whole numbers already, so the formula is K2S.
Example 2. What is the formula for nickel(IV) oxide?
1. The charge of the nickel ion is +4 (we can tell this from the IV) = Ni+4. The oxide ion is the ion
formed from an oxygen atom and is -2. = O-2
Ni+4
2.
2
O-2
4 . So far this gives Ni O , but we must do step 3.
2 4
3. Both subscripts can be divided by a number larger than 1 and still be whole numbers, so we
should do this. In this case, we divide by 2.
This gives NiO2. Now the subscripts are the smallest whole numbers.
Example 3: What is the formula of calcium acetate?
1. The calcium ion is +2. = Ca+2 . The acetate ion is a polyatomic ion. A good way to recognize
that this is a polyatomic ion is that it does not end in ide. We will have memorized that an acetate
ion is C2H3O2-1.
Ca+2
1
C2 H3 O2 -1
2
.
2. We need two acetate ions, so it must be placed in parentheses. This gives Ca(C2H3O2)2.
148
3. The subscripts 2 and 1 are already in smallest whole numbers. Note that we would never change
the subscript in the polyatomic ion itself. The final formula is Ca(C2H3O2)2.
Practice writing formulas of ionic compounds from their names with the following. Write what you
think is the correct answer, and then go through the rules to check yourself before “grading” your
answer. The numbers refer to the number on the key at the end of the handout.
14.
3.
10.
11.
Name
sodium sulfate
calcium nitrate
iron(III) sulfide
magnesium chloride
Formula
15.
2.
12.
9.
Name
strontium phosphide
aluminum oxide
manganese(II) bromide
iron(II) sulfide
Formula
Ionic Compounds: Names from formulas
On the surface, naming an ionic compound when you have its formula is quite simple. You
simply name the cation (without the word "ion") and then the anion. Thus SrCl2 is strontium
chloride and CaSO4 is calcium sulfate. It gets slightly more complicated if the cation is one whose
charge may vary; remember to name these cations, we must include their charge in Roman
numerals. To figure out a charge on one of these ions, remember that the charge on each metal ion
is the total positive charge divided by the number of metal ions indicated in the formula.
Suppose we need to name FeS. Iron (Fe) can have different charges in different compounds,
so we cannot just say iron sulfide. We must figure out the charge on the iron ion. The total negative
charge is that of one sulfide ion, so it is -2; the total positive charge in a compound must be equal in
magnitude to the total negative charge, so the total positive charge is +2. There is only one iron ion
and +2 charge/one ion = +2 charge on each iron ion. Thus the cation name is iron(II) and FeS is
iron(II) sulfide. Let's do some more.
What is the name of Cr2O3? Cr is not in group 1 or 2 and isn't Al, Zn or Cd, so its ion
charge can vary. Each oxide ion is –2, so there is a total negative charge of –6; this is "balanced" by
the charge of two chromium ions. +6/2 chromium ions = +3 on each chromium ion. Thus Cr2O3 is
chromium(III) oxide.
What is the name of NiSO4? Nickel ions can have various charges, so we must use the
2–
charge of the anions present to determine the charge of the nickel ion there. The anion is SO4 ,
the sulfate ion. Since the charge of one nickel ion balances the -2 of the sulfate, the nickel ion must
be +2, and NiSO4 is nickel(II) sulfate. What is Ni2(SO4)3? There are two nickel ions and three
sulfate ions. Three sulfates have a charge of (3) (-2)=-6, so the total positive charge is +6. Then
+6/2 nickel ions = +3 on each nickel ion. Thus Ni2(SO4)3 is nickel(III) sulfate.
More Practice: Name the compounds whose formulas follow, double-check yourself by going over
the rules above for each case, then check your answer against the “key”.
13.
4.
16.
8.
Formula
MgSO3
CrSO4
ZnO
HgCl2
Name
1.
6.
5.
7.
149
Formula
Al(NO2)3
CuF
Cu(NO3)2
FeO
Name
Key for Ionic Compound Nomenclature Practice
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
Al(NO2)3
aluminum oxide
calcium nitrate
CrSO4
Cu(NO3)2
CuF
FeO
HgCl2
iron(II) sulfide
iron(III) sulfide
magnesium chloride
manganese(II) bromide
MgSO3
sodium sulfate
strontium phosphide
ZnO
aluminum nitrite
Al2O3
Ca(NO3)3
chromium(II) sulfate
copper(II) nitrate
copper(I) fluoride
iron(II) oxide
mercury(II) chloride
FeS
Fe2S3
MgCl2
MnBr2
magnesium sulfite
Na2SO4
Sr3P2
zinc oxide
150
Supplement 15 Lewis Structures
There are several ways to symbolize molecules. Each way tells us certain things about the
molecule; which we use depends on the features we want to show or emphasize. If we want to
show which atoms are bonded together and the location of the outer-shell electrons in a molecule,
then we would use a Lewis structure. Note that Lewis structures do not show the actual threedimensional shape of the molecule. First let’s look in more detail at what a Lewis structure tells us,
then we’ll look at how to draw simple Lewis structures “from scratch”.
Each element symbol in a Lewis structure represents one atom of that element. A single line
between two atom symbols represents a pair (two) electrons shared by those atoms. A double line
represents two pairs of shared electrons and a triple line represents three pairs of shared electrons.
A single line may also represent a pair of unshared electrons, in which case it will be with one atom
symbol rather than between two atom symbols; alternately, an unshared pair of electrons can be
represented by two dots. Double and triple lines are not used for unshared electrons.
Let’s look at two examples:
O
H C H
H
C N
In each Lewis structure above, the H’s represent hydrogen atoms, the O represents an
oxygen atom, the C represents a carbon atom and the N represents a nitrogen atom. In the structure
on the left, the carbon atom shares one pair of electrons with each H atom and two pair of electrons
with the oxygen atom. Also, the oxygen atom has two pair of electrons that are not shared with
other atoms. In the structure on the right, the carbon atom shares one pair of electrons with the
hydrogen atom and three pair of electrons with the nitrogen atom. The nitrogen atom also has one
pair of electrons that is not shared with other atoms.
We can also draw Lewis structures for polyatomic ions. Lines and atomic symbols mean the
same thing as with molecules, but we also indicate that the ion is charged by writing its algebraic
charge to the top right of the structure. For example, the Lewis structure of the cyanide ion is
1-
C N
The C represents an atom of carbon, the N an atom of nitrogen. Each of the single lines
beside the C and N represent a pair of unshared electrons. The triple line between the C and N
represents three pairs of electrons that are shared by the C and N. The -1 at the upper right is the
charge on the cyanide ion.
Since we’ve looked at what a Lewis structure means, let’s consider how to draw them for
molecules. There are a few basic rules below for drawing Lewis structures. The first involves
drawing a framework, and deserves some discussion. Drawing a framework involves deciding
which atom(s) are in the center and which are bonded to one of the central atoms. Making this
decision takes some practice, and creates some difficulty for beginning students.
(a) Organic formulas are sometimes given in a way to indicate some of the structure. In this
type of formula, the symbols that follow a central atom are bonded to it. For example, writing the
framework for CH2CH2 involves starting with a C and bonding 2 H's to it and then another C with
its own 2 H's.
H
H
H
C
C
151
H
This type of formula is recognizable because the symbols of some types of atoms will repeat;
CH2CH2 could have been written as C2H4 to tell there were 2 C and 4 H, but C2H4 does not
indicate structure at all. The fact that C is written twice in CH2CH2 says something about structure
is indicated.
(b) If you are not given a formula that tells something about the structure, you will have to
decide which is the central atom(s) and which are bonded to it. Often the most symmetrical
arrangement of atoms is correct. Neither hydrogen nor a metal atom will ever be a central atom for
Lewis structures we do in 101, and halogens will be not be central except in anions with only a
halogen and oxygen (such as ClO4–).
(c) The more pairs an atom is willing to share, the more likely it is to be central. Nonmetal
atoms farther from F on the periodic table are generally more willing to share than those closer to F.
Carbon atoms are often central, and may often be bonded to each other as well. Nitrogen, oxygen
and sulfur may be central atoms, particularly in molecules involve only carbon hydrogen or
halogens as their other atoms.
(d) It will help to remember the following:
H shares one pair of electrons and has zero nonbonding pairs
C tends to share 4 pairs of electrons
N tends to share 3 pairs of electrons and have one pair of nonbonding electrons
O and S tend to share 2 pairs of electrons and have 2 pairs of nonbonding electrons
Halogens (F, Cl, Br, I) tend to share 1 pair of electrons and have 3 nonbonding pairs.
Steps for Drawing Lewis Structures
1. Connect the central atoms to each other using either a line or two dots to represent a pair
of electrons. Then connect the rest of the atoms to the central atoms, again using either a line or
two dots to represent a pair of electrons. Often a more symmetrical arrangement is the correct one.
2. Add up the total number of valence electrons contributed to the Lewis structure by the
individual atoms. If the formula is of an ion, then a.) add a number of electrons equal to the
negative charge or b.) subtract a number of electrons equal to the positive charge. Now divide the
number of valence electrons by two to get the total number of electron pairs in the Lewis structure.
3. Find how many pairs of electrons you have left to distribute by subtracting the number
you have used already (in step 1) from the number available (from step 2). If you used all the
electron pairs, the structure is done. Usually at this point you will have electrons left over, and will
need to do step 4.
4. For any atom not having a full valence shell (one pair for H or four pair for other atoms),
use the extra electron pairs you determined in step 3 to finish the Lewis structure as follows:
(a) Add pairs of electrons to the outer atoms first to complete those atoms' valence shells of
electrons. Spread pairs of electrons out among the outer atoms, but don’t do the central atom yet.
After all the outer atoms get four pair (or one pair for H), then finish with electrons on the central
atom if needed. DO NOT LET THE TOTAL NUMBER OF ELECTRON PAIRS GO OVER THE
TOTAL YOU CALCULATED IN STEP 2.
(b) If you run out of electrons before all atoms have completed valence shells, modify the
Lewis structure by moving an unshared pair of electrons from an outer atom to make a multiple
bond between the center atom and an outer atom.
5. Never have more than four pair of electrons about any element whose symbol is in row
one or two on the periodic table.
6. Go back and check your work. (a) Did you get the correct central atom and bond the
other atoms around it? (b) Does your Lewis structure use exactly the number of electron pairs
allowed? (c) Is each atom's valence shell filled?
152
Supplement 16 Polarity of Molecules
When two atoms share electrons, they often do not do so equally. One atom may pull the
electrons toward itself more than the other can. If the shared electrons spend more time around this
atom, then it will have a slight (but less than one) negative charge. The atom at the other end of
such a bond must have a slight positive charge since it was neutral until the first atom hogged the
electrons. We say that a covalent bond with a positive and negative end is polar. To show that the
charge is less than 1, we often write a lower case Greek delta, , instead of a number with the
charge sign.
For example, if we look at HCl, the shared electrons spend more time around the Cl than
around the H. As a result of this, the Cl end of the molecule has a slight negative charge and the H
atom has a slight positive charge. If we could look at where the electrons are every few instants
and mark a dot at that place, after a period of time we might have a diagram that looks like the
following:
+
H
Cl
-
It is time consuming to draw this type of picture all the time, so instead we can indicate that a bond
is polar by placing the partial charge ’s at the atom symbol or by drawing an arrow from the
positive atom to the negative one.


