Pages 3 to 33
“Quantum Chemistry”
Target Completion Date: October 1
About Slide Icons
important information.
Pages with a PINK
background are
supplementary . Not
material for a test!
• You should either note or highlight items from this slide. Some items
from this slide WILL be on tests!
Very Important Sample Problems
• Always hand-copy important sample problems in your note book, and
refer back to them when doing assignments. Similar problems will be on
tests!
Look at this! (usually charts, diagrams or tables)
• You don’t need to copy or memorize this, but you must read and
understand the diagrams or explanations here. Concepts will be tested,
but not the details.
Information only. Don’t copy!
R
• This is usually background information to make a topic more interesting
or to fill in details, or to give examples of how to use a table. Not directly
tested.
Review Stuff
• Not part of the material you will be tested on, but you are expected to
remember this from grade 10. It may be indirectly tested.
Supplementary stuff
• Not material covered this year.
Measurement and Conversion Basics
• All sciences, including chemistry, depend on
observations. Measuring is an important part of
observing.
• There are many important types of measurement
in chemistry, but the four most important are...
–
–
–
–
Mass
Volume
Temperature
Time
kilograms (kg) or grams (g)
Litres (L) or millilitres (mL)
degrees Celsius (°C) or kelvins (K)*
seconds (s), minutes (min), hours (h)
* Degrees Fahrenheit (°F) are used in the United States, but are never used in
chemistry. Although kelvins are similar in magnitude to degrees celsius, kelvins do not
need the degree symbol (°)
Measuring Tools
• You must know how to use the following to
measure volume:
– Graduated cylinder
– Pipette
– Burette
Burette
R
Conversions
• You must be able to do ALL standard metric
conversions, especially:
– Litres to millilitres, millilitres to litres
– Grams to kilograms, kilograms to grams
Kill
Hector the
Decathlete
Until he’s
Deceased
with Centipedes and
Millipedes
Quick Conversions
Prefix
means
mega (M)
million
…
100000
…
10000
kilo (k)
1000
hecto (h)
100
deca (da)
10
… unit
1
deci (d)
0.1
centi (c )
0.01
milli (m)
0.001
…
0.0001
…
× 10-5
micro (μ)
× 10-6
The table on the left gives the eight most commonly
used prefixes in the metric system.
It also includes five rows that do not have prefixes.
The middle row is for the unit: metre, litre, gram,
newton, or any other legal metric unit.
This table can be used to quickly convert from one
metric amount to an equivalent. Make a copy of this
table on the margin of the front cover of your
notebook, and learn how to use it.
Lets do an example. Let’s find how many centimetres
there are in 2.524 km
Conversion: 2.524 km  ? cm
2 524 00 cm
km
Add extra zeros if necessary
There are five steps in the table between “kilo” and
“centi”, so we have to move the decimal five places to
the right. If we were going up the table we would
move left.
Answer: 2524 km = 252 400 cm
Other conversions
• Later this year you will need these:
– Temperature: degrees Celsius (℃) to kelvins (K)
• Add 273 to degrees Celsius to get kelvins.
• Subtract 273 from kelvins to get degrees Celsius.
– Pressure: kilopascals (kPa) to millimetres (mmHg)
• Multiply kilopascals by 760 and divide by 101.3 to get
mmHg
• Multiply mmHg by 101.3 and divide by 760 to get
kilopascals
R
Density
• Density is the relationship between the volume of an object
and its mass. Density is an important characteristic
property of matter.
• This is a review formula from last year:
m
or
ρ
V
Where: ρ = the density of the object, in g/cm3 or g/mL
m = the mass of the object, in g
V = the volume of the object, in cm3 or mL
ρw = 1 g/mL = 1 g/cm3
The density of water is 1 g/mL. This is not true
of other substances. Objects with less density
than water will float. Objects with greater
density will sink.
Solving Problems
• When solving Chemistry problems on a test or exam, it
is important not only to find the correct answer, but to
justify it. While solving the problem you should:
1. Show your data, the information you used to solve the
problem.
2. Show your work, including the formulas you used and
the substitutions you made.
3. Write an answer statement, a sentence that clearly
states your final answer.
4. Include the correct units for your answer. Never just give
a number—you must specify what the number means!
Showing Your Solution
1.
2.
3.
4.
5.
On the final examination, you must not only be able to find
solutions to problems, you must also justify your answers by
showing what you did.
divide the area where you will write your
solution into four sections.
In the first section, write your data and
the items you need to find
In the second section, write the
formula(s) you think you need to use.
In the third section, show your
calculations
In the final section, write your answer in a
complete sentence with the correct units.
Suggested Solution Method
Problem: A block of material has a length of 12.0 cm, a width of 5.0 cm,
and a height of 2.0 cm. Its mass is 50.0 g. Find its density.
Arrange your solution like this:
List all the information you
find in the problem,
complete with units, and
the symbols.
Data:
l
w
h
m
V
= 12 cm.
= 5.0 cm
=2.0 cm
=50.0g
=?
To Find:
ρ (density) =?
Write down all the formulas you intend to use:
Formulas:
V = lwh
𝑚
ρ=
𝑉
Show the substitutions you make, and enough of your calculations to
justify your solution:
Calculations:
V
=12cm x 5cm x 2 cm
= 120 cm3
𝜌
= 50 g / 120 cm3
= 0.417g/cm3
Always state your answer in a complete sentence, with appropriate units.
Answer: The density of the block is 0.417 g/cm3 (or 0.417 g/mL)
Problems on Conversions and Density
1. Convert the following:
a) 125 mL to L
b) 450 g to kg
c) 2.5 L to mL
d) 30 mL to L
e) 4500 mL to L
f) 1.35 kg to g
g) 75 mL to L
h) 0.035L to mL
i) 0.56L to mL
2. Find the density of a 4cm x 3cm x 2cm block that
has a mass or 480 g. Justify your solution.
