Not the
real Mr. Cooper
“Good morning, and welcome to
introduction to chemistry.”
Info
• Class: Chemistry
• Instructor: Mr. Cooper
• Office: A112 - I’m pretty much in my
classroom before and after school
• E-mail: gcooper@lps.org
• WebPage: lsw.lps.org click on teachers
and find my name.
Class Format
• 1. Daily Quizzes (D.Q.):
– Each day will start with a quiz over the previous
days material.
– Clear your desk, have out a piece of paper and be
ready at the beginning of class.
– Recommendation: Write out the question and
answer and keep a running list of the quizzes on
the same sheet of paper.
– At the end of the term you will have created a
review sheet. I will not be giving you a review
sheet at the end of the term.
– Quizzes will not be picked up but your score will
be recorded on your self-evaluation sheet.
Class Format
• 2. Lecture:
• I will provide you one set of Lecture notes.
They are posted on the web. You can print
them off if you lose your set.
• This allows you to listen and formulate
questions and take your own notes rather
than just copying down my outline.
• If all of the class has notes ahead of time, we
can spend more time on answering
questions, lab work, book work and individual
help rather than copying notes.
Class Format
• 3. Laboratory:.
• One typed lab report of your choice will be
submitted per unit. You choose the lab you want to
do the report on. It is suggested that you choose
one at the beginning of the unit so that you can get
it done ahead of time. It will be due day of test.
• Leave 5 minutes at the end of each period for
clean up. Labs are not expected to be homework.
If you work diligently in class, you should get them
done. The labs are on-line if you lose the one
provided lab book.
• Labs may be replaced by worksheets, group work,
a video, or demonstrations
Class Format
• 4. If there is time left at the end of class
you are expected to be working on book
problems, which are assigned on a daily
basis. See unit outline.
• This is also an excellent time for you to
get individual help.
Grading
• Lab Quizzes: The day before each unit
test will be a unit lab quiz. It is open lab
book.
Grading
• Quizzes and tests will be announced in
advance.
• Tests and quizzes will be closed note and
closed book. Exception: lab quizzes are open
lab book.
• Your textbook is your first resource; so read
it!!! All materials for this course are based off
of your text book.
• Also, calculators will be allowed on tests and
quizzes. It is your responsibility to provide a
calculator.
Grading
• Tests and quizzes are m.c., short
answer, problem solving, make you
think exams; not memorizing
exams(although you will need to have
some things memorized.)
Grading
• There will be no test retakes!!!!
• However, you will have a take home
practice test. So, your unit test is
actually your retake.
Grading
•
•
•
•
•
The grading scale is as follows:
A= 90-100
B+= 85-89
B= 80-84
C+=75-79
C= 70-74
D+= 65-69
D= 60-64
F= Below 60.0
Misc.
• Any assignments or test missed for
truancy results in 60.0% of earned
grade. This is district policy.
• Late work policy – Once I have your
assignment graded and handed back, I
will no longer accept that assignment.
Misc
• Tardies - Building policy is followed.
• Cell Phones – Building policy is followed.
• iPods - iPods are not to be used during
instructional time or lab time. You may use
them during individual work time at your desk.
I reserve the right to revoke privileges.
Student Expectations
• Do your job as a student which means:
• 1. Bring all needed materials to class.
• ex) books, notebooks, writing utensils, brain, good
attitude, etc.
• 2. Respect each students right to learn and their
property.
• 3. Listen carefully and follow instructions given.
• 4. READ, STUDY, PAY ATTENTION TO DETAIL
• 5. No food or drink.
• 6. Use class time to work.
• 7. ASK QUESTIONS in class or see me after school
for help.
All work will be shown at all
times or no credit will be
given.
I should never have to make
this announcement again or
even write in on a worksheet
or test.
•Do not give me
late work
without a roll of
the dice.
• Summary of Grades
– Lab Quiz will be day before test
– Due Day of test
• Self-eval sheet
• One random homework pickup
• Lab Report
– Multiple Choice Test
– Short Answer Test
Misc.
• “I do not feel obliged to believe that the same God
who has endowed us with sense, reason, and
intellect has intended us to forgo their use.” - Galileo
Galilei (astronomer and physicist)
• Air Force Core Values
– Integrity First
– Excellence in All
– Service Before Self
• Remember, I am working hard for you. I expect that
you will work hard for me.
• I find it offense when at the end of the term you are
begging me to round or expecting me to do you some
extra credit favor when you didn’t give me your best
to begin with.
Mr. Cooper
Equipment Use Review
• What lab equipment is used for handling a hot
beaker?
• What lab equipment would be used to hold a piece of
metal in a flame?
• What piece of lab equipment is used to measure
volume?
• A BEAKER OR FLASK IS NEVER USED AS A
MEASUREMENT DEVICE!!!
Tirrill (Bunsen)
Burner
How the parts work.
Turning the barrel
Controls type of flame
(orange or blue) by
opening and closing
the air vent.
Always use blue flame
(open vent);
however, vent
should not be wide
open for initial
igniting.
How the parts work.
Gas Flow Control
• Controls the height
of the flame through
controlling the
amount of gas
flowing.
• Use appropriate
flame height. NO
TORCHES.
Operation of the Tirrill (Bunsen)
Burner
• Hook the hose to the gas
inlet and gas jet
• Place spark ignitor next to
top of barrel
• Turn on gas and ignite with
sparker
• Make barrel and gas flow
control adjustments for
proper flame
Trouble Shooting
• You should be able
to hear the gas
flowing. If not:
• Check if gas flow
control valve is open
• Check if jet valve is
clogged. If so see
your teacher.
Troubleshooting
• Gas attempts to light
but goes out.
Possible cause is:
• Air vent is too far
open. Turn the
barrel down.
Formatting a Lab Report
• Title: The word “title” is written and underlined; followed then by the
name of the lab.
• Purpose: The word “purpose” is written and underlined; followed by the
purpose of the lab.
• Procedure: Usually extremely detailed. You can summarize. Just a
couple of sentences is fine. Procedure questions will be on quizzes.
• Data: The word “data” is written and underlined. For this section you
will either be filling out charts or questions will be asked to help you
gather data. Write out the question, underline it, leave a space, then
answer the question.
Formatting a Lab Report
• Conclusion: The word “conclusion” is written and underlined. For
this section you will be asked questions. Write out the question,
underline it, leave a space, then answer the question.
• Application: Where is this concept used in the real world or in
the scientific community? How does this affect your life or why is
this important to have this knowledge for society or other real
world application or future predicted use?
– Minimum 3 sentences and Maximum 5 sentences
– Must have one source to accompany this section. If you use
a website please make sure you do not have a typo in the
address.
– No opinions. I am not your source nor are you a source. This
is a research component. Do some research and quote your
source other than your text. Do not use any “I” statements.
Sample lab report
Title: Place title here.
Purpose: Place the purpose here.
Procedure: A couple of general sentences summarizing lab steps.
Data:
1.
I am writing out the question and underlining it.
A space was left and question 1 is answered.
2. Another question is written out and underlined.
A space was left and question 2 was answered.
Sample lab report
Conclusion:
1.
I am writing out the question and underlining it.
A space was left and question 1 is answered.
2. Another question is written out and underlined.
Do you see a pattern here?
Application: Use or Application
Source: WWW.SCIENCERULES.COM
Keep answers clear and concise. Length is not
important. I care about content and good
communication.
Bad Student Example - Very Bad
• Application:
• After doing this lab, I sat and wondered how I would apply what I learned to
something that expands outside of our classroom. Thinking about the candle burning
sent me into deep contemplation. And then, out of complete randomness, I started
thinking about our environment and the things that we burn which pollute it. I then
thought of where all the statistics we hear about come from, and how the claims are
substantiated. How do scientists know exactly what percent our ozone layer has
deteriorated, and what percent of our atmosphere is made up harmful pollutants?
Well when fossil fuels are burned, or maybe even things like wood or who knows,
scientists most likely calculate the molecules given off so they can come up with
these statistics. Well maybe they deal with moles or liters of gas at STP, who knows,
but I’m sure somewhere in there scientists will have to convert from moles to
molecules, or grams to moles, or grams to molecules, and in a sense that is what we
have done in this lab. We found out how many moles of wax were burned over 3
minutes, and if we know what wax is made of then we can figure out what exactly
was released into the atmosphere.
• Source: My brain .. no seriously .. my brain.
Acronyms
•
•
•
•
•
•
•
•
KISS
Keep It Simple Stupid
SOP
Standard Operating Procedure
HUA
Heard Understood Acknowledged
WAG
Wild Ass Guess
Metric Conversions
grams (g) is used for mass (weight)
liters (l) is used for volume
meters (m) is used for distance
kilo
k
hecto
h
deka
dk
SI
g
l
m
deci centi milli
d
c
m
Metric Conversions
kilo
hecto
deka
SI
deci centi milli
• We will use a problem solving process
called dimensional analysis (tracks).
• Example 25.0 cg = _______ g
• 25.0 cg 1 g
100 cg
= 0.250 g
Example 0.351 hl = _______ ml
0.351 hl 100,000 ml
1 hl
= 35,100 ml
Metric Conversions
kilo
hecto
deka
SI
deci centi milli
• Your turns - convert the following:
• 15.72 g = ______ mg 15.72 g = _____ kg
Density = m  v
D = m/v
• intensive property
– density is the same no matter size
• 50 grams of gold has the same density
as 150 grams of gold.
• Important density to remember
– water is 1.0 g/ml at 4 oC
– 1 cm3 = 1 ml
Density Problem
• A substance has a volume of 1.74 ml
and a mass of 20.0 grams. What is the
density?