H- Cl
H- Cl
After we learn to tell which atoms are negative we won’t always even mark the charges, but
we will know they are there if the atoms are of two different types (except if they are C and H).
Which atom is negative? To decide which atom is negative, we must know which attracts
the shared electrons more (= more electronegative). This means we must compare the two atoms
rather than just looking at one. Atoms whose symbols are closer to the top of the periodic table
attract shared electrons more than those with symbols lower. Atoms whose symbols are more to the
right attract shared electrons more than those whose symbols are to the left. In using these rules, we
usually ignore the noble gases in the last column since they rarely ever share electrons with
anything.
If I look at the H–Cl and the H–F bonds, the F attracts the electrons it shares with the H
more than the Cl attracts the electrons it shares with the H. The H pulls the electrons the opposite
way with the same strength in the two cases. This means the F has more of a negative charge than
the Cl, so the H–F bond is more polar than the H–Cl bond. Let’s look at one other example.
Suppose we consider a P–O bond and compare it to a Si–F bond. The Si and F are quite far apart on
the periodic table, so there is more difference in how they attract their shared electrons than there is
with P–O. The higher the difference in electronegativity, the more polar the bond--so Si–F is more
polar than P–O.
You might think of this attraction comparing as though it is a tug of war. Atoms that are
more electronegative are stronger pullers than less electronegative atoms, so the stronger ones
“win” by hogging the electrons more of the time. (But not all the time).
153
Polar molecules
If a molecule is negative on one end and positive on the other, we call it polar. Determining
whether a molecule is polar is not always as easy as it is with a bond, though. Having polar bonds is
not sufficient to make a molecule polar, because sometimes the effects of the partial positives and
negatives cancel. This happens when the arrangement of the charged atoms is such that the center
of positive charges is at the same place as the center of negative charges. If a molecule has no polar
bonds or if the centers of the positive and negative charges are at the same place then it is nonpolar.
To determine if the centers of charges are in the same place, we have to look at the geometry of the
molecule. Let’s look at some examples.
Examples of determining molecule polarity
CO2
The Lewis structure of CO is O=C=O . Because there are two directions that the electrons
2
around the C can point, they will be 180o apart, and the three atoms will lie along a straight line.
O
C
O
. The symbol for the O atoms is closer to the top right of the periodic table than
that of C, so the O atom is more electronegative and will have a slight negative charge while the C
 

C
O
has a slight positive charge. O
. While the bonds are polar, we must look closer to
see if the molecule is polar or not. There is only one atom that is positive (the C), so the geometric
center of positive charge is on the C. The geometric center of negative charge is half-way between
the two O atoms, so the center of neg. charge is also on the C. This means the effects of the charges
on other molecules close by will cancel and the molecule acts as though it has no charges at all--it is
nonpolar.
H2O
_
H- O_ H
If we draw the Lewis structure of a water molecule, we get
. There are four groups
of electrons around the O, so these groups arrange in a tetrahedral shape. If we draw this in 3
dimensions, it will be easiest if we pick the two bonds to be in the plane of the paper and let the
unshared pairs of electrons point towards us and away from us. Since we don’t indicate directions
of unshared pairs on the paper, we can just draw shape A below. Once we’ve determined the
directions of the bonds, we can use the sketch of the molecule shown in B to determine
polarity/nonpolarity. Note that we can look only at the bonds to determine whether a molecule is
polar, but they must be pointing in the correct directions to do so!
H
H
O
H
H
O
B. molecule without
A. molecule with
showing lone pairs
lone pairs shown
(but their effect on
the shape is still there!)
If we look at structure B, we can see that there are two polar bonds, each of them an O-H bond. The
O is more electronegative, so it has a negative charge and the H’s have a positive charge.
154
+
H
H
+
O
The center of negative charge is on the O atom. The center of positive charge is half-way between
the H atoms; this is not on an atom at all, but is in the middle of space. It is not at the same place as
the center of positive charge, though, so the molecule is polar.
Another way to visualize polarity: Suppose we make a model of the molecule and pull on
the negative atoms in a direction away from the positive atoms that are in the same bond. Also,
let’s suppose we pull harder for more polar bonds. If the pulling would move the model, then the
molecule is polar. If the pulling would not move the model, then it is nonpolar. We can use the two
examples above to illustrate the “pulling” rule. We will draw arrows in the direction of the pull.
H
H
O
O
C
O
Pulling on the water model would move it down the paper. Pulling on the CO2 model wouldn’t
move it as the pulls would cancel.
A molecule is nonpolar if it has only C–H and C–C bonds as then it has no polar bonds at
all. Any molecule without any polar bonds is a nonpolar molecule.
A “rule of thumb”: If a molecule has no lone pairs on its center atom and all the bonds from that
atom are alike (same number of shared electrons and to the same type of atom), then the molecule is
nonpolar. Example: CCl4.
Uses of polarity:
(a) To predict solubilities of substances
(b) To rank melting points and boiling points
Practice: Draw 3-D structures and decide if each of the following is polar or nonpolar: (answers
below): CH4, COH2, CCl2H2, NH3, CBr4.
CH4, nonpolar; COH2, polar; CCl2H2, polar; NH3, polar; CBr4, nonpolar
155
Supplement 17: Nomenclature of Binary Molecular Compounds That Are Not Neutral Acids
Naming compounds is a matter of following rules. It is not difficult, but it does require that
you recall information, that you apply the rules, and that you go through the several steps involved
in arriving at a correct name.
Many systems for naming chemicals have been used over time. Two procedures are
commonly used for inorganic compounds; one for ionic compounds and another for binary
molecular compounds. This sheet outlines procedures for naming binary molecular compounds. A
separate sheet describes procedures for naming ionic compounds.
Without checking the properties of a substance in the laboratory, it is not possible to say
whether a compound is ionic or molecular. However, it is generally true that compounds between
two nonmetals (elements to the right of the "stair steps" in the periodic table) are molecular. We
will assume this to be true and use the system of nomenclature described here to name all
compounds containing only nonmetals.
Binary molecular compounds are named by naming the elements in the compound and using
prefixes to indicate the number of atoms of each element present. (If a single atom of an element is
present, the mono- prefix is often omitted for the first element named.) The prefixes used are on the
following table:
1
mono
6
hexa
2
di
7
hepta
3
tri
8
octa
4
tetra
9
nona
5
penta
10
deca
The element farther to the left and/or lower on the periodic table is written (and named) first. In
binary molecular compounds the name of the second element is modified so that the name ends in ide. The following checklist outlines the steps:
1.
2.
3.
4.
What are the elements in the compound?
Which element is farther to the left or lower on the periodic table?
How many atoms of each element are in one molecule?
What prefixes are needed to indicated the number of atoms?
Example: What is the name of a compound containing two atoms of oxygen and one atom of
carbon in each molecule?
1.
2.
3.
4.
The elements in the compound are oxygen and carbon.
Carbon is to the left of oxygen on the periodic table, so carbon is written first.
There is one atom of carbon and two atoms of oxygen.
The prefix for one is mono- and the prefix for two is di-
The name is monocarbon dioxide but the mono- prefix is usually omitted to give carbon
dioxide for the name. Notice that oxygen is changed to oxide by dropping the -ygen ending in
oxygen and then adding -ide.
Write the correct formula for each of the following, and do it this way: 1) Write the formula
for the first compound without looking at the checksheet. 2) After you have written the
formula and think you have it right, look at the checksheet and check each step to be sure you
have done it correctly. 3) Now check your answer against the answer at the end of this
handout. 4) Now write the formula for the rest of the names in the list and repeat the same
156
procedure: Write the formula first, then go over the checklist, and finally check your answer
to see if it is correct. (The number in front of the name is to locate the correct answer on the
answer sheet at the end of this document.)
11.
2.
7.
13.
9.
four atoms of chlorine bonded to one carbon atom
one atom of sulfur and three atoms of oxygen
one atom of sulfur and two atoms of hydrogen
two atoms of sulfur and one atom of carbon
six atoms of bromine and two atoms of silicon
Writing the name when you are given the formula follows the same logic:
Example: What is the name for N2O4?
1.
The elements in the compound are nitrogen and oxygen.
2.
The order of elements in the name will be the same as the order in the formula.
3.
There are two atoms of nitrogen and four atoms of oxygen.
4.
The prefix for nitrogen will be di- and the prefix for oxygen will be tetraThe name of the compound is dinitrogen tetraoxide
Following the procedures outlined above, name these binary molecular compounds.
4.
1.
12.
SiC
PCl3
SiO2
3.
14.
BrF3
N2O3
Writing the formula for a binary molecular compound when you know its name is very
simple. Ask yourself these questions:
1.
2.
What are the elements represented by the name?
How many atoms of each element are represented?
Example: What is the formula for nitrogen trihydride?
Step 1: What are the elements? "Nitrogen is clearly one. The other is less obvious. The 'tri-' prefix
is telling me how many atoms of the element are present, and the '-ide' suffix just tells me that this
is a binary compound. 'Hydr-' is the only part of the last word in the name that is identifying the
element involved. Must be 'hydrogen' since that is the only element I can think of that starts out
'hydr-'."
Step 2: How many atoms of each element are represented? "There is no prefix on 'nitrogen' so I can
assume that the formula contains a single nitrogen atom.
The formula must by NH3 (Which is commonly called ammonia)
Following the procedures outlined above, write formulas these binary molecular compounds.
6.
tetraphosphorus heptasulfide
5.
tribromine octoxide
8.
dichlorine heptoxide
15.
disilicon hexachloride
10.
nitrogen monoxide
157
Answers for Binary molecular Nomenclature Worksheet
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
phosphorus trichloride
sulfur trioxide
bromine trifluoride
silicon carbide
Br3O8
P4S7
dihydrogen sulfide
Cl2O7
disilicon hexabromide
NO
carbon tetrachloride
silicon dioxide
carbon disulfide
dinitrogen trioxide
Si2Cl6
158
Supplement 18: Chemical Equations
What is a Chemical Equation ?
A chemical equation is a method of representing chemical changes accurately. An example
is shown below.
2 NaOH + CO2 Na2CO3 + H2O
On the left of the arrow are the formulas for the substances that undergo the change. These
are called reactants. The arrow means that a chemical change has occurred. On the right of the
arrow are the formulas for the substances that are the result of the chemical change. These
substances are called products.
We could also represent a physical change using the same arrows as with the evaporation of
water shown below.
H2O(l) H2O(g)
Of course we recognize this is not a chemical change, in spite of the use of the arrow, because the
same formula unit is present on each side of the arrow.
Another example of a chemical equation is shown below.
CH4 + 2 O2 CO2 + 2 H2O
This equation represents the change that occurs when natural gas – which is mainly
methane, CH4 – is burned in air. The reactants are methane and oxygen. The products are carbon
dioxide and water. We could say in words that for every molecule of CH4 that reacts, two
molecules of oxygen react to produce one molecule of CO2 and two molecules of water.