Density Problem
• One more to test your algebra:
• What is the volume of ice in a container
if the density is 0.920 g/ml and the mass
is 58.39 g?
Properties of Matter
Definitions
• Matter - anything that has mass and
takes up space
• Mass - amount of matter an object
contains
• Substance (pure) - matter that has a
uniform composition
– Ex. Sugar - C12H22O11
– Lemonade is not a pure substance
Properties of Matter
States of Matter (Solid)
• Definite shape
• Definite volume
• Is incompressible (atom or molecules
can not be pushed closer together)
• Examples – coal, sugar, ice, etc
Properties of Matter
States of Matter (Liquid)
•
•
•
•
•
Matter that flows
Has a fixed volume
Takes the shape of its container
Incompressible
Examples
– Water, milk, blood, etc
Properties of Matter
States of Matter (Gas)
• Matter that takes the shape and volume
of its container
• Easily compressed
• Examples
– Oxygen, nitrogen, helium, etc
States of Matter Video
Properties of Matter
Physical Property
• An observed condition of the substance
• Physical properties help identify substances
• Examples include:
Properties of Matter
Physical Change
• A change which alters a given material
without changing its composition
• Nothing new is made
• Example
– Ice melting - new state of matter but
substance is still H2O
– Vapor - a substance that is in a gaseous
state but liquid at room temp
Change of State - a physical change
Properties of Matter
Chemical Property
• The ability or inability of a substance to
rearrange its atoms.
• Example
– Gasoline has the ability to react violently
with oxygen
Physical and Chemical Properties
Properties of Matter
Chemical Change
• The actual rearrangement of atoms
• Example
– The combustion of gasoline to make
carbon monoxide, carbon dioxide, carbon,
water (this produces a great amount of
energy)
Classifying a physical or
chemical change
• Ask yourself these questions:
• 1. Has something new been made?
– If yes than a chemical change occurred
– Indicators - color change, formation of precipitate,
absorption or release of energy, formation of a gas
• 2. What does it take to get back to the original
form?
– If a physical process can revert it back than the
change was physical.
A chemical change
Classify the following as a
physical or chemical change
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•
•
•
•
•
•
•
A sidewalk cracking
Blood clotting
Getting a tan
Making Kool-Aid
Making a hard boiled egg
Plastic melting in the sun
Autumn leaf colors
Digestion of food
• The ripening of a
banana
• Making ice cubes
• Milk curdling
• Turning on the
television
• Making toast
• Mowing the grass
• Paint fading
• Grey hair
Categorizing our environment
Classifying Matter
Pure Substances
Element
Compound
Mixtures
Heterogeneous
Homogeneous
aka - solution
Classifying Matter
Mixtures
• A physical blend of two or more
substances.
• Examples:
– Beef stew, air - mixture of gases
Classifying Matter
Mixture (Heterogeneous)
• Not uniform in composition
• One portion of the mixture is different
from the composition of another portion
• Example:
– Soil - sand, silt, clay, decayed material
Classifying matter
Mixture ( homogeneous)
• Completely uniform composition
• Components are evenly distributed
throughout the sample
• Example
– Alloys - mixture of metals (brass, steel)
• AKA - solution
– Example - ammonia, alloys, kool-aid
Classifying Matter
Mixtures can be separated by physical means
• Examples
– A spoon can separate beef stew
– Sulfur and iron can be separated with a magnet
• Tap water is a mixture that can be separated by
distillation.
• Distillation - a separation techniques based on the
physical property of boiling points.
– Liquid is boiled to produce a vapor
– Then condensed to a liquid leaving impurities behind
Categorizing our environment
Classifying Matter
Pure Substances
Element
Compound
Mixtures
Heterogeneous
Examples
river water
milk
beef stew
Homogeneous
aka - solution
Examples
pop
steel
kool-aid
Classifying Matter
Elements
• Simplest form of matter
• Cannot be broken down into anything
else
• Building blocks for all other substances
• Examples
– Hydrogen, oxygen, carbon
Classifying Matter
Compounds
• Two or more elements combined
through a chemical bond
• Can only be separated into simpler
substances by chemical reactions
• Example
– Sugar - C12H22O11
Chemical and physical properties of compounds
are different from their constituent elements.
• Examples
• Sugar
– Carbon is black
– Hydrogen is a gas
– Oxygen is a gas
• Salt - NaCl
– Na (sodium) soft metal that explodes with water
– Cl (Chlorine) pale yellow-green poisonous gas
Classifying matter review
Remember
• Substance - all of one kind of matter
– Examples: element or compound
• Mixture - has more than one kind of
material
– Examples - two or more compounds or
elements that are mixed, not chemically
combined
Categorizing our environment
Classifying Matter
Pure Substances
Mixtures
Element
Compound
Heterogeneous
Examples
Carbon (C)
Gold (Au)
Neon (Ne)
Examples
Sodium Chloride (NaCl)
Sugar (C12H22O11)
Dihydrogen Monoxide (H20)
Examples
river water
milk
beef stew
Homogeneous
aka - solution
Examples
pop
steel
kool-aid
The anatomy of the periodic table
• Get out your periodic tables
• Know where the following are on your
periodic table (p.t)
• Group A (representative elements)
• Group B
• Metals
• Nonmetals
• Metalloids (Semimetals)
– Note - aluminum is not considered a metalloid
The anatomy of the periodic table
• Know where the following are on your periodic table
(p.t) continued
• Transition metals
• Inner transition metals
• Alkali metals
• Alkaline metals
• Halogens
• Noble gases
Naming Compounds
Ionic
Molecular
i.e. covalent
Naming Molecular
Compounds
• Molecules are made up of nonmetals
• Prefixes are used to represent numbers of
atoms. See your text for prefixes
• Binary compounds end in -ide
• Examples
• Name? - Cl2O8 and OF2
• Formula for? - dinitrogen tetroxide
• Answers -
Naming Molecular
Compounds
• Your turn. Try these.
• Name or write the formula for:
– Boron trichloride
– Dinitrogen tetrahydride
– N2O5
PF5
S4N2
• Answers
CCl4
SO3
H2O
Take ten minutes and work a
few problems on the “Naming
covalent compounds” side of
your worksheet.
Ions
• An atom that carries a charge
• The charge on the ion is called the
Oxidation state or Oxidation Number
• Cation - positively charged atom
– Metals form cations
– CATions are PAWsitive
• Anion - negatively charged atom
– Nonmetals form anions
Naming Cations
Name the metal followed by the word
ion
• Example
– Na - sodium - neutral element
– Na1+- sodium ion - cation of the element
• Another example:
– Mg - magnesium
Mg2+ - Magnesium ion
Naming Anions
• Ending changes are used for Anions
• Elemental anions will end in -ide
• Example
– Cl2 - chlorine - neutral element
– Cl1- - chloride - anion of the element
• Another example
– O2 - oxygen
O2- Oxide
Writing Formulas for Binary Ionic
Compounds
• The periodic table tells you the charge for
group A (aka - the representative elements)
• Group 1A - 1+ Group 2A - 2+
• Group 3A - 3+ Group 4 - depends
• Group 5A - 3- Group 6A - 2• Group 7A - 1- Group 8A or (0) - does
not form ions
Naming
• Your turn:
– Name or write the symbol for the following:
•
•
•
•
Aluminum
Calcium Ion
Ga3+
K
Phosphide
Iodine
Nitrogen
Sulfide
Naming Binary Ionic Compounds
Name the metal then the nonmetal with
the ending changing to -ide
– The -ide tells the person it is a binary
compound and the anion portion.
• Examples: MgCl2
• Magnesium Chloride
• Potassium Sulfide
K2S
Writing Formulas for Binary Ionic
Compounds
• All compounds are electrically neutral
• To write the formula, figure out how many
cations and anions are needed so that the
number of positives and negatives are equal.
Find the least common multiple to figure
out the total number of +’s and -’s. Then
divide by the charge to find out how many
of each atom is needed!
• If X1+ and Y2-, what would be the formula?
• X2Y - Charges total 2 +’s and 2 -’s
Writing Formulas for Binary Ionic
Compounds
• If X3+ and Y2-, what would be the
formula?
• X2Y3 - Charges total 6 +’s and 6 -’s
• Find the formula for the following pairs
of ions:
– Na1+ , P3-
• Answers:
Sr2+ , N3-
• Now:
– Finish side 1 of worksheet
– Work sections 1 - 4 on back of worksheet
– Work homework problems
Writing formulas for multivalent ionic
compounds
• Transition metals have the ability to
form more than one cation
• Therefore, a roman numeral is placed in
the name to signify the charge on the
cation
• Example:
– Iron (III) Chloride
• Write the formula?
Writing formulas for mulitvalent ionic
compounds
• Write formulas for the following:
• Copper (I) Oxide
• Copper (II) Oxide
Answers -
Naming compounds with multivalent
metals
• If the metal is in group B it requires a
roman numeral in the name.
• You will have to deduce the roman
numeral based on the formula.
• Example
– Name CoI2
• Answer -
Naming compounds with multivalent
metals
• Deducing the roman numeral
• Multiply the charge on the anion by the number of
anions and then divide by the number of cations to
get the roman numeral.
• Write the names for Fe2S3 SnO2
• Answers -
• Take ten minutes and work on sections
5 and 6 on the back side of your
worksheet.