The main purpose of a chemical equation is to accurately represent a chemical change or a
chemical reaction. The use of symbols is probably much clearer than the use of words. If we
interpret the chemical equation on a submicroscopic level, then we see each set of symbols
representing a substance as either a single molecule or as a formula unit (for non-molecular
compounds). However, you must be aware that a chemical equation can represent the change of
any amount of substances. We could just as well say that the equation above also represents the
reaction of 1000 molecules of CH4 with 2000 molecules of O2 to yield 1000 molecules of CO2 and
2000 molecules of H2O. No matter how much of a reaction occurs, we always will use the same
chemical equation to represent it. We will simply infer that the reaction on the single molecule
scale occurs enough times so that the larger scale reaction of more molecules can take place. Still,
no matter how much reaction occurs, the same chemical equation is used to represent the change. It
is more accurate to state the equation in words as: For every molecule of CH4 that reacts, two
molecules of O2 react resulting in the formation of one molecule of CO2 and two molecules of
H2O.
Skeletons and Equation Balancing
It is common that a chemist will know which substances are used as reactants and which
substances are formed as products. When this occurs, he may write a skeleton that shows the
reactants and products separated by arrows and "+" signs as in a chemical equation. In our example
above, he might know that CH4 reacts with O2 to produce CO2 and H2O. In this case, he would
write a skeleton as below.
skeleton:
CH4 + O2 CO2 + H2O
You can recognize this as a skeleton because, while it may represent the substances going
into and coming out of the reaction well, it does not satisfy the principle of conservation of atoms.
For instance, there are 4 atoms of H shown on the left and only two shown to the right. This cannot
be an accurate representation of any change because we believe that atoms cannot be destroyed or
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created in any chemical reaction. For this reason, we must go through the process of balancing the
skeleton to make it represent the equation accurately. Since there are 4 H atoms on the left, there
must be 4 H atoms on the right so we insert a 2 IN FRONT of the H2O.
Step 1:
CH4 + O2 CO2 + 2 H2O
As soon as this has been done, we see that there are a total of 4 atoms of O shown on the right (two
in CO2 and two from two molecules of H2O). Since O atoms cannot be created, we must have a
total of 4 O atoms in the reactants so we place a 2 IN FRONT of O2 as shown below.
Step 2:
CH4 + 2 O2 CO2 + 2 H2O
We now have a chemical equation. There are 1 atom of C, 4 atoms of H, and 4 atoms of O
shown on each side of the equation. Indeed, this equality in the number of atoms of each kind on
each side of the equation is how the name "equation" came to be used for this representation of
chemical reactions.
We can summarize this section by saying that it is important to know the actual substances
that react and are produced in a chemical reaction. These are represented by a skeleton. However.
we do not have an accurate representation of the reaction until we have written a chemical equation
in which the number of atoms of each element involved in the process is shown to be the same as
the number of atoms coming out of the process.
The Use of Chemical Equations to Find Weights of Substances
So far, a chemical equation is mainly an aid in visualizing what happens in a chemical
reaction. It can be seen to apply for just a few molecules or formula units. However, a chemical
equation can be used to predict the amounts of substances that react and the amounts of substances
that are produced on a laboratory scale or even on an industrial scale. This use of the chemical
equation requires that we have a connection between the numbers of single units of a substance and
weight. This connection is made by using the weight of a mole of particles. In rough language, the
mole is the chemist's "dozen", his counting unit. In order to bring the scale of chemical processes to
the laboratory scale, he will refer to the reaction we have described above as: for every mole of
CH4 that reacts, 2 moles of O2 react to form 1 mole of CO2 and 2 moles of H2O. Using moles this
way is just as true as using "dozen" or "thousands" or any other counting unit.
You should recall that the mole is 6.0 x 1023 units. This is a convenient counting unit for
the very tiny particles that are molecules and formula units. The method of going from numbers of
particles (moles of particles) is to use the molar mass which is the mass in grams of one mole of
particles. For example, a mole of CH4 molecules contains 1 mole of C atoms and 4 moles of H
atoms. Since these weigh 12.01 g/mole C atoms and 4.03 g/4 moles of H atoms, one mole of CH4
molecules weighs 12.01 + 4.03 = 16.04 g/mole CH4 molecules. This number, the molar mass of
CH4 can be used to calculate the number of CH4 molecules present in any sample of CH4 by
simple division. If we had a sample of 4.01 g of CH4, the number of CH4 molecules present would
be:
4.01 g CH4/16.04 g CH4/mole CH4 = 0.250 mole CH4
You may not recognize this as a number of molecules but since a mole is a counting unit it is
indeed a number. This is very like saying that we have 0.250 gross of nails. While it is not a simple
number, it is easy to convert the number of moles to a simple number of particles using Avogadro's
number.
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The aim of all this is to say that a chemical equation can be used to show a relation between
the weights of substances reacting or produced in a chemical equation. We can show this by
writing our old equation for the burning of methane with various numbers of particles under it and
then by writing it with the corresponding weights of substances underneath it.
CH4 + 2 O2  CO2 + 2 H2O
1 mole 2 mole
1 mole 2 mole
or
100 mol 200 mol 100 mol 200 mol
or
0.01 mol 0.02 mol 0.01 mol 0.02 mol
If we write the weights corresponding to these numbers of particles,
16. g
64. g
44 g
36 g
or
1600 g 6400 g
4400 g 3600 g
or
0.016 g 0 .064 g
0.044g 0.036 g
We can summarize this section by saying that a chemical equation shows the relationship
between the NUMBERS of particles reacting and produced. It can be used to show the
relationships between weights of substances reacting and produced if the molar mass, the weight of
a mole of substance, is used.
Limiting Reagents
You may get a somewhat mistaken impression from the above that you can only mix certain
magical amounts of substances in order to get a reaction to occur. This is not true. You should read
a chemical equation as saying that: TO THE EXTENT A REACTION OCCURS, what is
represented is what happens. Regardless of how substances are mixed, the equation shows what can
occur. The reaction of methane we have shown is a good example. If we were to burn 0.16 g of
methane in a whole lab full of air that contained 60,000 g of oxygen, O2, then only 0.16 g of
methane and only 0.64 g of oxygen would react.
or
CH4 + 2 O2  CO2 + 2 H2O
0.01 mol 0.02 mol 0.01 mol 0.02 mol
0.16 g 0.64 g
0.44 g
0.36 g
It is easy to see why this is true. Once all the methane has been used in the equation as
written, there is simply no more to react with the huge excess of oxygen present and the reaction
stops. This represents the common situation where one reagent limits the amount of reaction simply
by being consumed as expected before the other reagent can be used. The reagent that is
completely consumed is called the limiting reagent and the other substances that remain after the
limiting reagent is consumed are said to be present in excess. The reaction must simply stop once
there is no more of one of the reactants present. Most reactions in the laboratory and all reactions in
industry are carried out this way. In industry, it makes sense to use an excess of a cheaper
substance to make sure the more expensive substance is used entirely.
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Supplement 19: Polarity and Solubility of Substances
Table salt, primarily NaCl, dissolves well in water, yet CaCO3 dissolves very little in water
and causes bathtub ring. Painters can clean an “oil-based” paint from brushes with turpentine, but
can’t clean them with water. Why do some substances dissolve in water while others dissolve better
in nonaqueous solvents? Having an answer to this question would let us pick which solvents to use
for particular substances. To get an answer, we need to look at what goes on at an atomic scale
when something dissolves.
The Process of Solution Formation
We must examine three forces: those between solvent molecules, those between solute
molecules, and those between solvent and solute molecules. (The same ideas applied to molecules
below also apply to ions of ionic substances and to atoms of monoatomic substances.)
Suppose we look at an example in which the solute is a solid and the solvent is a liquid. Remember
that solids and liquids remain in a condensed state because their submicroscopic particles attract
each other enough to keep them from “flying off” into the space above their top surfaces.
The solute molecules are held together by attractions between the many
molecules,
while the solvent molecules,
, are held together by their attractions to each other. In order for a
solution to form, the submicroscopic particles must get all mixed together:
solution
Hence, for a solution to form, the solute molecules must be separated from other solute
molecules and the solvent molecules must be separated from each other. They can be separated if
there is an attraction between molecules of the solute and solvent that is stronger than or similar in
strength to attractions within the solute and within the solvent.
We might consider this in terms of groups of boys and girls. When children are in the lower
grades, they tend to form social groups by gender--the boys hang out together and so do the girls,
but they don’t mix (they’re insoluble?). When they reach an age at which there is more attraction
between the boys and the girls than within each gender group, then they mix together.
To return to molecules, if the attraction between the solute particles is higher than between
the solvent and solute molecules, then the solute particles stay together. Or, if the attraction
between solvent molecules is higher than between solute and solvent molecules, then the “solvent”
particles stay together and there is no mixing. In both of these cases, extremely little solute would
dissolve. If, however, the attraction between solvent molecules and solute molecules is higher than
solute-to-solute or solvent-to-solvent attractions, then a solution can form.
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The pictures below illustrate what happens as a solid dissolves in a liquid. Liquid molecules
partially surround a solid molecule at the surface (picture 1) and carry the solid particle away from
the rest of the solid (picture 2). When this occurs, other solid molecules are exposed to the liquid
and they, too, can be surrounded (3) and carried away(4) . This process continues until all of the
solid molecules are mixed in among the liquid molecules.
(1)
(2)
(3)
(4)
Polarity and Predictions of Solubility
We can make predictions about the attractions between molecules based on their polarity.
Polar molecules attract other polar molecules, while there is less attraction between polar and
nonpolar molecules. Thus molecules of a polar solute would have more attraction for molecules of
a polar solvent than those in a nonpolar solvent. A polar solute is more likely to be dissolved by a
polar solvent than by a nonpolar solvent. If the solute molecules attract each other very strongly,
then even polar solvent molecules may not be able to pull the solid molecules away from each
other. Thus, a polar solute may or may not dissolve in a polar solvent, but it is very unlikely that it
will dissolve in a nonpolar solvent.
In a different case, if both solute and solvent molecules are nonpolar, then the attractions
between solute molecules and the solvent molecules are often about the same strength as those
within the solute and those within the solvent, and a mixture can form. If the solid molecules are
nonpolar and the liquid molecules are polar, then separating the liquid molecules is too difficult, and
no solution forms. Thus nonpolar substances tend to dissolve better in nonpolar solvents and not in
polar solvents.
Let’s apply our rules to pick some solvents. Suppose you want to make a dilute solution of
wood alcohol, which is CH3OH. Should you use water or a nonpolar solvent, such as hexane,
C6H14? The 3-D structure of a molecule of CH3OH is illustrated by the following.
H
C H

H O H
The C–H bonds are nonpolar. The C–O bond is polar, and so is the –O–H bond, so the molecule is
polar. You should choose water as a solvent over a nonpolar solvent.
We can also use our rules to deduce polarities from solubilities if the substance dissolves.
For example, bubble gum dissolves in the mixture of nonpolar molecules in gasoline-- this implies
the molecules in bubble gum are likely to be predominately nonpolar.
Question: Which of the following would be most soluble in water? I2, CH2CH2, CH3Cl?
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Supplement 20: Osmosis
Osmosis is the movement of a substance through a semipermeable membrane from a region
of high concentration to one of low concentration. Generally this occurs in the presence of other
substances that do not move through the membrane. We can see how this movement would occur if
we examine how we would expect the submicroscopic particles to behave.
A semipermeable membrane is one that allows some things to move through it while other
things cannot move through. If we examine a cross -section at a molecular level, we might see the
following.
A semipermeable membrane has pores which allow some molecules through and not others.
Suppose in this example the small molecules can go through, but the larger ones cannot. The
molecules are constantly moving and the smaller molecules sometimes happen to be moving in a
direction to pass through the pore in the membrane. In the solution with a higher concentration of
small molecules, there are more of them close to the pores and so more can pass through the pores.
There are also some of the smaller molecules moving through from the side of lower concentration,
but fewer move from low to high concentration than move the other way. This means there is an
overall movement of the molecules that can go through the membrane from a region of their high
concentration to one of their low concentration. Nothing really happens to the molecules that
cannot go through the membrane.
Osmosis occurs frequently in everyday life. Pickles are smaller than the cucumbers from
which they are made because water in the cucumbers moves through the membrane into the brine in
which the pickles are made. In this case, the water starts at high concentration in the cucumber and
at lower concentration in the brine. Celery and many other vegetables that have become limp
during storage can be re-crisped if they are placed into water. In these cases, the water is at higher
concentration outside the vegetable and a lower concentration inside, so the water moves across the
vegetable membranes into the vegetable. During kidney dialysis, there is a net movement of waste
molecules move across an artificial membrane from a region of high concentration to one of low
concentration (of the waste molecules). In order to prevent the osmosis of useful molecules, their
concentration is kept the same on both sides of the membrane.
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Supplement 21: An Outline of Acids and Bases
Definitions :
An acid gives up a H+.
A base accepts a H+ .
In a neutralization reaction, a H+ is transferred from the acid to the base.
Some of the common reactions of acids and bases:
When an acid is placed in water, H+ (=H3O+) ions are formed.
HCl + H2O H3O+ + Cl–
When a base is placed in water, OH– ions are formed or go into solution.
NH3 + H2OOH– + NH4+
NaOH(s) + H2OOH– (aq) + Na+ (aq)
H+ is transferred from the acid to a base or a part of the base (= a neutralization reaction)
HCl + NH3  NH4Cl
HNO3 + NaOHH2O + NaNO3
Recognizing the formula of an acid:
(1) Compounds that do not contain C and have a H written in front of their formulas are generally
acids . For example, H2S is an acid.
(2) Compounds that have CO2H or COOH in their formulas are acids, e.g., CH3COOH is acetic
acid.
(3) A positive ion that contains a N atom and has a H bound to the N is frequently an acid.
NH4Cl is actually two ions, NH4+ and Cl–. The ammonium ion is an acid.
H
H N
H
+
H
Cl -
Recognizing the formula of a base:
(1) An ionic compound that contains a hydroxide (OH–) ion. NaOH, Mg(OH)2
(2) A neutral compound that contains a N atom.
H H
A H+ can bond to the lone pair of electrons,
CH3NH2 is
| |
forming
H- C- N- H
H
.
so the lone pair on the N allows the molecule to accept an H+ (which is the definition of base).
165
,
(3) An ionic compound that does not contain the following: Cl–, Br–, I–, NO3–, SO42–,ClO4–, or H+
Neutral compounds: Neither donate nor accept hydrogen ions when placed with an acid or a base.
Recognizing Neutral compounds: Ionic compounds that do not contain hydrogen ion or hydroxide
ion and do contain Cl–, Br–, I–, NO3–, SO42– or ClO4– are neutral. Molecular compounds that do
not ionize in water are also neutral.
Strong and Weak Acids--Definitions:
If all of the molecules of a compound give up one or more hydrogen ions to water when the
compound is put into water then it is a strong acid.
Recognizing the formula of a strong acid: Memorize these strong acids!
HNO3 , H2SO4 , HCl, HBr, HI , HClO4
If only part of the molecules of a compound give up one or more hydrogen ions to water when the
compound is put into water then it is a weak acid.
Recognizing a weak acid: all acids that are not strong are weak.
Strong and Weak Bases:
If all of the molecules of a compound react to produce a hydroxide ion when placed in water the
compound is a strong base.
Recognizing strong bases: Soluble ionic compounds that contain hydroxide ion (cations from
Periodic columns 1 or the bottom of 2) are strong bases.
If only part of the molecules of a compound react to produce hydroxide ion or accept hydrogen ion
from water then it is a weak base.
Recognizing weak bases: If a base is not a soluble hydroxide, it is weak.
Practice: Examples to classify:
NaCl
H3PO4
RbOH
NH3
HBr
CH3COOH
CH3NH3+
KNO3
CH3NH2
NH4Cl
CaSO4
NH4Br
HNO2
NaOH
Write the chemical equation for the reaction of each of the above with water.
Write the chemical equation for the reaction of each base above with HCl.
Write the chemical equation for the reaction of each acid above with NaOH.
Other “Stuff” to remember about acids and bases:
Water solutions of acidic compounds have a pH less than 7
Water solutions of basic compounds have a pH greater than 7
Water solutions of neutral compounds have a pH equal to 7
While a compound may be called an acid, in water solutions it is the H+ that causes the test results
we generally associate with acids (for example, turning litmus colors, reacting with bicarbonate to
166
produce gas bubbles). In water solutions it is the OH– that causes the test results we generally
associate with bases.
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Supplement 22: Buffers
Buffers are solutions that resist (but not totally prevent) changes in pH when small amounts
of acid and base are added due. This occurs due to the presence of both weak acid and weak base.
Buffers are quite important biologically as they help maintain blood pH (in this case the weak acid
is HCO3– and the weak base is CO32–). Buffers are also used to control the pH during reactions in
the lab.
Acid Addition
If a small amount of acid is added, the weak base present reacts with the added acid,
preventing much change in H+ molarity. With much less change in H+ molarity, the pH doesn’t
change nearly as much as in an unbuffered solution.
For example, with blood, the reaction when acid is added is CO32– + H+  HCO3–
Base Addition
If a small amount of base is added to a buffered solution, the pH doesn’t change nearly as
much as in an unbuffered solution. In the buffer, the added base reacts with the weak acid. In an
unbuffered solution, addition of base makes a higher OH– molarity, thus changing the pH.
To return to the blood example, when base is added the reaction is
HCO3– + OH–  H2O + CO32–
Because the added OH– reacts, its molarity doesn’t change much and the pH is constant.
What kinds of solutions make buffers?
Any solution which contains a weak base and a weak acid is a buffer if the molarity of weak
acid divided by the molarity of weak base is between 0.1 and 10. If the ratio is less than 0.1 (or
more than 10), the amount of acid (base) will be too small to resist much change in pH as the weak
acid (weak base) gets used up too quickly.
As a practical matter, buffers are usually made of acid-base pairs that differ from each other
by only one H+. Common examples are solutions containing HCO3– + CO32–-, solutions
containing PO43– + HPO42–, solutions containing HPO42–- + H2PO4– and solutions containing
NH4+ + NH3.
Q1: For each of the solutions listed, write the reaction equation(s) for what happens if (a) a small
amount of acid is added and (b) a small amount of base is added. (a) a solution containing HNO2
and NaNO2 (b) a solution containing HC2H3O2 and KC2H3O2 .
Q2: What would you use with HF to make a buffer? with CH3NH2? with NaH2PO4?
pH of Buffers
The pH of a buffer can be calculated if the molarity of the weak acid and the weak base are
known along with a constant called the Ka or pKa of the weak acid (this is typically looked up in a
handbook).
pH = pKa + log [weak base molarity/weak acid molarity]
Example: Calculate the pH of a buffer that is 0.10 M in acetic acid and 0.20 M in sodium acetate.
The pKa of acetic acid is 4.7.
pH = 4.7 + log [0.20/0.10] = 4.7+0.3=5.0
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Supplement 23: Predicting Whether a Reaction Will Be Spontaneous
A spontaneous reaction, when over, produces a fair amount of products for the amount of
reactants used. (Note, this is not the same as “happens by itself”!) A nonspontaneous reaction, when
over, produces few products for the amount of reactants used. Examples of spontaneous reactions are
the melting of ice when it is placed in surroundings that have a temperature above 0 oC and the change
of egg white from almost clear and gelatinous to a white solid when it is heated. Examples of
nonspontaneous reactions are the dissolution of bathtub ring or oil in plain water. Classifying a reaction
as nonspontaneous does not mean there is no product, just that there is little product compared to the
amount of reactant. (A little bit of the oil may dissolve in the water).
If no reactants or products are removed or added and conditions are left the same, reactions
reach a point where the amount of product present no longer changes. The amount of time needed
to reach this situation has nothing to do with whether a reaction is spontaneous or nonspontaneous.
Some spontaneous reactions are slow, some are fast.
What are some features of spontaneous reactions? There are two features of reactions
which, when taken together, seem to determine whether a reaction is spontaneous. One feature is
the relative amounts of disorder of the particles in the reactants and products, the other is whether
heat is absorbed or released in the reaction.
First let’s talk about what I mean by disorder and how to tell if heat is absorbed or released.
When I say something has larger disorder, I mean that its submicroscopic particles are more
jumbled or scattered. For example, the submicroscopic particles of a gas are more disordered than
those of a liquid because those in a gas are more scattered about. A liquid is more disordered than a
solid because its component molecules are more jumbled and have no regular arrangement.
Not surprisingly, if a reaction gives off heat, the temperature of its surroundings increases.
When a reaction absorbs heat, the temperature of its surroundings decreases. If a reaction occurs in
a water solution, then the surroundings are the water and the container. Thus when a reaction gives
off heat, the water gets warmer and when a reaction absorbs heat, the water gets cooler. Now, how
do these things help us to predict whether a reaction will be spontaneous?
If the amount of disorder in the submicroscopic particles of the products is larger than the
amount of disorder in the particles of the reactants, then the reaction is more likely to be
spontaneous, though we cannot be sure. If the reactants are more disordered, then the reaction is
more likely to be nonspontaneous, though we cannot be sure.
If the reaction gives off heat, then it is more likely to be spontaneous, though, again, we cannot
be sure. If the reaction absorbs heat, it is more likely to be nonspontaneous—but we can’t be sure.
In the wake of all this uncertainty, all is not lost. It is the combination of the relative
disorder and heat change which allows us to predict with certainty whether a reaction will be
spontaneous or nonspontaneous. If a reaction’s products are more disordered than its reactants and
it gives off energy, then the reaction will always be spontaneous. If a reaction’s reactants are more
disordered and the reaction absorbs energy, then the reaction will not be spontaneous. If one feature
tends to promote spontaneity and another feature to promote nonspontaneity, then the temperature
will determine whether the reaction is spontaneous. We can summarize this information as follows:
Product Or Reactant Disorder
Is Greater?
product
reactant
product
reactant
Energy Absorbed Or Released?
Is Reaction Spontaneous?
released
absorbed
absorbed
released
yes
no
sometimes ( at high T)
sometimes ( at low T)
We can also predict spontaneity based on the change in a quantity called the Gibb’s Free
Energy, or just Free Energy. It is given the symbol G and it incorporates the effects of disorder
and heat. If the G is negative, the reaction will be spontaneous and if the G is positive, the
reaction is not spontaneous.