Polyatomic Ions
• A group of atoms that carry a charge
• Examples:
– SO42-
NO31-
• Names of polyatomic ions that contain oxygen will end
in -ate or -ite
• -ite is one less oxygen then ate
• Example
– Sulfate is SO42- Sulfite is SO32– Chlorate is ClO31- Chlorite is ClO21-
• Other polyatomic ions
– NH41+ Ammonium
– OH1- Hydroxide
CN1- cyanide
Writing formulas using polyatomic ions
• The polyatomic ion is treated as one unit.
• Balance the charges
• Place parenthesis around the polyatomic ion
if there is more than one
• Example
– Write the formula for Iron (II) Nitrate
Naming using Polyatomic ions
• Name the metal then name the
polyatomic ion. If you need a roman
numeral; include it.
• Treat the polyatomic ion as one unit (as
if it were one atom)
• Example - Name CuSO4
Exceptions for roman numerals
• Silver, Cadmium and Zinc do not get
roman numerals.
• Ag is always +1, Cadmium and Zinc are
always +2
• Tin and Lead need roman numerals.
They are multivalent (multiple oxidation
states)
Naming Acids
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•
•
•
•
•
Memorize
HCl - Hydrochloric Acid
H2SO4 - Sulfuric Acid
HNO3- Nitric Acid
H3PO4 - Phosphoric Acid
Note - Acids give off H1+ (Hydrogen ions) and
bases give off OH1- ions
• What do you get when an acid and base
combine?
Naming Compounds
Is there a metal?
Yes
No
Ionic
Molecular
Does the compound contain a
multivalent ion?
aka - transition metal or
group B element
Use prefixes to represent
the number of atoms.
Example: H2O Dihydrogen Monoxide
CO2 Cabon Dioxide
No
Yes
Name the cation first
then name the anion
Example: Lithium Fluoride
Magnesium Carbonate
Name the cation first
Place a roman numeral
Name the anion
Example: Iron (II) Sulfate
Check for understanding
• Name or write the formula for:
–
–
–
–
Potassium Sulfate
Chromium (III) Cyanide
Fe(ClO3)3
CuCl
• Answers
• Now finish your worksheet and work on your
homework.
• Get help
• Make sure and check your answers. You will be
writing formulas all year and doing math based on
these formulas. You get the formula wrong you get
the math wrong.
Helpful hints for balancing
chemical equations
•
•
•
•
•
Balance hydrogens second to last
Balance oxygens last
Check for lowest ratio
Coefficients must be whole numbers
Don’t break up your compounds with coefficients
– NaCl cannot become Na6Cl
• Do not change your subscripts
• Balance the polyatomic ions as one unit (if it didn’t
break apart)
• Perform a final check
Balance the following
•
C2H6 +
•
Na3PO4 +
O2 -->
CO2 +
H2O
Mg(NO3)2 --> NaNO3 + Mg3(PO4)2
Types of Reactions
Including reaction prediction
Generals about writing Equations
• Reactants on the left and products on the right
• Symbols - see text for symbols that are
included in equations.
– Ex: g for gas, l for liquid, s for solid
– aq for aqueous
• Catalyst goes above the arrow
•
KI
– Ex H2O2(aq) ---> H2O(l) + O2(g)
• Diatomic Molecules - BrINClHOF
– Elemental state - Br2I2N2Cl2H2O2F2
1. Synthesis (Combination)
• Two or more substances react to form a
single substance
• R + S --> RS
• Ex) SO3(g) + H2O(l) --> H2SO4(aq)
• Usually gives off energy when forming bonds
• Example: Write the balanced equation for:
magnesium ribbon reacting with oxygen
• Mg(s) + O2(g) ---> MgO(s)
• 2 Mg(s) + O2(g) --> 2MgO(s)
1. Synthesis (Combination)
• Your turn. Write balanced equations for
the following:
– Aluminum (s) reacts with oxygen (g)
– Hydrogen (g) reacts with oxygen (g)
• Answers:
2. Decomposition
• A single compound is broken down into simpler
products
• RS --> R + S
• Ex) BCl3 --> B + Cl2
• Requires energy to break chemical bonds (heat, light,
electricity)
• Example - Write the balanced equation for mercury
(II) oxide decomposing;
• HgO --> Hg + O2
• 2HgO --> 2Hg + O2
2. Decomposition
• Your turn. Write balanced equations for
the following:
• The decomposition of water
• The decomposition of lead (IV) oxide
• Answers
3. Single Replacement Reactions
• An element replaces an element of a
compound
• T + RS --> TS + R
• Ex) Zn(s) + H2SO4(aq) --> ZnSO4(aq) + H2(g)
• A metal may replace a metal or a nonmetal
may replace a nonmetal
• Activity Series - list of metal in order of
decreasing activity
• Nonmetals reactivity decreases as you go
down the periodic table
• This is limited to the halogens -group 7A
3. Single replacement reactions
• Ex) Write the balanced equation when
aluminum reacts with sulfuric acid
• Al(s) + H2SO4(aq) --> Al2(SO4)3(s) + H2(g)
• 2Al(s)+ 3H2SO4(aq) --> Al2(SO4)3(s) + 3H2(g)
3. Single replacement reactions
• Your turn. Write balanced equations for the following:
• When chlorine reacts with potassium iodide
• When copper (assume Cu2+) is added to Iron (II)
Sulfate
• Answers
–
4. Double Replacement
• Exchange of positive ions between two
compounds. Just swap the positive ions and
write the new formula.
• R+S- + T+U- --> R+U- + T+S• Ex) FeS(s) + 2HCl(aq) --> H2S(g) + FeCl2(aq)
• Ex) Write the balanced equation for barium
chloride added to potassium carbonate
• BaCl2(aq) + K2CO3(aq) --> BaCO3(s) + KCl(aq)
• BaCl2(aq) + K2CO3(aq) --> BaCO3(s) + 2 KCl(aq)
4. Double Replacement
• Your turn. Write balanced equations for
the following.
• Iron (III) Sulfide reacting with
hydrochloric acid
• Answer
5. Combustion Reactions
• Oxygen reacts with another substance, often
producing heat and light
• Often involve hydrocarbons
– Compounds of hydrogen and carbon
• Combustion of hydrocarbons produces a lot of
energy, therefore, hydrocarbons are used as fuels.
• Examples: methane, propane, butane, octane
5. Combustion Reactions
•
•
•
•
•
Two types of combustion
1. Complete combustion
CxHy + O2(g) --> CO2(g) + H2O(g) + energy
2. Incomplete combustion
Two more products: CO and C
• CxHy + O2(g) --> CO2(g) + H2O(g) + CO(g) + C(s) + energy
5. Combustion Reactions
• Ex) Write a balanced equation for the
complete combustion of C3H8.
• C3H8(g) + O2(g) --> CO2(g) + H2O(g) + energy
• C3H8(g) + 5 O2(g) --> 3 CO2(g) + 4H2O(g) + energy
• Your turn: Write a balanced equation for the complete
combustion of C8H18.
• Answer
Precipitation Reactions
• Most ionic compounds dissociate into
cations and anions when dissolved in
water.
• A complete ionic equation (basically a
double replacement reaction) shows
ionic compounds as free ions.
• In other words, write in the charges.
Precipitation reactions
Predicting the precipitate
• Use the chart on the back of your periodic table.
• Which of the following compounds are not soluble
–
–
–
–
–
Calcium Sulfate
Sodium Acetate
Silver Chloride
Aluminum Hydroxide
Potassium Phosphate
Precipitation Reactions
Complete Ionic Equations
• In an aqueous solution, substances exist as
free ions. The equation shows this.
• Example for AgNO3(aq) + NaCl(aq)
Ag+1(aq) + NO31-(aq) + Na1+(aq) + Cl1-(aq) --> AgCl(s) + Na+1(aq) + NO31-(aq)
Precipitation Reactions
Net Ionic Equation
• A net ionic equation indicates those ions
that took part in the reaction.
• Net ionic equation for the reaction from
the previous slide is:
• Ag1+(aq) + Cl1-(aq) --> AgCl(s)
Precipitation Reaction
• Example: Write a complete and net ionic
equation for the reaction of aqueous
solutions of iron (III) nitrate and sodium
hydroxide.
Fe3+(aq) + NO31-(aq) + Na+1(aq) + OH-(aq) --> Fe(OH)3(s) + Na+1(aq) + NO31-(aq)
• Fe3+(aq) + OH1-(aq) --> Fe(OH)3(s)
Precipitation Reactions
• Your turn. Write a complete ionic equation
and a net ionic equation for the reaction of
aqueous solutions of silver nitrate and
potassium sulfate.
• Answer
I. Molar Conversions
The Mole
• 1 mole of hockey pucks would
equal the mass of the moon!
• 1 mole of basketballs would fill a
bag the size of the earth!
• 1 mole of pennies would cover the
Earth 1/4 mile deep!
Molar Conversions
Converting from moles to grams to representative particles
and vice versa. Use the following conversion factor:
1 mole = 6.02 x 1023 representative units = molar mass (g)
or formula weight
Representative units a. ionic compounds are called formula units
b. molecular compounds are called molecules
c. atoms are called atoms.
Example of representative units 6.02 x 1023 atoms Cu
6.02 x 1023 molecules O2
6.02 x 1023 units NaCl
Molar Conversion Examples
• How many moles of carbon are in
26.0 g of carbon?
Molar Conversion Examples
• How many molecules are in 2.50
moles of C12H22O11?
Molar Conversion Examples
• Find the mass of 2.1  1024
formula units of NaHCO3.
Molar Conversion Examples
• Find the number of units of Iron
(III) Chlorate in 98.6 g of Iron (III)
Chlorate.