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Supplement 24: Risk Management
Most activities involve some risk. When you drive a car, you risk having an accident. When you go
swimming, you risk getting cramps and drowning. The question we must ask ourselves is not “How do we
get rid of all risk”, but “How do we decide how much risk is acceptable?” In deciding , we should be aware
of several things: what things affect how people perceive risk and how is risk assessed?
People’s willingness to take a risk seems to depend on several factors. One is the actual chance of
being hurt, based on factual information. An example would be something like “One person in one hundred
that participates in an activity is injured”. However, if the risk is connected with something people do not
understand, or with something not under their direct control, they find the same amount of risk less
acceptable than they would if it was due to things they understood or controlled. Also, many people are
more willing to be exposed to a risk connected with something that occurs in nature. If the report of an
event has been in the newspapers or on TV, people seem more willing to believe it could occur again soon.
We should clearly consider what is important when assessing whether a certain risk is acceptable--should a
risk be accepted or rejected only on the basis of factual data or should other factors be considered?
How is risk assessed?
Sometimes by looking at records of what has happened in the past. For example, people exposed to
cadmium waste in Japan developed Itai-Itai disease and miners exposed to high levels of radon developed
lung cancer at a higher rate than those not exposed in this way. The risk associated with the use of certain
drugs or substances is sometimes assessed is by exposing animals to the substance.
There are several problems associated with testing drugs and other substances by using animals.
Different animals respond in different ways, so man might respond differently than a rat. Most animals don’t
live as long as we do, so testing for long-term effects is difficult as the animals may die or respond
differently due to age. Even when possible, testing over a long time period is more expensive, and in the
meantime we may need the results. Seeing small effects may require the use of many animals to be sure that
an effect is really different than the behavior of the control group rather than a small variation due to
something beyond our control (for example, the DNA of a particular rat). While these are difficult questions,
we don’t want to test people, and the other alternative is sometimes to run no tests. We must decide which is
worst--results hard to interpret and possibly wrong, or no testing at all.
There are some techniques that are used to try to solve some of the above problems. In a few cases,
the substance may be tested on more than one type of animal. If it causes cancer in rats and hamsters and
rabbits, it seems more likely that it may cause cancer in humans than it would if it only caused cancer in rats.
Sometimes the type of animal used is chosen because it responds similarly to humans; for example, because
pigs digestive systems react somewhat more like those of humans than those of many other animals, they are
used to test substances likely to affect digestive systems.
To simulate the long term exposure to small doses of a substance, animals may be given very large
doses of the substance for a shorter time. The effects are then interpreted in ways that may or may not have a
sound basis: (1) If a large dose results in some effect, it is assumed that a small dose will result in an effect,
too, though not as much. (2) If a large dose creates an effect in a short time, it is generally assumed that
smaller doses would have an effect over a long time. Whether the above assumptions are true depends on
the identity of the substance and on our bodies--is there a limit below which our body can repair any damage
done? Is there an exposure small enough so it does no damage? These questions cannot be answered
without a good deal of expense and time. Sometimes results are revised after further investigations.
As a consumer, you should realize that there are ways to assess risks, but these methods are not
perfect. When you decide whether a risk is acceptable, you should be aware of the common prejudices and
not fall into the trap of reacting without thought.
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Supplement 25:
Concentrations in Percent, Parts per Thousand and Parts per Million
You are probably familiar with thinking of percent as “parts per hundred”. For example, if
you make an 80% on an exam, then you scored 80 points out of each possible 100 points. The score
represents a ratio of points obtained to points possible; there may actually have been 200 points on
the exam and you scored 160 of them, but this is still an 80 percent since 160/200 equals the same
number as 80/100.
The concentration units percent, parts per thousand and parts per million all represent a ratio
of how much of one type of substance is there to how much total sample there is. The units of “how
much” can sometimes be particles, as indicated by your text, but in practice the units are more likely
to be grams or milliliters. Until we discuss solutions in a later chapter, we will assume the units
used in percent, parts per thousand and parts per million are always grams.
Percent problems can be worked easily using factor label if we realize that percent means
the number of grams substance out of each 100 grams sample. Thus if a sample is 5 percent silver,
it means that there are 5 g of silver for each 100 g of sample. If we have double the sample, we
5 g silver
expect double the silver. Thus we can write the ratio:
and use it to solve
100 g sample
problems.
For example, if we have 345 g of sample that is 5 % silver, how many g of silver are in it?
We know the 345 g sample and we want to know g silver, so we would like to have a ratio
involving those two types of units so we can set up the problem as follows:
345 g sample g silver
=
g sample
? g silver.
We know the numbers from the 5%.
345 g sample
5 g silver
=
100 g sample
17.25 g silver.
Similarly, suppose we have silver ore that is 3.2% silver and we want to take enough of it to get 125
g of silver. The initial setup is
125 g silver
g ore
g silver
=
? g ore, but we must put the numbers in the ratio.
The 3.2% silver tells us there are 3.2 g silver per 100 g ore, so we insert
125 g silver 100 g ore
3.2 g silver
=
3906. g ore.
Parts per thousand (ppt) is a concentration unit that expresses the number of grams of a substance
per 1000 g of sample. For example, if an air sample is 5 ppt ozone, it means for each 1000 g of air
there are 5 g of ozone. If we have twice as much air, we expect twice as much ozone, so we can
5 g ozone
1000 g air
write the ratios
or
.
1000 g air
5 g ozone
171
Suppose we have 6000 g of air that is 5 ppt ozone. How many grams of ozone is in this sample?
6000 g air
5 g ozone
= 30 g ozone.
1000 g air
Parts per million (ppm), is a concentration unit that expresses the number of grams of a substance
per 1,000,000 g of sample. For example, if stream water is 200 ppm chromium, for each 1,000,000
g of stream water there are 200 g of chromium. If we double the amount of water, we’d double the
1,000,000 g water
amount of chromium, so we can write the ratio
.
200 g chromium
Suppose your well water is 200 ppm chromium. How much water would you have to drink to
ingest 1 g of chromium?
1 g chromium
1,000,000 g water
200 g chromium
=
5, 000 g water
Converting between percent, ppt and ppm.
It is easy to convert between these units if we just remember what they mean as ratios and
rephrase the questions. For example, if we wanted to know what percent chromium the 200 ppm
chromium was, we’d first decide what the percent really means. We actually want the g chromium
in 100 g of water. Thus we know the 100 g water and can use the ratio implied by 200 ppm
chromium:
200 g chromium
(100 g water)
= 0.0200 g chromium (in the 100 g water.)
1,000,000 g water
The trick is to remember that the ? g chromium to solve for is out of 100 g water and since percent
chromium means the number of g chromium in 100 g water.
Similarly, suppose we want to know how many ppt ozone are equivalent to 45 percent ozone. We
remember that ppt ozone means the number of g of ozone in 1000 g of sample. Also, we know that
45g ozone
45% ozone means
. The factor label setup is thus
100 g sample
45g ozone
1000 g sample
= 450 g ozone.
100 g sample
One more example--ppt to ppm. Suppose we had air that is 56 ppt ozone and we want to know how
many ppm ozone this is. We want to know how many g ozone in 1,000,000 g air.
1,000,000 g air 56 g ozone
1,000 g air
= 56,000 g ozone (in 1,000,000 g air) = 56,000 ppm
Practice question: If air is 200,000 ppm oxygen, how many ppt oxygen is it?
172
Supplement 26: Chemical Equilibrium
Experimentally, we observed that as a chemical reaction occurs the amounts of reactants
decrease and the amounts of products increase. This process slows down and eventually stops so
that at the end of the process, the amounts of reactants and products quit changing. This state is
called a chemical equilibrium, that is, the state in which amounts of reagents no longer changes is
called a chemical equilibrium. All chemical reactions will eventually reach a state of chemical
equilibrium. Often this state of the chemical system is an extreme one: either the only reagents
remaining in measurable amounts are products and we say the reaction is "essentially complete" or
the reaction does not seem to go at all and there are no measurable amounts of products. In some
reactions, it is found that the final state of the system is one in which measurable amounts of both
reactants and products are present. These are all possibilities at chemical equilibrium, that is, the
position of equilibrium may lie at the extremes of essentially no reaction ("equilibrium position to
the left") or essentially complete reaction ("equilibrium position to the right") or in the midrange of
some reaction.
How far towards products the reaction proceeds can be told somewhat from the size of the
reaction's equilibrium constant, K. The equilibrium constant is equal to the product molarities
raised to the power of their coefficients in the balanced equation divided by the reactant molarities
raised to the power of their coefficients in the balanced equation. Solids and pure liquids are not put
into the K expression. For example, for the reaction
M 2  M 2 2
Cu
Fe
K
2 Fe3+ + Cu (s)  Cu2+ + 2 Fe2+
2
M 3
Fe
You can see that the more product and less reactant formed, the larger K would be.
It might be supposed that the reaction is completely over and the molecules/ions of products
and reactants aren't doing anything, but it turns out not to be the case. They are involved in a
dynamic equilibrium, in which there are constant forward and backward reactions occurring at the
same rate--as product molecules react "backwards" to form reactants, at the same time reactants
react in the forward direction to re-form products. The net effect is the same total amount of
product and reactant, but which atoms are in products and which in reactants constantly switch
places.
The amounts of individual reactants and products present at equilibrium can be manipulated
to some extent by some tricks to be mentioned below. The position of equilibrium (how far towards
products the reaction goes) is effected by temperature and the concentrations of reagents used but it
is not affected by the presence of catalysts used to effect the rate of reaching equilibrium, or by
reagents that do not affect the concentration of either the reactants or products.
Changing the Position of Equilibrium
It is possible to change the amount of a reactant or product by manipulating the relative
amounts of the other materials in the reaction mixture. It is also possible to change the position of
equilibrium by changing the temperature. The way we predict the effect of these changes is by
using an idea called Le Chatelier's Principle. This idea, in simplified form, states that if a chemical
equilibrium is disturbed by changing anything in the K expression or by changing the temperature,
then the equilibrium shifts in a direction to try to compensate for the disturbance. The simplest
application of the principle can be seen using our previous example:
2 Fe3+ + Cu(s)  Cu2+ + 2 Fe2+
173
If the system is at equilibrium and more reagent Fe3+ is added, then the equilibrium will
shift to the right (producing more products) in order to decrease the molarity of Fe3+. Similarly, if
substance indicated in the K is removed from the equilibrium mixture, the reaction will go in a
direction to produce more that substance. If Fe3+ is removed, the reaction will go backwards to
make more of it (we say the equilibrium shifts to the left). If Cu2+ is removed, then the reaction
will go to make more Cu2+ and Fe2+ ( the equilibrium shifts to the right).
The effect of temperature on the position of equilibrium can be predicted if it is known
whether the reaction absorbs energy or gives off energy as it occurs. If the reaction absorbs energy
as it proceeds, then as the temperature is increased the equilibrium will shift more to the right to
offset the temperature increase. A temperature decrease for the same reaction would result in an
equilibrium shift to the left (less product formed). If the reaction gives off energy, the opposite
changes would occur: temperature increases would cause an equilibrium shift to the left and a
temperature decrease would give an equilibrium shift to the right.
The main commercial use of this one method of manipulating the position of chemical
equilibrium is that whenever a reaction is done industrially, a large excess amount of the cheaper of
the reactants is always used and the product is removed as it is formed as much as this is possible to
insure the reaction proceeds as far to the right as possible. Similarly, temperature is adjusted to
produce the most complete reaction possible. It is also important in feedback mechanisms of the
body, when any sudden production of a substance may be offset by an equilibrium shift.
174
Supplement 27: Gases--Pressure, Volume, Temperature & Amount of Gas
You have experienced the effects of pressure, volume and temperature on gases in everyday
life. We will first examine these effects more closely and then look at the theory that scientists use
to explain the behavior.
Pressure and Volume of Gases at Constant Temperature and Amount of Gas.
If the temperature and amount of a gas remains the same, and its volume is decreased, then
the pressure it exerts will increase. For example, when you use a bicycle pump, the pressure of the
air in the pump increases as its volume decreases. The reverse volume-pressure relationship is also
true: when the volume of a gas increases, its pressure decreases as long as the amount of gas and its
temperature remains constant.
Temperature and Pressure at Constant Volume and Amount of Gas.
If we put a fixed amount of gas into a container with rigid walls (so the volume doesn’t
change) and heat it, the pressure exerted by the gas will increase. For example, an aerosol can put
into a fire will explode due to the pressure buildup as the gas heats. If we cool a fixed amount of
gas, the pressure it exerts will decrease if its volume is kept the same.
The Amount of Gas and Its Pressure at Constant Temperature and Volume.
When you add more gas to a container with a fixed volume and temperature, then the
pressure will increase. This is what happens when you add air to a car tire. Removing the gas from
a container of fixed volume results in a pressure decrease if the temperature remains the same.
The Amount of Gas and Its Volume at Constant Temperature and Pressure.
Adding more gas to a container with flexible walls results in a volume increase if the
pressure and temperature at the end are the same as at the beginning. An example of this is when
you blow up a balloon.
Volume and Temperature at Constant Pressure and Amount of Gas.
If a gas is kept in a container with flexible walls, such as a balloon, then the pressure exerted
by the gas is just about equal to the pressure of the air outside the container. If this were not so then
one of the gases (the air or the gas in the balloon) would push on the container wall harder and the
wall would move.
If such a gas is heated, then its final volume will be larger than its initial volume if the
pressure is the same at the end as at the beginning. If a gas is cooled, then its volume decreases if
the pressure is the same at the end as at the beginning.
The Kinetic Molecular Theory of Gases
The kinetic molecular theory is a description of how gas molecules behave. This description
can then be used to explain the pressure-volume-temperature behavior of gases. Kinetic molecular
theory says the following:
(1) The submicroscopic particles of a gas move constantly. They move in a straight line
until they bounce off of something and then they change their direction of movement.
(2) The average speed at which the submicroscopic particles of a gas move increases when
the temperature of the gas is increased.
Let’s see how the simple behavior proposed by the kinetic molecular theory of gases
explains the interactions of the pressure, temperature, volume and the amount of gas. To do this, we
must first consider what causes a gas to exert a pressure.
175
When the submicroscopic particles of a gas collide with the walls of the container, they
exert a force on the container. To visualize this, imagine many tiny bb’s colliding with a wall. The
force exerted by the molecules is spread over the area of the container walls. The force exerted on
the walls divided by the area of the walls is the pressure that the gas is exerting on the container.
Pressure and Volume of Gases at Constant Temperature and Amount of Gas.
Let’s now imagine what happens if the volume of the container is decreased. The area of
the container walls will decrease, but the same number of molecules is colliding with the walls.
This means at any time and in any given region there must be more molecules colliding with the
wall after the volume is decreased than before. There will then be more force exerted on an area
and the pressure will be higher. This explains the behavior of increased pressure with decreased
volume. The reverse is also true, of course. If the volume is made larger then the same number of
molecules are colliding over a larger container wall area. Thus the force exerted in any one region
is lower after the volume is increased, and the pressure decreases.
larger volume
smaller volume
Temperature and Pressure at Constant Volume and Amount of Gas.
According to the kinetic molecular theory of gases, when the temperature of a gas is
increased, the average speed of its submicroscopic particles increases. This means the particles
would strike the container walls with more force, thus exerting more pressure since the area remains
the same. If the gas is cooled, then the molecules will slow down and strike the container walls
with less force, thus exerting less pressure.
The Amount of Gas and Its Pressure at Constant Temperature and Volume.
If the volume and temperature of a gas sample are kept the same and more of the gas is
added to a container, then there are more gaseous submicroscopic particles present. More of them
will then collide with the container walls at any given time, increasing the force on the walls and so
increasing the pressure. When the amount of gas is decreased, then there are fewer particles of gas
to collide with the container walls. The force on the same area is then decreased and the pressure
also decreases.
The Amount of Gas and Its Volume at Constant Temperature and Pressure.
If gas is added to a container with flexible walls, then the force exerted by the particles will
increase and push the walls outward. As the walls go out, the volume increases, which make the
pressure go back down to what it was before. At this point the pressure on the inside and outside of
the walls is the same and the walls quit moving. The end result is thus a larger volume at the same
pressure and temperature as at the start.
Volume and Temperature at Constant Pressure and Amount of Gas.
When a gas in a container with flexible walls is heated, the submicroscopic particles of the gas move
faster. They then exert more force on the walls of the container and push the walls outward. As the walls go
out, the volume increases, which makes the pressure go back down to what it was before. At this point the
pressure on the inside and outside of the walls is the same and the walls quit moving. The end result is
thus a larger volume at the same pressure and amount of gas as at the start.
176
Supplement 28: Determining Molecular Geometry from Lewis Structures
Molecular Geometry. One of the uses of Lewis structures is to be able to predict the shapes of
molecules. It is very simple to predict the shapes and then carry that further in predicting properties
of molecules. For now we will be concerned only with shapes or geometries. A simple set of rules
and drawings follows to help with predicting and illustrating molecular shapes:
1.
Draw a proper Lewis structure.
2.
Determine the number of groups surrounding any central atom. Any atom counts as one
group, no matter how many bonds are involved. So a single bonded atom counts as one group, but a
double bonded or triple bonded atom also counts as one group. Also every pair of non-bonded
electrons also counts as a group. So one non-bonded pair counts as one group, two non-bonded
pairs counts as two groups, etc.
3.
The total of bonded atoms plus the non-bonded pairs determines the number of groups.
These groups try to move as far apart as possible. The reasoning for them trying to move apart is
that the electrons all have the same energies and that they all are negatively charged [like charges
repel]. This limits the shapes or geometries that are available. Examples for simple cases are given
in the table below:
Number of
electron groups
about central atom
Electron
Geometry
around
Central Atom
2
linear
3
trigonal
planar
4
tetrahedral
Sketch
Molecular Geometry
The molecular geometry describes the way the atoms are arranged in space; it differs from an
electron geometry only if there are lone pairs on the central atom. There are a number molecular
geometries, but all of these can be obtained by first drawing the arrangement of the electron
groups as above (showing the correct 3-D angles), and then simply erasing the lone pairs. When
doing this, it will be easier to interpret the shape if you put lone pairs out of the plane of the
paper (with the wedges). For atoms with more than one central atom, it would be best to put
bonds between central atoms in the plane of the paper if possible, and leave any needed wedges
for perimeter atoms.
177
Let’s do an PH3 as an example. First we would draw the Lewis structure.
H P
H
H
We see that there are three bonds and one unbonded pair, so there are four groups of atoms around
the P. This means the geometry of the electron groups is shown by the figure on the left below.
To get the molecular geometry, we simply erase the lone pair and its pointing line and thus
obtain the figure on the right
Examples for simple cases of molecular geometry are given in the table below, where each O
represents an atom.
Geometries of Molecules
Sketch
Molecular geometry
Sketch
O
.
O
.
O
O
.
Molecular Geometry
.
linear
tetrahedral
O
O
O
O
O
.
trigonal
planar
O
O
O
.
.
O
O
O
.
O
.
O
.
O
.
O
trigonal
pyramidal
O
angular
or bent
angular
or bent
O
O
.
These geometries have certain bond angles associated with the groups. Think of the central
atom as being the locus of an angle defined by lines drawn towards any two of the groups on that
o
central atom. An example showing the 180 angle for a linear set of three atoms is shown below:
180o
O
O
O
Here is a model showing the angles for a trigonal planar group of atoms:
178
O
120o 120o
O
O
120o
O
o
The 109.5 angle of the tetrahedron is a little more difficult to show on paper because, unlike the
previous two structures which can have all of the groups on the plane of the page, the tetrahedron is
a three dimensional figure. A model might look something like this:
O
109.5o 109.5o
109.5o
O
O
O
109.5o
O
Four of the six bond angles have been marked to give you some feel for this structure.
179
Supplement 29: Concentration Units, Molarity and Dilutions
There are many occasions when it is very important to know how much solute is present in a given
amount of solvent. The amount of solute in a given amount of solvent is called the concentration.
There are lots of different units of concentration which are useful for different situations. Ones that
we will learn about at the appropriate times are wt/wt percent, wt/vol percent, vol/vol percent, ppm,
ppb, mole fraction, etc.
The unit that is most important for the stoichiometry of solutions is called molarity.
Molarity is defined as
Molarity =
mole of solute
Volume of solution in Litre
= M
.
In short, we write
M = mol
V
where it is understood that the V is the volume of the solution in liters. (Not volume of the solvent
added.)
We know all we need to know to be able to calculate the molarity of a solution.
Example.
What is the molarity of a solution obtained by diluting 10.0 g of NaOH to a final volume of 500
mL? FW for NaOH is 40.00 g/mol
Recall that
mol = wt x 1
10.0 g NaOH x 1 mol NaOH = 0.0250 mol NaOH
40.00g
So,
M = 0.0250 mol = 0.500 mol NaOH
0.500 L
L
Let's try a different version of the problem. How many moles and how many grams of HCl are
contained in 26.7 mL of 0.435 M HCl?
26.7 mL = 0.0267 L
Rearrange the definition of molarity to get
mol = V M.
Then,
0.0267 L  0.435 mol/L = 0.0116 mol HCl,
and the weight of HCl is
wt = mol  FW = 0.0116 mol HCl  36.46 g/mol HCl = 0.423 g HCl
180
Dilutions
Example : 0.100 L of pure water is added to 0.500 L of a 0.200 M solution
of NaCl. What is the new molarity after this dilution?
All we need to do here is maintain the same number of moles but change
the volume.
Number of moles of NaCl:
0.500 L x 0.200 mol NaCl = 0.100 mol NaCl
lL sol
New volume
New molarity
0.600 L
0.500 L+ 0.100 L = 0.600L
0.100 mol NaCl = 0.167 M NaCl
There is a short-cut for dilution problems. One can set-up the relation
M1V1 = M2V2
where the subscript 1 refers to the original solution and the subscript 2 refers to the final solution
after dilution.