Moles in a Gas
• 1 mole of gas takes up 22.4 L of space at
standard temperature and pressure.
• Conversion factor - 1 mole = 22.4 L
– Remember this is for a gas only
• Standard Temperature and Pressure (STP)
– Temp = 0oC
– Pressure = 1 atm (atmosphere)
– 1 atmosphere is defined as the amount of
pressure the earth’s atmosphere places on you at
sea level
Calculations w/ molar volume
• Determine the volume, in liters, of 0.60
mol SO2 gas at STP.
• Answer – 0.60 mol SO2 22.4 L SO2
1 mol SO2
= 13 L SO2
Calculations w/ molar volume
Your Turn
• How many atoms of He are contained in your
party balloon if the balloon takes up 4.2 L of
space? Of course, this is one cold party, as it
would be held at STP.
• Answer -
Molarity
• Unit of Concentration
– There are many units of concentration
• Molarity is most useful to the chemist
moles
of
solute
M=
Liters of solution
Liters of solution means the total volume of
water and solute.
If I want a liter of solution I will not use a liter of
water.
Molarity Problems
You work them.
• A saline solution contains 0.90 g NaCl in exactly 100
ml of solution. What is the molarity of the solution?
Molarity Problems
You work them.
• How many moles of solute are present 1.5 L of 0.24
M Na2SO4?
Preparing a solution
• How would you make 500.0 ml of a 0.25 M solution of
copper (II) chloride?
• 0.25 M = mol/0.5000 L - change ml to liters and solve
for moles.
• You need 0.13 moles of CuCl2. Converting to grams
equals 17 grams.
• Final answer
– Take 17 grams of CuCl2 and dissolve in enough water to make
500.0 ml of solution.
• Dissolve the 17 grams in say 400 ml of water. Once the
CuCl2 is dissolved add water up to 500.0 ml.
Preparing a solution
your turn
• How would you prepare a solution of 0.40 M KCl? If a
volume is not given assume 1 L.
Making Dilutions
• Making dilutions from known
concentrations:
• M1 x V1 = M2 x V2
• Volume can be in liters or mL as long as
the same units are used.
Dilution Problems
• How would you prepare 1.00x102 mL of
0.40 M MgSO4 from a stock solution of
2.0 M MgSO4?
• 0.40 M x 100 mL= 2.0 M x V2
• V2 = 20 mL
• Answer - Take 20 mL of 2.0 M MgSO4
and dilute with enough water to make
100 mL of solution.
Dilution Problems
Your Turn
• How would you prepare 90.0 mL of 2.0
M H2SO4 from 18 M stock solution?
• Answer
Dilution Problems
Your Turn - 1 more
• If 250 mL of a 12.0 M HNO3 is diluted to
1 L, what is the molarity of the final
solution?
• Answer -
Percent Composition
% composition
your turn
• Hydroxide makes up what percent of
Calcium Hydroxide?
• Answer
Hydrates
• Hydrates are substances that contain
water within the crystalline structure of
the compound.
• The water is not chemically bound; it is
trapped within the crystal.
• Ex. FeSO4 . 7H2O
Empirical vs. Molecular
Formula
Calculating Empirical Formulas
– Lowest whole-number ratio of the atoms of
the elements in a compound
• C6H12O6 (glucose)
• The ratio that glucose normally has for
carbon:hydrogen:oxygen is 6:12:6.
• The lowest ratio that glucose has for
carbon:hydrogen:oxygen is 1:2:1 (each
number can be divided by the smallest
number in the ratio which is 6).
• The empirical formula for glucose is
CH2O since this is the lowest wholenumber ratio of atoms for that
compound.
– May or may not be the same as the normal
molecular formula of a compound
• Next - Calculating empirical formulas
What is the empirical formula of a
compound that is 25.9% nitrogen and
74.1% oxygen?
• If 25.9% of the compound is nitrogen and 74.1% of
the compound is oxygen, then a compound with a
mass of 100 g has 25.9 g of nitrogen and 74.1 g of
oxygen.
• To calculate the empirical formula, we need to relate
the moles of each atom in the compound, so we need
to convert the masses of the elements to moles.
25.9 g N
1 mol N
mol N 

 1.85 mol N
1
14.0067 g N
74.1 g O
1 mol O
mol O 

 4.63 mol O
1
15.9994 g O
This would mean that the ratio of nitrogen to oxygen
is N1.85O4.63.
We can divide each number in the ratio of N1.85O4.63 by
1.85 to get N1O2.50.
Since we cannot have 2.50 atoms of
oxygen, we must multiply through each number by 2 to
even it out, getting N2O5 as our empirical formula.
• In calculating empirical formulas, remember
that the number of atoms is a whole number.
If the number of atoms for an element is close
to a whole number (i.e., 2.1 or 2.2 or 2.8, or
2.9), you can usually round up or down to get
a whole number.
• If you should get a number of atoms closer to
2.33 or 2.5, multiply each number in the
formula by a number that gets that to a whole
number. For example, if you calculated 2.33,
you would multiply this by 3 to get a value of
7 for that number.
• Give it a try
• Determine the empirical formula for a
compound containing 7.8% carbon and
92.2% chlorine.
Empirical vs. Molecular
Formula
Calculating Molecular Formulas
• Although sometimes a molecular
formula may be the same as a
molecule’s empirical formula, like in
carbon dioxide (CO2), we have seen
that the empirical formula for glucose is
not the same as its molecular formula.
• One can determine the molecular
formula of a compound by knowing its
empirical formula and its mass.
• Next - Example
Calculate the molecular formula of the compound
whose molar mass is 180.1583 g and empirical formula
is CH2O.
We know that the molecular formula will have a molar
mass of 180.1583 g. We also know, by calculating the
gmm of CH2O, that CH2O has an empirical formula
mass (efm) = 30.0264 g CH2O.
Now, in order to figure out what we must multiply each
number in the empirical formula by, we must figure out by
what number we must multiply the empirical formula mass
to get the molecular formula mass.
To get from 30.0264 to 180.1583, 180.1583
30.0264
6
• Therefore, we must multiply each number of atoms in
CH2O by 6 to get the molecular formula of C6H12O6.
• You can double-check your answer by recalculating
the molar mass of C6H12O6.
• gmm C6H12O6 = 6 x 12.0111 g + 12 x 1.00794 g + 6 x
15.9994 g = 180.1583 g C6H12O6
• This agrees with the molar mass we were given, so
the molecular formula we calculated is correct.
• Give it a try
• Determine the molecular formula of a
compound that is 40.0% C, 6.6% H, and
53.4% O and the molar mass is 120.0g.
Example (toughy)
• 1.00 g of menthol on combustion yields
1.161 g of H2O and 2.818 g of CO2.
What is the empirical formula?
• Solution:
Stoichiometry
Calculations of quantities in chemical
reactions.
The use of ratios to calculate quantities
The five step process
• 1. Start with the balanced equation
• 2. Set up the problem - put down the tracks
• 3. Convert to moles if needed. This means you would
be given grams, representative units or liters.
• 4. Convert to moles of what you want. You will use
the mole ratio from the balanced equation.
• 5. Convert to what you are trying to find (grams,
liters, representative units) if needed.
Stoichiometry
Example Problem #1
• How many moles of ammonia are
produced when 0.60 mol of hydrogen
reacts with nitrogen?
Stoichiometry
Example Problem #2
• Your Turn
• How many moles of aluminum sulfide
are produced when 1.2 moles of
aluminum reacts with sulfur?
Stoichiometry
Example Problem #2
• Answer
Stoichiometry
Example Problem #3
• How many grams of ammonia will be
produced by reacting 5.40 g of
hydrogen with nitrogen?
Stoichiometry
Example Problem #4
• Your Turn
• How many grams of aluminum are
needed to react with 2.45 g of copper
(II) chloride?
Stoichiometry
Example Problem #4
• Answer
% Yield
• Definitions
• 1. Theoretical Yield
• The maximum amount of product that
can be formed from a given amount of
reactants
• In other words, the calculated amount
predicted through stoichiometry
% Yield
• Definitions
• Actual Yield
• The amount that is actually formed
when the reaction is carried out in the
laboratory.
% Yield =
actual yield
X 100
theoretical yield
% yield will never be over 100%
Most likely it will never even be 100%
Why will % yield never be 100%
• Advantageous to add an excess of an inexpensive reagent
to ensure that all of the more expensive reagents reacts
• Reactant may not be 100% pure
• Materials are lost during the reaction
– If a reactions takes place in a solution it may be
impossible to get all of the reactants or products out of
the solution
• If the reactions takes place at a high temperature,
materials may be vaporized and escape into the air
• Side reactions may occur
– Example Mg burned in air. Some of Mg reacts with
nitrogen reducing the amount of MgO produced.
• Loss of product when filtering or transferring
• If reactants are not carefully measured
% yield example problem
• In a reaction between barium chloride
and potassium sulfate, 3.89 g of barium
sulfate is produced from 3.75 g of
barium chloride. What is the percent
yield?
% yield example problem
answer
• BaCl2 + K2SO4 --> BaSO4 + 2 KCl
3. 75 g BaCl2 1 mol BaCl2 1 mol BaSO4 233.4 g BaSO4
208.3 g BaCl2 1 mol BaCl2 1 mol BaSO4
= 4.20 g BaSO4
3.89 g BaSO4 x 100 = 92.6 %
4.20 g BaSO4
% yield example problem
your turn
13.35 grams of magnesium hydroxide is
produced when 42.50 grams of
magnesium nitrate reacts with an
excess of aluminum hydroxide. What is
the percent yield?
% yield example problem
answer
Limiting Reagent
• 1. Limits or determines the amount of product that
can be formed
• 2. The reagent that is not used up is therefore the
excess reagent
• These types of problems require 2 sets of tracks.
Quantities of both reagents will be given. Therefore,
you need to find out which one is the limiting reagent.
Limiting Reagent
• One track to determine limiting reagent
• A second track to determine product
Limiting Reagent Example problem
• How many grams of copper (I) Sulfide can be produced when 80.0
grams of Cu reacts with 25.0 grams of sulfur?
• 2Cu + S --> Cu2S
• Pick a reactant and calculate how much of the other reactant is
needed.
80.0g Cu 1mol Cu 1mol S 32.1g S
63.5g Cu 2mol Cu 1mol S
= 20.2g S
So, 20.2 g of S is needed; 25.0g is supplied
Plenty of S; therefore, Cu is limiting reagent.
Use Cu to solve the problem
80.0g Cu 1mol Cu 1mol Cu2S 159.1g Cu2S
63.5g Cu 2mol Cu
1mol Cu2S
= 1.00x102 g Cu2S
Limiting Reagent Example Problem - Your Turn
• How many grams of hydrogen can be produced when 5.00g of Mg is
added to 6.00 g of HCl?
Limiting Reagent Example problem- Your Turn
• Acetylene (C2H2) will burn in the presence of oxygen. How many grams
of water can be produced by the reaction of 2.40 mol of acetylene with
7.4 mol of oxygen?
The Development of Atomic Models
• Democritus was a preSocratic Greek
philosopher (born
around 460 BC).
• Democritus was
originator of the belief
that all matter is made
up of various
imperishable, indivisible
elements which he
called "atomos", from
which we get the
English word atom.
•According to legend,
Democritus was supposed to
be mad because he laughed
at everything, and so he was
sent to Hippocrates to be
cured. Hippocrates pointed
out that he was not mad, but,
instead, had a happy
disposition. That is why
Democritus is sometimes
called the laughing
philosopher.
BB - Model
Dalton’s Model
More to come
Plum Pudding Model
•
•
•
•
•
Proposed by J. J. Thomson
(1856 - 1940), the discoverer of
the electron in 1897.
The plum pudding model was
proposed in March, 1904 before
the discovery of the atomic
nucleus.
In this model, the atom is
composed of electrons
surrounded by a soup of positive
charge to balance the electron's
negative charge, like plums
surrounded by pudding. The
electrons were thought to be
positioned throughout the atom.
Electrons could move like letters
in alphabet soup
Instead of a soup, the atom was
also sometimes said to have had
a cloud of positive charge.
Thomson's model was compared
(though not by Thomson) to a
British treat called plum pudding,
hence the name. It has also been
called the chocolate chip cookie
model, but only by those who have
not read Thomson's original paper
Nuclear Model
• The Gold foil experiment
or the Rutherford
experiment was an
experiment done by
Ernest Rutherford (1871 1937) in 1909. This
experiment discovered the
nucleus.
• Led to the downfall of the
plum pudding model of
the atom.
• Alpha particles (positive
particles--Helium Nuclei)
were shot at gold foil.
• Particles passed through
the gold foil. A few shot
back.
• Conclusions:
1. Atom is mostly empty space
2. Dense center called the nucleus
3. Electrons were stuck surrounding
the nucleus.
Planetary Model
• Introduced by Niels Bohr,
a Danish physicist (1885 1962), in 1913.
• Because of its simplicity,
the Bohr model is still
commonly taught to
introduce students to
quantum mechanics.
• The Bohr model depicts
the atom as a small,
positively charged nucleus
surrounded by waves of
electrons in orbit — similar
in structure to the solar
system, but with
electrostatic forces
providing attraction, rather
than gravity.
"The opposite of a correct statement is a
false statement. But the opposite of a
profound truth may well be another
profound truth." Niels Bohr
Quantum Mechanical Model
• Erwin Schrödinger (August
12, 1887 – January 4, 1961)
• An Austrian physicist,
achieved fame for his
contributions to quantum
mechanics, especially the
Schrödinger equation, for
which he received the Nobel
Prize in 1933.
• This model is based on
probability
• Where are you going to find
and electron 90% of the time.
• Atom is viewed as a fuzzy
cloud.
• Schrödinger equations create
electron clouds (orbitals) with
specific shapes.
Main Points For “Atoms” Video
• What is the key to understanding atomic structure?
• The discovery of what particle is associated with the
Crook’s Tube?
• What did Rutherford expect to happen in the gold foil
experiment?
• What was Rutherford’s genius?
• What conclusions did Rutherford draw from the Gold
foil experiment?
• How much smaller is the nucleus than the electron
cloud?
• What determines the shape of the electron cloud?
Dalton’s Atomic Theory
• Democritus - Greek philosopher who
first suggested atoms
• John Dalton (1766-1844)
• Studied ratios in which elements
combine
• Dalton put together the first atomic
theory
Dalton’s Atomic Theory
• All elements are composed of tiny indivisible particles called
atoms
• Atoms of the same element are identical. The atoms of any one
element are different from those of any other element
• Atoms of different elements can physically mix together or can
chemically combine with one another in simple-whole number
ratios to form compounds
• Chemical reactions occur when atoms are separated, joined, or
rearranged. Atoms of one element, however, are never changed
into atoms of another element as a result of a chemical reaction
Finding how many subatomic
particles for each atom
• Atomic Number - whole number on p.t.
• Gives the number of protons
• Atoms are electrically neutral; there,
positives equal negatives.
• Atomic number also equals number of
electrons
Finding how many subatomic
particles for each atom
• Mass number = protons plus neutrons
– Mass number = p+
+
no
• Mass number is not found on the
periodic table
• So, nO = mass number - p+
• If carbon has a mass number of 14, how
many e-’s, p+’s, and no’s does it have?
Symbols
• Mass number is in top left and atomic
number is in bottom left
9Be
• 16O
8
4
How many subatomic particles in each?
Oxygen Beryllium -
Isotopes
• Atoms with the same number of protons
but different numbers of neutrons
• Ex) Carbon 12 vs. Carbon 14
• These atoms have a different mass
• Chemically alike because still have the
same number of protons
Isotopes of Hydrogen
• Hydrogen -1 simply called hydrogen
• Hydrogen - 2 called deuterium
• Hydrogen - 3 called tritium
Development of AMUs
•
•
•
•
•
•
Atomic Mass Units (AMUs)
Protons have a mass of 1 amu
1.67 x 10-24 g
Neutrons have a mass of 1 amu
Electrons have a mass of 0 amu
9.11 x 10-28 g
Atomic Mass
• The weighted average mass of the isotopes
in a naturally occurring sample of the element
• Don’t confuse with “mass number”
• To calculate atomic mass you need 3 pieces
of information
• 1. The number of stable isotopes
• 2.The mass of each isotope
• 3.The natural percent abundance of each
isotope
Atomic Mass
• Example Problem - Calculate the atomic mass for
element X. One isotope has a mass of 10 amus (10X)
and is 20% abundant. The other has a mass number
of 11 amus (11X) and an abundance of 80%.
• To solve: Multiply the mass number times the
abundance than add them together.
Atomic Mass
• 10 x 0.20 = 2.0
• 11 x 0.80 = 8.8
• Add 2.0 + 8.8 = 10.8
– The atomic mass of element X is 10.8
amus
Atomic Mass
• Your turn. Solve:
– What is the atomic mass of Element Z?
The isotopes are 16Z, 17Z, 18Z; with percent
abundances of 99.759, 0.037, 0.204.
Atomic Mass
• Answer
Atomic (Hotels) Theory
• Floors of the hotel are known as energy levels
• The period corresponds to the floor. The number of
the floor is called the Principle Quantum Number
• Energy levels are divided into sublevels. These are the
rooms of our atomic hotel. There are different types of
rooms.
• Atomic Orbital - Region of high probability of finding an
electron
• There are 4 types of rooms (orbitals). s, p, d, f rooms
• s - Spherical p - dumbbell d- clover leaf f - too
complex to describe.
• s - superior p - preferred d- desirable f- fair
– Sharp
principal
diffuse
fundamental
Atomic (Hotels) Theory
•
•
•
•
•
Three Managers each with their own rule:
1. Aufbau principle
Rooms closest to the basement are checked out first
Electrons enter orbitals of lowest energy first
2. Pauli Exclusion Principle - no more than 2 per
room
• An atomic orbital contains a maximum of 2 electrons
• Roommates must have opposite spins to share a
room
Atomic (Hotels) Theory
• 3. Hund’s Rule
• Similar rooms must have an occupant before pairing
up as roommates
• When electrons occupy orbitals of equal energy, one
electron enters each orbital until all the orbitals
contain one electron with parallel spins
S-orbitals of the 1st and 2nd energy levels.
P-orbital
2nd energy level
The anatomy of the periodic table
•
•
•
•
•
•
•
•
•
Get out your periodic tables
Know where the following are on your periodic table (p.t)
Group or Family
Period
Group A (representative elements)
Group B
Metals
Nonmetals
Metalloids (Semimetals)
– Note - aluminum is not considered a metalloid
The anatomy of the periodic table
• Know where the following are on your periodic table
(p.t) continued
• Transition metals
• Inner transition metals
• Alkali metals
• Alkaline metals
• Halogens
• Noble gases
• Atomic Number
• Atomic Mass
Focus Questions to upcoming video
8 Questions
• What is the periodic law?
• Who is Dmitri Mendeleev and what did he do?