For this example we have
0.200 M × 0.500 L = X M× 0.600L
solving for the unknown final concentration we get
X M = 0.200 M × 0.500 L = 0.167 M NaCl
0.600 L
Molality
Molality is similar to but not the same as molarity. Molality, m, is defined by,
m = mol solute
kg solvent
Note that for water solutions 1.00 kg of water has a volume of 1.00 L. If the solution is dilute the
volume of the solution formed from 1 L of water is still approximately 1 L, so that the molarity and
molality are about the same. However, in concentrated water solutions and in solutions where the
solvent is not water the molarity and molality are very different.
181
Mole Fraction
Mole fraction is essentially self-defined. In equation form the mole fraction (usually symbolized by
X) is
XA = mol A
total mol
Other units which are important in medicine and medical technology, food, and in monitoring the
environment are:
Weight percent (wt%)
Weight percent is defined as,
wt% = mass of solute
x 100
mass of solution
Volume percent or percent by volume (vol%)
Percent by volume is usually used for liquid-liquid solutions, as in alcoholic beverages. It is defined
by,
vol% = volume of solute
x 100
volume of solution
(The "proof" of an alcoholic beverage is twice the vol% concentration of ethyl alcohol. A 100 proof
beverage is 50 vol% alcohol.)
wt/vol%
Perhaps the most common of the "%" type concentration units is the wt/vol%. This unit is defined
by,
wt % =
mass of solute(g)
x 100
vol
volume of solution(ml)
Note that the units in this one matter. The units of wt/vol% are g/100mL.
Parts per million (ppm)
Parts per million is usually used to describe the concentration of trace contaminants in otherwise
pure materials. ppm is defined by
ppm =
mass of solute x 106
mass of solution
Since for very dilute solutions the mass of the solution and the mass of the solvent are very nearly
the same one could just as well write,
ppm = mass of solute x 106
mass of solvent
182
Parts per billion (ppb)
Nowadays, with increasingly accurate methods of analysis and increasing concern over minute
amounts of some suspected pollutants, we frequently see reports with concentrations in the ppb
range. ppb is defined similarly to ppm as,
ppb =
mass of solute x 109
mass of solvent
Molarity is useful in chemical analysis because volumes of solutions are much easier to control or
measure out than weights.
Now we can do stoichiometry with solutions.
So if we consider the possibilities for the reaction pathway we would get a diagram like,
We can start with either a weight or a volume and end up with either a weight or a volume (or even
a new molarity). Stoichiometry is used a lot in quantitative analyses.
Example.
A sample is said to contain some sodium carbonate. If it takes 32.6 mL of 0.435 M HCl to react with all the
sodium carbonate in the sample, how many grams of Na2CO3 are in the sample?
We have seen this reaction before,
Na2CO3 + 2 HCl  2 NaCl + CO2 + H2O.
(FW of Na2CO3 is 106.0 g/mol)
We must use the path: vol  mol  new mol  new wt
Find mol,
convert to new mol,
convert to new wt (weight of Na2CO3).
183
One can make known molarity solutions of redox reagents.
Let's say we have a 0.228 M solution of KMnO4. We want to use this to analyze for the concentration of iron
in another solution.
Take a 20.0 mL sample of the solution with the unknown iron concentration.
(Make sure that all of the iron is in the Fe2+ oxidation state by mixing the sample with Zn metal and then
filtering out the Zn metal remaining. The reaction for this reduction is
2 Fe3+(aq) + Zn(s)  2 Fe2+(aq) + Zn2+(aq).)
Now our Fe2+ can be oxidized by the permanganate ion back to Fe3+ according to the reaction
Fe2+ + MnO4 +8 H+  Fe3+ + Mn2+ + 4H2O
Do the experiment and find that it takes 36.82 mL of the permanganate solution to react with the iron (II)
ions.
Our road map is vol  mol  new mol  Oops again!
If we knew the molarity of the new solution we could get the new volume. But this time we know the volume
of the new solution and we know the number of moles so we can get a new molarity.
.
So, convert the volume of the 0.228 M permanganate solution into moles of permanganate ion.
Use the reaction fraction to convert mol of permanganate to mol of Fe2+.
Use the mole of Fe2+ and the volume of the iron solution sample to find the molarity of the iron in the
sample.
184
Supplement 30: The Mole
Once we know the mass of a sample, we can use the mass of the atoms or molecules in the sample
to determine how many atoms or molecules are in the sample.
For example, how many 12C atoms are in a 1.00 kg block of pure 12C isotope?
1.00 kg 12C
x 1000g x 1 amu
1 kg
1.660551 x 10-24 g
12
C atom
= 5.02 x 1025 12C atom
12.000000 amu
x
Obviously atoms and molecules are not a convenient unit of measure when we're working with
macroscopic (i.e., human size) objects. For this reason chemist define a new unit of measure called
the mole.
1 mole is defined as the number of carbon atoms in exactly twelve grams of pure 12C.
From the definition of atomic mass unit we get
1 mole = 6.0220943 x 1023 chemical units
The number 6.0220943 x 1023 is called Avogadro's number.
How many moles are in a 1.00 kg block of pure 12C isotope?
1.00 kg 12C x 1000g x 1 mole
1 kg
12.000000 g
12
C
= 83.3 moles 12C atom
Of course we know a 1.00 kg block of naturally occurring carbon will contain a mixture of 12C, 13C,
and even some 14C isotopes. So the number of moles of carbon atoms in a 1.00 kg block of naturally
occurring carbon is
1.00 kg 12C x 1000g x 1 mole
1 kg
12.011 g
12
C
= 83.26 moles 12C atom
We can also calculate the mass of a mole of molecules. For example, what is the mass of a mole of
methane (CH4) molecules?
mass of 1 mole of C is12.011 g
mass of 4 moles of H is
4 x 1.00g
mass of 1 mole of CH4 is
16.043g
or
16.043g CH4 /mol
In other words the molar mass of CH4 is 16.043g/mole.
molar mass = the mass in grams of 1 mole of a molecule.
For ionic compounds, which do not exist as individual molecules we use the term
Formula mass = The mass in grams of 1 mole of the chemical formula
185
For example, what is the formula mass of CaCO3?
mass of 1 mole of Ca is
mass of 1 mole of C is
mass of 3 moles of O is
mass of 1 mole CaCO3 is
Formula mass of CaCO3 is
40.08 g
12.011 g
3 x 15.999g
100.09g
or
100.09g CaCO3 /mol
Another example: How many hydrogen atoms are there in 2.50 g of NH3?
2.50 g NH3 x 1 mole x 6.022 x 1023 molecules x 3 H atoms
17 g NH3
1 mole NH3
1 NH3 molecule
186
= 5.02 x 1025 H atom
Supplement 31 Mass Percent (Composition)
The Mass Percent of a component is defined as
Mass % = mass of component x 100%
total mass of sample
For example, what is the mass percents of carbon, hydrogen, and oxygen in pure ethanol C2H6O?
First we calculate the mass of one mole of C2H6O...
mass of 2 moles of C is
2 x 12.011 g
mass of 6 moles of H is
6 x 1.008 g
mass of 1 mole of O is
15.999 g
Molecular weight of C2H6O is
46.069 g C2H6O / mol
Next we calculate the mass percents
mass % C = 24.002 g x 100 % = 52.14 %
46.069 g
mass % H = 6.048 g x 100 % = 13.13 %
46.069 g
mass % O = 15.999 g x 100 % = 34.73 %
46.069 g
Note that the mass percentages should add up to 100%.
187
Supplement 32--Combustion Analysis
When chemists make new compounds one of the first things they often do is determine the mass %
for the different elements in the compound. To analyze the mass percent of carbon and hydrogen
chemist use a combustion device.
The sample is burned in the presence of excess oxygen which converts all the carbon to carbon
dioxide and all the hydrogen to water. The CO2 and H2O produced are absorbed in two different
stages and their masses determined by measuring the increase in weight of the absorbers.
For example, Ascorbic acid (vitamin C) contains only C, H, and O. Combustion of 1.000 g of
Ascorbic acid produced 40.9% C and 4.5% H. What is the empirical formula for Ascorbic Acid?
First we need to calculate the mass percent of Oxygen. Since the sample contains C, H, and O, then
the remaining
100% - 40.9% - 4.5% = 54.6% is Oxygen
Now we need to express the composition in grams and determine the number of moles of each
element:
1.000g sample
x 40.9% x
mole C
100%
12.011g
1.000g sample
x 4.5% x
mole H
100%
1.008g
1.000g sample
x 54.6% x
mole O
100%
16.011g
=
=
=
0.0340 moles C
0.045 moles H
0.0340 moles O
Next we divide by the smallest number of moles to obtain the mole ratio which is also the atom
ratio. In this case carbon has the smallest number of moles so...
C:
H:
O:
0.0340 moles/0.0340 moles =
0.045 moles/0.0340 moles =
0.0341 moles/0.0340 moles ~
1
1.32 ~ 1 1/3
1
Finally we calculate the smallest whole integer ratios by multiplying each number above by 3 to get
C:3 H:4 O:3
thus we obtain the empirical formula C3H4O3.
Remember the empirical formula has the smallest whole integer ratios, the molecular formula can
be different, e.g., C6H8O6, or C9H12O9, or C12H16O12, ... are all possible molecular formulas.
Now let's ask a related question. What is the molecular formula if the molecular weight of Ascorbic
Acid was formed to be 176 g/mole?
In this case we need to find the multiplication factor between the molecular formula and the
empirical formula:
factor = (molecular weight)/(empirical formula weight)
188
The empirical formula weight of Ascorbic Acid is
mass of 3 moles of C is
mass of 4 moles of H is
mass of 3 moles of O is
mass of 1 mole of CH4 is
3 X 12.011 g
4 X 1.008g
3 X 15.999g
88.062g
Therefore the multiplicative factor is (176 g/mole)/(88.062 g/mole) ~ 2, and the molecular formula
for Ascorbic Acid is C6H8O6
189
Supplement 33--Chemical Equations
The most fundamental building blocks of matter for chemists are atoms (that is, the elements you
see in the periodic table). Atoms can combine with other atoms by chemical bonding to form
molecules. Recall that the process where a molecule is transformed into a different molecule is
called a chemical change. This process of chemical change is represented by a chemical reaction.
For example,
CH4 (g) + 2 O2 (g)
reactants
CO2 (g) + 2 H2O (g)
products
It is important to keep in mind that while in a chemical reaction molecules are destroyed and created
by breaking and forming chemical bonds, atoms are neither created nor destroyed in a chemical
reaction. In other words - there must be the same number of each type of atom on the product and
reactant sides of the arrow. Making sure that this rule is obeyed is called "balancing the chemical
equation". In the above example the equation would be unbalanced without the 2 in front of the O2
and H2O. We can make a table to confirm that the number of atoms on each side of the arrow are
the same:
Reactants
1C
4H
4O
Products
1C
4H
4O
Balanced?
yes
yes
yes
Notice that we never change the chemical formula of any product or reactant when trying to balance
a chemical equation.
A balanced equation is essential to known the stoichiometry for the chemical reaction.
Stoichiometry - Relationship between quantities of matter that participate in chemical reactions.
-orHow much of this, plus how much of that gives how much of something else?