• How is today’s periodic table different from
Mendeleev’s?
• What characteristic is common among the noble gases
• What are the names for vertical columns and horizontal
rows
• What characteristics are common amongst group 1A?
• What are the 2 most important things about the periodic
table?
• Why is fluorine the Tyrannosaurus rex of the periodic
table?
-1 0
10
m
Atoms
Nucleus (protons and neutrons)
Space occupied by electrons
Proton
Neutron
10
-1 5
m
Periodic Table and Electron Configurations
• Build-up order given by position on periodic table; row
by row.
• Elements in same column will have the same outer
shell electron configuration.
The relation between orbital filling and the periodic table
Electron Configuration
• Orbitals have definite shapes and
orientations in space
(insert Fig 2.11 of text)
(if it will not all fit on one screen, put part
(a) on one screen and part (b) on the
next )
Orbital occupancy for the first 10 elements, H through Ne.
Trends in the
Periodic Table
Atomic radii of the maingroup and transition
elements.
Trend for atomic radii
• Left to right atoms get smaller
• Why?
– Increase in nuclear charge
– More protons and more electrons means greater
electrostatic attractions (stronger magnet)
• Top to bottom atoms get larger
• Why?
– Increase in energy levels (You are adding floors to your
hotel). Electrons are further from the nucleus
Atomic Radius
• Atomic radii actually
decrease across a row in
the periodic table. Due
to an increase in the
effective nuclear charge.
• Within each group
(vertical column), the
atomic radius tends to
increase with the period
number.
Atomic Radii for Main Group
Elements
Trend for Ion Size
• Ion is a charged atom.
• Metals lose electrons and nonmetals gain electrons to create
ions.
• Cations are pawsitive (positive) and Anions are negative.
• Cations are smaller than their corresponding atom. Why?
• Loss of electrons means the positive nucleus has a greater
attraction on the remaining electrons
• Anions are larger than their corresponding atom. Why?
• Gain of electrons means the nucleus has less attraction for the
electrons as well as the electrons are repulsing each other
causing an increase in the size of the electron clouds
Radii of ions
This is a “self-consistent” scale based
on O-2 = 1.40 (or 1.38) Å.
Ionic radii depend on the magnitude
of the charge of the ion and its
environment. (more later)
Positively charged ions are smaller
than their neutral analogues because
of increased Z*.
Negatively charged ions are larger
than their neutral analogues because
of decreased Z*.
Same periodic trends as atomic
radii for a given charge
Trend for ion size
• Decrease across a period then jumps in
size at nonmetals and continues to
decrease
• Increases on the way down a group as
you are adding energy levels (electrons
are farther from the nucleus)
Ionization energy
• The energy required to remove an
electron
First ionization
energies of the
main-group
elements
Trends in the
Periodic Table
Ionization Energy
• Ionization energy is a periodic property
Ionization energy
• In general, it increases across a row. Why?
• increasing attraction as the number of protons in the
nucleus increases (stronger magnet)
• it decreases going down a group. Why?
• Outer shell electrons are further from the nucleus so
less electrostatic attraction. Nucleus has less pull on
them.
• Shielding also plays a factor.
6) The trend across from left to right is
accounted for by a) the increasing nuclear
charge.
Electronegativity - This is the most
important trend to understand for this class.
• The tendency for an atom to attract electrons
when chemically bonded.
• Same trend as ionization energy.
– In general, it increases across a row. Why?
– increasing attraction as the number of protons in
the nucleus increases (stronger magnet)
– it decreases going down a group. Why?
– Outer shell electrons are further from the nucleus
so less electrostatic attraction. Nucleus has less
pull on them. Shielding also plays a factor.
Trends in three atomic properties
See chart in book for
summary
Check for understanding
• Which of the following atoms has the
largest atomic radii, ion size,
electronegativity, and ionization energy
• Na, Mg, K, Ca, S, Cl, Se, Br
Bonding
Ionic
Metallic
Covalent
Valence Electrons
• Atoms in a group behave similarly
because they have the same number of
valence electrons.
• Valence electrons - electrons in the
highest occupied energy level
• To find the number of valence electrons
just look at the group number
Lewis Electron Dot Structures
• Symbol of element with dots around it representing
valence electrons
• Example:
C
 2 electrons per side totaling 8
 No pairs until each room has an electron
Lewis Electron Dot Structures
• Example:
O
Unshared
Pair or Lone
Pair
• Pairs of electrons are adjacent not
across from one another. This will help
with identifying shape.
Octet Rule
• In forming compounds, atoms tend to achieve the
noble gas electron configuration.
• They lose, gain, or share electrons with another atom
to achieve 8.
• Metals lose electrons leaving a complete octet in the
lower energy level
• Nonmetals gain electrons to fill the energy level to
achieve 8.
Electron Configurations of Ions
•
•
•
•
•
•
•
•
•
•
•
Na - 1s22s22p63s1 Na1+ - 1s22s22p6
O - 1s22s22p4
O2- - 1s22s22p6
Write the following on your periodic table.
Group 1A - 1+
- loses 1 electron
Group 2A - 2+
- loses 2 electrons
Group 3A - 3+
- loses 3 electrons
Group 4A - Depends on the atom
Group 5A - 3- gains 3 electrons
Group 6A - 2- gains 2 electrons
Group 7A - 1- gains 1 electron
Group 8A (0) - does not form ions
VSEPR Theory
•
•
•
•
•
•
Valence Shell Electron Pair Repulsion Theory
“Electron pairs around atoms tend to be as far apart as possible.”
Similar charges (I.e., negative charges from electrons) tend to repel each other
and want to be spaced apart at maximum angles.
Used to predict molecular geometries
Bond angles
– Angles between bonds
– Spacing apart as far as possible
Lone pairs of electrons will repel bonded atoms a bit more than expected
toward each other around the central atom
Lewis Dot Structure - A symbolic description of the
distribution of valence electrons in a molecule. Dots are
used to represent individual electrons and lines are used to
represent covalent bonds. Lines are not drawn for ionic
bonds!!!
Drawing Lewis Dot Structures
1. Add up the total number of valence electrons in the
molecule by totaling the valence electrons on each atom in
the molecule or polyatomic ion.
For example let’s work the following together:
PO43-, O3 , BrF5
• 2. Draw the skeleton structure of the molecule or
polyatomic ion in which the covalent bonds between
the atoms are drawn as single lines. Each bond
equals two valence electrons. If the molecule has
more than two atoms, the atom with the lowest
electronegativity is generally the central atom and is
written in the middle.
• 3. Distribute valence electrons around the outer
atoms as nonbonding electrons until each atom has a
complete outer shell (i.e. 8 electrons except for H
which has only 2 valence electrons).
•
4. Add the remaining valence electrons to the
central atom.
• 5. Check the central atom.
•
* If the central atom has eight electrons surrounding it, the
Lewis Structure is complete.
•
* If the central atom has less than eight electrons, remove a
nonbonding electron pair from one of the outer atoms and form a
double bond between that atom and the central atom. If needed,
continue to remove nonbonding electron pairs from the outer atoms
until the central atom has a complete octet.
•
* If the central atom has more than eight electrons, then this
means that the central atom has expanded its valence shell to hold
more than eight electrons. This is allowed for atoms with valence
shells in the third energy level or higher (i.e. in or beyond the third
period of the periodic table).
• Other notes
– Length of bond
• Single bonds are longer than double which are
longer than triple.
• C-C =154 pm, C=C =134 pm, C≡C =120 pm)
– Energies of bonds
• Triple bonds have more energy than double and
double have more energy single.
• C-C =348 kj/mole, C=C =614 kj/mole,
• C≡C =839 kj/mole)
• When the electron number is odd the
octet rule is broken. Example NO:
Covalent Bonding
Polar Bonds and Molecules
Covalent Bonding
-- Polar Bonds and Molecules -Bond Polarity
• “The Tug of War”
– The pairs of electrons that are bonds between atoms are pulled
between the nuclei of the atoms in a bond.
– The electronegativities of the atoms determines who is winning
Yet; there is no winner. The tug of war never ends.
• Classifications for Bonds
– Nonpolar covalent
• When atoms pull the bond equally
• Happens with two atoms of equal electronegativity, most often
using the same atoms
• Examples: H2, O2, N2
– Polar covalent
• When atoms pull the bond unequally
• Happens with two atoms of different electronegativities
• Example: HCl, HF, NH
Covalent Bonding
-- Polar Bonds and Molecules -Bond Polarity
•
•
In a polar molecule, one end of the molecule is slightly more electronegative
than the other atom, resulting in one atom being slightly negative (-) because
of higher electronegativitiy, and the other atom being slightly positive (+)
because of lower electronegativity.
 is known as a partial charge since it is much less than 1+ or 1- charge.
Covalent Bonding
-- Polar Bonds and Molecules -Bond Polarity
• Electronegativities and Bond Types
– H: 2.1 Cl: 3.0 Since hydrogen is less, it will have the positive
partial charge while chlorine has the negative partial charge.
– 3.0 – 2.1 = 0.9 HCl is polar covalent.
0.0 – 0.1 difference
Nonpolar covalent bond
H – H (0.0 difference)
0.1 – 1.7 difference
Polar covalent bond
H – Cl (0.9 difference)
1.7 + difference
Ionic bond
Na+Cl- (2.1 difference)
Covalent Bonding
-- Polar Bonds and Molecules -Polar Molecules
•
•
Dipole
– Molecule that has two poles
– Example: HCl from the previous page
Polar vs. Nonpolar
H2O and CO2
Both have 3 atoms; yet,
One is polar and one is
nonpolar.
Why?
Structure (with bond
polarity) determines the
molecules polarity.
3 video clips coming up
• Get out your bonding sheet of chemistry
• Yes, we are going to discuss more on bonding so try
and hold your enthusiasm as rioting is not tolerated
and things need to be accomplished.
• Bonding appreciates your cooperation and will sign
autographs at the end of the period.
• Thank you.
Intermolecular Attractions
Attractions Between Molecules
• van der Waals forces
– Two types: dispersion forces and dipole interactions
• Dispersion forces
– Weakest of all molecular interactions
– Caused by movement of electrons
– Occurs in the BrINClHOF’s
Intermolecular Attractions
Attractions Between Molecules
2nd van der Waals force
• Dipole interactions
• Occurs when polar molecules are attracted to one
another
• Partial charge (+) of one polar molecule is
attracted to the opposite partial charge (-) of
another molecule
Intermolecular Attractions
attractions between molecules
• Hydrogen bonding
– Hydrogen covalently bonded to a very electronegative atom is also
weakly bonded to an unshared electron pair of another
electronegative atom
– Example: water
Short Lab - No lab report
• In the back in a tray is a micropipet and a penny.
Place 1 drop of water on the penny and then take a
guess as to how many drops of water you can fit on a
penny. Write your guess on the board at the front of
the room. Place your name next to your guess. Then
count the drops of water you fit on the penny before it
overflows. Record this count on the board next to
your guess. As your water drop grows, watch it from
the side. Clean up and have a seat at your desk.
Gases
• There are four variables that affect a
gas.
• 1. Pressure
• 2. Volume
• 3.Temperature
• 4. Number of molecules
The variables
• Pressure units - there are many units for pressure.
– kPa - kilopascal (101.3)
– Atm - atmospheres (1 )
– mm Hg - millimeters of Hg (760) - torr
•
•
•
•
Volume is measured in Liters
Temperature is in Kelvin
K = oC + 273 or oC = K - 273
If the Kelvin temperature doubles the K.E. doubles.
The pressure-volume relationship
Boyle’s Law
• Pressure and volume are inversely
related.
• One goes up the other goes down
• P1 x V1 = P2 x V2
Boyle’s Law
sample problem
• A high-altitude balloon contains 30.0 L
of helium gas at 103 kPa. What is the
volume when the balloon rises to an
altitude where the pressure is only 25.0
kPa?
• 103 kPa x 30.0 L = 25.0 kPa x V2
• V2 = 124 L
Boyle’s Law
sample problem
• Your turn
• Your birthday balloon travels with you
from Lincoln to Denver. The balloons
volume is 4.0 L with an atmospheric
pressure of 101.3 kPa. You arrive in
Denver where the atmospheric pressure
is 90.0 kPa. What is the new volume of
your balloon?
Boyle’s Law
sample problem
• Answer
The temperature-volume relationship
Charles’s Law
• Volume and temperature have a direct
relationship.
• One goes up the other goes up
– One goes down the other goes down
V1
T1
=
V2
T2
Charles’s Law
sample problem
• A balloon inflated in a room at 24oC has
a volume of 4.00 L. The balloon is then
heated to a temperature of 58oC. What
is the new volume if the pressure
remains constant?
4.00L = V2
297 K 331 K
V2 = 4.46 L
Charles’s Law
sample problem
• Your turn
• A balloon is inflated in a room at 24oC
and has a volume of 4.00 L. The balloon
is placed in a freezer and then removed
the volume is now 3.25 L. What was the
temperature of the freezer in Celsius?
Charles’s Law
sample problem
• Answer
The Temperature-Pressure Relationship
Gay-Lussac’s Law
• The pressure of a gas is directly
proportional to the temperature of a gas
• Temperature goes up; pressure goes up
P1 = P2
T1 T2
Gay-Lussac’s Law
example problems
• The gas left in a used aerosol can is at
a pressure of 103 kPa at 25oC. If the
can is thrown into a fire, what is the
pressure of the gas when it reaches
928oC?
103 kPa = P2
298 K
1201 K
P2 = 415 kPa
Gay-Lussac’s Law
example problems
• Your turn
• A container of propane has a pressure
of 108.6 kPa at a morning temperature
15oC. By mid afternoon the temperature
has reached 32oC. What is the pressure
inside the propane tank?
Gay-Lussac’s Law
example problems
• Answer
The combined gas law
P1 x V 1
T1
=
P2 x V 2
T2
The combined gas law
example problem
• The volume of a gas-filled balloon is
30.0 L at 40oC and 153 kPa. What
volume will the balloon have at STP?
153 kPa x 30.0 L = 101.3 kPa x V2
313 K
273 K
V2 = 39.5 L
The combined gas law
example problem
• Your turn
• A gas-filled balloon is 25.0 L at 35oC
and 145 kPa. What is the temperature if
the volume increases to 28.0 L and a
pressure of 152 kPa?
The combined gas law
example problem
• Answer
The Ideal Gas Law
PV=nRT
•
•
•
•
•
P = Pressure
V = Volume
n = Number of Moles
R = ideal gas constant
T = temperature in Kelvin
R -The ideal gas constant
•
•
•
•
Depends on unit of pressure
0.0821 L . Atm / K . mol
62.4 L . mmHg / K . mol (torr is mm Hg)
8.31 L . kPa / K . mol
Ideal Gas Law
example problem
• Calculate the pressure of 1.65 g of helium gas at 16.0oC and
occupying a volume of 3.25 L?
• You will need g to moles and Celsius to Kelvin:
• 1.65 g He 1 mole He
•
4.0 g He
= 0.413 mol He
• K = oC + 273 ; 16. 0 + 273 = 289 K
• For this problem you will need to pick an R value. For this
problem I will choose to use the R value containing kPa.
Ideal Gas Law
example problem
• P x 3.25 L = 0.413 mol x 8.31 kPa . L x 289 K
•
mol . K
• Do the algebra and solve; if you do it right, guess
what? You get the answer right. Neat concept, huh?
Maybe your mommy will give you a cookie.
• = 305 kPa
• Your turn
• How many moles of gas are present in a sample of
Argon at 58oC with a volume of 275 mL and a
pressure of 0.987 atm.
Ideal Gas Law
example problem
• Answer
Surface Tension
Intermolecular Forces Bulk
and Surface
Phase Changes
Energy Changes Accompanying Phase Changes
• All phase changes are possible under the right
conditions (e.g. water sublimes when snow
disappears without forming puddles).
• The sequence
heat solid  melt  heat liquid  boil  heat gas
is endothermic.
• The sequence
• cool gas  condense  cool liquid  freeze  cool solid
is exothermic.
Phase Changes
Heating Curves
Heating Curve Illustrated
Vapor Pressure
Explaining Vapor Pressure on
the Molecular Level
• Dynamic Equilibrium: the point
when as many molecules
escape the surface as strike
the surface.
• Vapor pressure is the pressure
exerted when the liquid and
vapor are in dynamic
equilibrium.
Vapor Pressure
Volatility, Vapor Pressure, and
Temperature
• If equilibrium is never established then the
liquid evaporates.
• Volatile substances evaporate rapidly.
• The higher the temperature, the higher the
average kinetic energy, the faster the liquid
evaporates.
Liquid Evaporates when no Equilibrium is
Established
Vapor Pressure
Vapor Pressure and Boiling Point
• Liquids boil when the external pressure equals the vapor
pressure.
• Temperature of boiling point increases as pressure
increases.
• Two ways to get a liquid to boil: increase temperature or
decrease pressure.
– Pressure cookers operate at high pressure. At high
pressure the boiling point of water is higher than at 1 atm.
Therefore, there is a higher temperature at which the food
is cooked, reducing the cooking time required.
• Normal boiling point is the boiling point at 760 mmHg (1
atm).
Water
Ice
• Liquid water’s density is greatest 4oC.
• Ice has a 10% greater volume;
therefore, lower density.
Ice
Why does ice behave so differently?
• As kinetic energy (speed of the
molecules) decreases, hydrogen bonds
hold water molecules in place.
• The water molecules are held in an
open framework creating a hexagonal
symmetry of molecules. This increases
the volume.
Aqueous Solutions
• Water samples that contain dissolved
substances.
• Solvent - dissolving medium
• Solute - dissolved particles
• Example - salt water - NaCl is the solute and
the water is the solvent
• Characteristics:
–
–
–
–
Homogeneous
Stable
They do not settle out
Both solvent and solute will pass through a filter
Describe the process of
solvation.
• Watch the next 2 video clips and then
answer the above question.
Other solution questions:
• What substances don’t dissolve in
water? Why?
• Why does grease dissolve in gasoline
and not water?
• How does soap work?
• How is the relationship summed up?
Electrolytes
• Compounds that conduct an electric
current in aqueous solution or molten
state.
• All ionic compounds are electrolytes
nonelectrolytes
• Do not conduct an electric current in aqueous
solution or molten
• Many molecular compounds are nonelectrolytes
– Ex sugar, rubbing alcohol
• Some very polar molecules become electrolytes
when dissolved in water.
• Why? Because they ionize in solution
– HCl + H2O --> H3O+ + Cl• Weak electrolyte - only a fraction of solute exists as
ions
• Strong electrolyte - large portion of the solute exists
as ions
• Henry’s Law
• - The solubility of a gas in a liquid is
directly proportional to the pressure of the
gas above the liquid.
• - pressure goes up, solubility goes up
• Mathematically stated
•
S1 = S2
•
P1
P2
• Example problem – you try
• If the solubility of a gas in water is 0.77 g/L
at 3.5 atm of pressure, what is its solubility
(in g/L) at 1.0 atm of pressure?