Let's consider the following example of water being converted into hydrogen and oxygen gas using
a Hoffman Apparatus:
2 H2O(l) --> 2 H2(g) + O2(g)
The balanced chemical equation tells us that two molecules of water in the liquid state react to form
two molecules of H2 in the gas state and one molecule of O2 in the gas state. The Hoffman
apparatus allows us to trap the gaseous H2(g) and gaseous O2(g) in separate volumes. By examining
these volumes and using Avagadro's law we can see the stoichiometry of the chemical equation.
Avogadro's Law: The volume of a gas is directly proportional to the number of molecules in the
gas.
In the demonstration you would see that there is twice the amount of H2 gas and O2 gas, as
predicted by the stoichiometry of the chemical equation. These integer relationships between gas
volumes were some of the earliest proof we had that matter existed in discrete packages called
atoms and molecules.
Just as any good chef wants to add together the correct amount of each ingredient for the best recipe
190
so does the chemist also want to add together the correct number of molecules called for by a
chemical reaction to produce the desired product. For example, to separate water molecules into H2
and O2 gas our Hoffmann apparatus had to put electrical energy into the reaction. We could have
written the chemical reaction for this as
2 H2O(l) + 571.6 kJ --> 2 H2(g) + O2(g)
That is, we use 2 moles of H2O and 571.6 kJ of energy to make 2 moles of H2 gas and 1 mole of O2
gas.
Now let's say we want to do the reverse, that is, convert H2 gas and O2 gas into water and energy.
2 H2(g) + O2(g) --> 2 H2O(l) + 571.6 kJ
The stoichiometry tells us that combining two moles of H2(g) with one mole of O2(g) will give 2
moles of H2O(l) plus 571.6 kJ of energy.
Let's look at another example. Baking soda (NaHCO3) is used as an antacid. It neutralizes excess
HCl secreted by the stomach.
NaHCO3 + HCl(aq) --> NaCl(aq) + H2O(l) + CO2(aq)
How many moles of HCl are neutralized per gram of NaHCO3?
We set this problem up similar to the previous one:
1.00 g NaHCO3 x 1 mole NaHCO3 x 1 mole HCl
84.01 g NaHCO3 1 mole NaHCO3
Convert g to moles
= 1.19 x 10-2 moles HCl
neutralized
mole ratio
Now try this one at home. Milk of Magnesia (Mg(OH)2) is also an antacid, and the chemical
reaction is
Mg(OH)2(s) + 2HCl(aq) --> 2 H2O(l) + MgCl2(aq)
Which neutralizes more HCl per gram, NaHCO3 or Mg(OH)2?
191
Supplement 34--Limiting Reagents (Limiting Reactants)
Most everyone understands the concept of a limiting reagent. It's the same idea as when you go to
the grocery store and buy a package of hot dog wieners and a package of hot dog buns. When you
get home you realize that you bought 10 wieners but only 8 buns (you're two buns short!). If you
want to put one wiener in one bun then your limiting ingredient will be the hot dog buns. Likewise
your excess ingredient will be the hot dog wieners.
The same thing happens in chemistry.
Limiting Reagent - Reagent that limits the amount of products that can be formed.
For example, nitrogen gas is prepared by passing ammonia gas over solid copper(II) oxide at high
temperatures. The other products are solid copper and water vapor.
2 NH3(g) + 3 CuO(s) --> N2(g) + 3 Cu(s) + 3 H2O(g)
If 18.1 g of NH3 are reacted with 90.4 g of CuO, which is the limiting reagent? How many grams of
N2 will be formed?
First we compute the number of moles of NH3 (M.W. = 17.0 g/mole) and the number of moles of
CuO (M.W. = 79.5 g/mole).
18.1 g NH3 x 1 mole NH3 = 1.06 mole NH3
17.0 g NH3
available
90.4 g CuO x 1 mole CuO = 1.14 mole CuO
79.5 g CuO
available
To determine which reagent is limiting we use the mole ratio from the chemical equation to convert
moles NH3 to moles CuO.
1.06 mole NH3 x 3 mole CuO = 1.59 mole CuO
2 mole NH3
needed
So, only 1.14 moles of CuO is available, therefore CuO is the limiting reagent. That is, CuO will
run out before the NH3 does.
The mass of N2 produced will be
1.14 mole CuO x 1 mole N2 x 28.0 g N2 = 10.6 g N2 produced
3 mole CuO 1mole N2
192
Supplement 35--Solution Chemistry
If one of the substances is present in much greater quantities than all the other substances then it is
called the solvent. The other substances in solution are known as solutes. For example, when a
small amount of NH4Cl is dissolved in a large quantity of water we refer to water as the solvent and
NH4Cl as the solute. Another example is Naphthalene (used in mothballs) can be dissolved in
benzene. In this example benzene is the solvent and naphthalene is the solute.
Solutes dissolved in water (solvent) are called aqueous solutions. Not all substances are soluble in
water. Why do some substances dissolve in water and others don't? It has to do with the structure of
the water molecule.
O
H
H
oxygen and hydrogen share electrons in the covalent bond, but they
don’t share electrons equally.
Oxygen has a greater attraction for electrons, so the shared electrons (bonding electrons) spend
more time close to oxygen then to either of the hydrogens. This gives oxygen a slightly excess
negative charge and hydrogen a slightly more positive charge.

O
H

H

This unequal charge distribution makes water a polar molecule, and gives water its ability to
dissolve compounds. When an ionic solid dissolves in water, the positive ends of the water
molecule are attracted to the negatively charged anions and the negative ends of the water molecule
are attracted to the positively charged cations.
So when an ionic substance (salt) dissolves in water, it is broken up into individual cations and
anions which are surrounded by water molecules. For example, when NH4 NO3 is dissolved in
water it breaks up into separate ions.
H2O
NH4NO3(s)
NH4+ (aq) + NO3 - (aq)
NH4+ and NO3- ions are floating around in H2O essentially independent of each other.
Water also dissolves non-ionic substances. For example, C2H5OH (ethanol) is very soluble in H2O.
This is because C2H5OH has a polar OH bond that the water molecules like to hang around.
193
H
H
H
H
H
H
O
H
H
O
O
H
H
Many substances do not dissolve in water and that is because they are non-polar and do not interact
well with water molecules. A common example is oil and water. Oil contains molecules that are
non-polar, thus they do not dissolve in water.
How do we know that ionic solids dissolve in water and form cations and anions that float around
separately? One clue comes from conductivity experiments. Anions and Cations should act as
charge carriers in solution. Therefore a solution with dissolved ions should conduct electricity. Let's
look at a few examples. Pure (distilled) water contains no dissolved ions. Therefore pure water will
not conduct electricity. In a simple conductivity experiment as shown below we would not expect
the light to be on.
voltage
voltage
Na+
Cl-
An aqueous NaCl solution, however, will have dissolved ions present and therefore will conduct
electricity. Therefore the light in our conductivity experiment will be on if dipped in an aqueous
NaCl solution.
Substances that exist in solution almost completely as ions are called strong electrolytes.
Substances that do not form ions when they dissolve in water are called non-electrolytes. And
example of a non-electrolyte is sugar. Sugar will readily dissolve in water but doesn't form cations
and anions in solution. That is, there are no charge carriers formed.
H2O
C12H22O11
C12H22O11
Dissolves, does not fall apart
Substances that only partially ionize into ions when dissolved in water are called weak electrolytes.
For example, Acetic Acid (HC2H3O2) dissolves in water, but only partially dissociates into ions.
H2O
HC2H3O2 (s)
HC2H3O2
(aq)
Dissolves
H2O
194
HC2H3O2
(aq)
H+ (aq) + C2H3O2 -
(aq)
Dissolves, and slightly ionizes
The double arrows tell us that the equation is in equilibrium with the ionized and un-ionized
compound.
Be careful not to confuse how soluble a substance is in water with whether it is a weak, strong, or
non-electrolyte. For example, sugar dissolves completely in water but it is a non-electrolyte.
195
Supplement 36: Solution Stoichiometry, Titrations
For balanced chemical equations involving solutions we calculate the number of moles by knowing
the concentration (moles/liter, or Molarity) and volume (in liters).
Example:
How many moles of water form when 25.0 ml of 0.100 M HNO3 (nitric acid) solution is completely
neutralized by NaOH (a base)?
1. Begin by writing the balanced equation for the reaction:
HNO3(aq) + NaOH(aq)  NaNO3(aq) + H2O(l)
2. The relationship between HNO3 and H2O is 1:1, for one mole of HNO3 that is completely
consumed (i.e. neutralized) in the reaction, one mole of H2O is produced.
3. How many moles of HNO3 are we starting with?
0.250 L x 0.100 mol = 0.00250 mol HNO3
L
-3
2.50 x 10 mol HNO3 x 1 mol H2O = 2.50 x 10-3 mol H2O
1 mol HNO3
Therefore, we should have produced 2.50 x 10-3 moles of H2O
Titrations
How can we know the concentration of some solution of interest? One answer to this problem lies
in the method of titration. In titration we will make use of a second solution known as a standard
solution which has the following characteristics:
The second solution contains a chemical which reacts in a defined way, with known stoichiometry,
with the solute of the first solution
The concentration of the solute in this second solution is known.
Classic titrations include so-called acid-base titrations. In these experiments a solution of an acid
with an unknown concentration is titrated with a solution of known concentration of base (or vice
versa). For example, we may have a solution of hydrochloric acid (HCl) of unknown concentration
and a standard solution of NaOH. To a fixed amount of the HCl solution is added incremental
amounts of the NaOH solution until the acid is completely neutralized - i.e. a stoichiometrically
equivalent quantity of HCl and NaOH have been combined. This is known as the equivalence point
in the titration. By knowing the concentration of the standard solution, and the amount added to
achieve stoichiometric equivalency, we can determine the amount of moles of HCl in the original
sample volume.
196
1.0 M NaOH(base)
Unknown conc. of HCl(acid)
How do we know when we have reached the equivalence point in such a titration experiment? In
this type of acid-base titration, so called indicator-dyes are used. For example phenolphthalein is
colorless in acidic solutions and turns pink in basic solutions. Thus, in the above experiment we
will add a small amount of this indicator-dye and add base until we barely begin to see a color
change to pink.
25.00 ml of a solution of HCl with an unknown concentration is titrated with a standard solution of
0.5000 M NaOH. The phenolphthalein indicator dye begins to turn color after the addition of 12.88
ml of standard solution. What is the concentration of the HCl?
Balanced equation for the reaction:
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
First, calculate the number of moles of NaOH added.
0.01288 L NaOH x 0.5000 mol NaOH = 6.440 x 10-3 mol of NaOH added
L
From the stoichiometry of the balanced chemical equation we can now calculate the number of
moles of HCl in the sample.
6.440 x 10-3 mol of NaOH x 1 mol HCl
1 mol NaOH
=
6.440 x 10-3 mol of HCl in sample
Remember that the volume of HCl solution used was 25.00 mL or 0.02500 L.
Therefore, the Molarity of the HCl solution is:
6.440 x 10-3 mol HCl = 0.2576 mol HCl
0.02500 L
L
or
0.2576 M HCl
197
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