• Answer
• How does temperature affect the solubility
of a gas?
Thermochemistry
-- The Flow of Energy: Heat --
Thermochemistry: the study of heat
changes in chemical reactions
Chemical potential energy: energy stored
within the structural units of chemical
substances
Thermochemistry
-- The Flow of Energy: Heat --
Chemical System Types
System type
Endothermic
Exothermic
Description
System absorbing heat
from the surroundings
System releasing heat to
the surroundings
q (change
in heat)
q > 0
q < 0
Thermochemistry
-- The Flow of Energy: Heat --
Law of Conservation of Energy:
In any chemical or physical
process, energy is neither
created nor destroyed
Thermochemistry
-- The Flow of Energy: Heat -The calorie
• Expressed as a c (lower case)
• Quantity of heat needed to raise the temperature
of 1 g of pure water 1C
Calorie
• Expressed as a C (upper case)
• Dietary Calorie
• 1 Calorie = 1 kilocalorie = 1000 calories
Thermochemistry
-- An Intro Video --
Thermochemistry
-- The Flow of Energy: Heat --
Specific Heat
- a physical property (intensive property)
- describes how much heat a substance can
hold
Water vs Metal
- how does the temperature of a swimming
pool compare from 3 p.m. to 3 a.m.?
- how does the surface of your car hood
compare from 3 p.m. to 3 a.m.?
Thermochemistry
-- The Flow of Energy: Heat -Joule
• SI unit of heat and energy
• Raises the temperature of 1 g of pure water 0.2390C
• 4.184 J = 1 cal
Heat Capacity
• Amount of heat needed to increase the temperature of an object
exactly 1C
• Will change depending on the mass and chemical composition
Specific Heat
Quantity of heat needed to raise the temperature of
1g of substance 1oC
Thermochemistry
-- The Flow of Energy: Heat --
Specific Heat Capacity
Heat (q)
Mass (m)
specific heat capacity (C)
change in temperature (T)
q = mC T
Thermochemistry
-- The Flow of Energy: Heat -Example:
How many kilojoules of heat are absorbed when 1.00 L of water is
heated from 18C to 85C?
Solution:
q = mCT
q = 1000g x 4.18 J x 67oC
goC
q = 2.8E5 J 1 KJ
1000 J
= 280 KJ
Thermochemistry
-- The Flow of Energy: Heat -Example:
A chunk of silver has a heat capacity of 42.8 J/C. If the silver
has a mass of 181 g, calculate the specific heat of silver.
Solution:
q = mCT
42.8 J = 181g x C x 1OC
C = 0.236 J/goC
Thermochemistry
-- Measuring and Expressing Heat Changes -Your Turn:
The temperature of a piece of copper with a mass of
95.4 g increases from 20.0oC to 43.0oC when the
metal absorbs 849 J of heat. What is the specific
heat of copper?
Thermochemistry
-- Measuring and Expressing Heat Changes --
Calorimeter
Heterogeneous Aqueous Systems
1. Difference in types of heterogeneous aqueous
systems is particle size
Suspensions 1. particles settle out of solution
2. can be filtered
ex muddy water
Colloids 1. particles stay suspended in dispersion medium
2. many are cloudy - may look clear when diluted
3. exhibit Tyndall effect - scattering of light
ex - paints, glues, gelatin desserts, etc
Solutions - we already went over solutions
Colloidal Systems
Emulsions Colloidal
dispersion of liquids in
liquids
needs an emulsifying
agent
ex - soap
Properties of acids and Bases
• Taste
– Acids taste sour ex - lemons
– Bases taste bitter ex - soap
• Feel
– Acids feel like water; but have you ever
gotten fruit juice on a canker sore or cut
– Bases feel slippery ex - soap and water
Properties of acids and bases
• Reaction with metals
– Acids - Hydrogen gas is produced when
reacted with certain metals
– Bases - typically don’t react with metals
• Both are electrolytes
• React to form salt and water
• Milk of Magnesia (Magnesium
Hydroxide) is a base used to treat
excess stomach acid problems.
Definitions of Acids and Bases
• Arrhenius Acid/Base - focused on
products
– HCl + H2O --> H3O+ + Cl– NaOH + H2O --> Na+(aq) + OH-(aq)
• Acids form H+’s and Bases form OH-’s
Definitions of Acids and Bases
• Bronsted-Lowery Acid/Base - focused
on what happens during formation
– HCl + H2O --> H3O+(aq) + Cl-(aq)
• Acid - substance that donates a proton
• Base - substance accepts a proton
• In the above example, what is the base
and what is the acid?
• How about this example?
– NH3 + H2O --> NH4+(aq) + OH-(aq)
Definitions of Acids and Bases
• Lewis acid/base
– H+ + OH- ---> H2O
• Acid - a substance that can accept a
pair of electrons to form a covalent bond
• Base - a substance that donates a pair
of electrons to form a covalent bond
• What is the Lewis acid and base?
– AlCl3 + Cl- --> AlCl4-
Problem
• Write an equation for the ionization of
nitric acid and explain how it fits each
definition?
Hydrogen Ions from Water
This is the “bases” to start understanding pH
• Water is considered neutral
• Collision between water molecules can cause
a hydrogen ion to transfer from one molecule
to another.
H2O
H2 O
H3O+
hydronium ion
OHhydroxide ion
Self-Ionization of Water
Water self ionizes to the concentration of 1.0 x 10-7 mol/L.
When the concentration of each ion equals 1.0 x 10-7 mol/L
the solution is said to be neutral
Therefore, since water is considered neutral the
concentrations of the ions can be calculated through the
Ion product constant.
The Ion-Product Constant
Kw
• Notation
– [H+] - concentration of hydrogen ions
• Or hydronium ions
– [OH-] - concentration of hydroxide ions
• When [H+] and [OH-] are multiplied we get the ionproduct constant.
• Kw = [H+] x [OH-] = 1.0 x 10-14 (mol/L)2 or M2
• This is an inverse relationship.
– One goes up, the other goes down.
The Ion-Product Constant
Kw example problem
• If [H+] = 1.0 x 10-5 mol/L, is the solution
acidic, basic, or neutral? What is the
[OH-] of this solution?
• Answer
– Acidic - the [H+] is greater than 1.0 x 10-7
mol/L
– 1.0 x 10-5 mol/L x [OH-] = 1.0 x 10-14 M2
– [OH-] = 1.0 x 10-9 mol/L
The Ion-Product Constant
Kw example problem
• If [OH-] = 2.8 x 10-8 mol/L, is the solution
acidic, basic, or neutral? What is the [H+] of
this solution?
• Answer
The pH concept
• [H+] is cumbersome so the pH scale
was created.
• pH is the negative logarithm of the
hydrogen-ion concentration.
• pH = -log[H+]
Sample pH problems
1 of 3
• The hydrogen-ion concentration of a
solution is 2.7 x 10-10 mol/L. What is the
pH of the solution?
• Answer
Sample pH problems
2 of 3
• The pH of a solution is 6.8. What is the
[H+]?
• Answer
Sample pH problems
3 of 3
• What is the pH of a solution if the [OH-]
= 4.0 x 10-11 mol/L
• Answer
pOH
• pH + pOH = 14
• -log[H+] + -log[OH-] = 14
• Example – If the H+ is 7.2 x 10-9 mol/L
what is the pOH?
• 5.9
Naming Acids
And writing formulas
General Form - HX (X is an
anion or polyatomic ion)
• Rules - 3 of them
• 1. When anion ends in -ide, acid name
begins with hydro and -ide is changed
to -ic with the word acid
• Ex - HCl is hydrochloric acid
• Try - H2S
• Rule 2 • When anion ends in -ite, ending changes to ous with the word acid. (No hydro)
• Name these - H2SO3 , HNO2
• Rule 3 • When anion ends in -ate, ending changes to ic with the word acid
• Name these - HNO3 , HC2H3O2
Work backwards to get the
formula
• Write the formula for the following:
– Chloric acid
– Hydrobromic acid
– Phosphorous Acid
• Don’t confuse phosphorous and phosphorus
• Answers
Neutralization Reaction
• Base and acid react to produce a salt
and water
Titration
• A lab technique where a neutralization
reaction is performed to determine the
concentration of an unknown.
The anatomy of a titration:
• Standard solution - solution of known
concentration.
• End Point - the point at which neutralization
is achieved
• Indicator - chemical that changes
color with a change in pH.
We will be using phenolphthalein.
Clear in an acid pink in a base
You want a light pink
• Buret - Measurement device
Titration Calculations
4 steps
•
•
•
•
Start with the balanced equation
Find the moles in the standard solution
Set up ratio to find moles of unknown
Find molarity(mol/L) or volume
Titration Calculations
example problem
• A 25.75 ml solution of H2SO4 is neutralized by
18.23 ml of 1.0 M NaOH. What is the
concentration of H2SO4?
• H2SO4 + 2 NaOH --> Na2SO4 + 2 H2O
• 0.01823 L NaOH 1.0 mol NaOH 1 mol H2SO4
1 L NaOH
• = 0.35 mol/L H2SO4
2 mol NaOH 0.02575L H2SO4
Titration Problems
your turn
• What is the molarity of phosphoric acid if 15.0
mL of the solution is neutralized by 8.5 mL of
0.15 M NaOH?
Titration Problems
1 more your turn
• How many milliliters of 0.45 M hydrochloric acid
must be added to 25.0 mL of 0.15 M NaOH?