3. Find the width of a cube whose density is 5 g/cm3
and whose mass is 135 g. Justify your solution.
Also: Do the worksheets entitled “Density” and
“Metric Conversions”
Appendix 5
Page 394
Overview: Uncertainty
Inherent Errors in Measuring Devices
In chemistry we often use instruments to
measure quantities. Unfortunately, all
instruments have some degree of inaccuracy or
error. I prefer to call this error “uncertainty”,
since it is not a mistake on the observer’s part,
but an unavoidable inaccuracy that comes from
the instrument. In this section we will see how
to write measurements that show we recognize
the limitations of our instruments.
App.5
Page 394
Absolute Uncertainty (AU)
In math, numbers are considered pure, abstract things. In math, 2.00, 2.0 and
2 are considered the same, they all represent number 2.
In science, numbers are considered to be measurements, and all
measurements have some degree of uncertainty. They are seldom
considered perfect!
The difference is in the precision of the instrument used to measure.
All instruments that we use to make measurements have an
inherent error or absolute uncertainty. On some instruments,
the absolute uncertainty is marked, on other instruments we
make the following assumption:
Assumption: The absolute uncertainty of a measurement is usually*
one half of a measuring instrument’s smallest graduation.
*at university level, or when using high-quality equipment AU measurements may be expected
to be one fifth of the measure between the smallest marking instead of one half!
App.5
Page 394
Example of Uncertainty.
• At first glance, the two
graduated cylinders here
seem identical, but look
closer.
• The first one has a
measurement of 32.0 ± 0.5 mL
• The second one 32.5 ± 0.5 mL
• It is NOT correct to say that
the first measurement is just
32 mL!
How to Record Absolute Uncertainty
Plus-minus sign
(32 .5 ±0.5 )mL
Doubtful, but Significant digit
Unit
A pair of parenthesis may be placed
around the measurement.
Absolute Uncertainty
• When you first look at a graduated cylinder, it appears
to contain 32 mL of liquid.
• Looking closer, you see it is about halfway between 32
and 33 mL, so you record the .5
• If you judged it to be ⅓ of the way you could write 32.3. If it was
below the 32 instead of just above it, you might record 31.5. If you
still see it as exactly 32 even after a closer look, then record it as
32.0
• Then you write the absolute uncertainty, the allowable
error of the instrument – usually* half the measure
between the smallest markings.
• In this case, the smallest markings represent one millilitre, so half
the measure ( 0.5 mL) is the uncertainty.
Adding and Subtracting with
Absolute Uncertainties
• Frequently we make two measurements and
subtract them to find a difference (Δ).
When we subtract numbers that have an
uncertainty we must ADD the absolute
uncertainty values!
• Eg. While doing a density experiment we add an
object to a graduated cylinder. The reading of the
cylinder changes from (20.5±0.5)mL to (24.0±0.5)mL.
The volume difference (ΔV) is (3.5±1.0)mL
• When adding two numbers with
uncertainties, we also ADD the uncertainty.
AUT = ΣAU
or
AUT = AU1 +AU2+…
Relative Uncertainty
App.5
Page 394
• Sometimes it is useful to know how much uncertainty
we have compared to the original measurement. To do
this we can calculate the relative uncertainty (RU).
• RU of a measurement equals the absolute uncertainty
divided by the absolute value of the original
measurement.
• The resulting decimal number is usually converted to a
percentage (by multiplying it by 100)
RU% =
percentage
relative
uncertainty
RU % 
Click here for details on UNCERTAINTY.
AU
measure
 100%
App.5
Page 394
Example of Relative Uncertainty
• The graduated cylinder has a reading of
(32.5±0.5) mL (absolute)
• To find its relative uncertainty, divide:
0.5 ÷ 32.5 = 0.01538461
• Round off to a reasonable number of
decimal places and convert to a percent:
0.015 x 100 = 1.5%
• Write it like this:
Parenthesis is NOT used for
32.5 ml ±1.5%
Relative Uncertainty
The nice thing about relative uncertainties it
that they show you how small your error
actually is.
Not in
TEXT!
Multiplying and Dividing with
Uncertainties
• When you multiply and divide measurements,
you cannot use Absolute Uncertainties.
• Instead, we must add the Relative
Uncertainties after we multiply or divide.
• Example: an object weighs (58.3±1.0)g and
has a volume of (32.5±0.5)mL. Find its density.
Use RU
• (58.3±1.0) g ÷ (32.5±0.5) mL = Can’t do it this way!
• find percents: 1.0 ÷ 58.3 x 100 =1.7%, 0.5 ÷ 32.5 x 100=1.5%
• 58.3 g ±1.7% ÷ 32.5 mL ±1.5% = 1.79g/mL ± 3.2%
• Answer: The density is 1.79g/mL ± 3.2%
Correct precision
• It is considered improper in science to imply
that a measurement is more precise than it
really is.
• If you have a graduated cylinder that is marked in 1 mL
increments, you can record it to between the two
smallest marks: eg. 32.0 ±0.5 mL or 32.5 ±0.5 mL are
acceptable readings.
• With the same graduated cylinder, it would be wrong to
write 32 ±0.5 mL or 32 ±0.5 mL or even 32.00 ±0.5 mL
• In science 32 mL, 32.0 mL and 32.00 mL have
different meanings with respect to
.
Exercise on Uncertainty
• Do the sheet “uncertainty”
• The sheet will be corrected in class.
• Procedures for an in-class exercise
• Make sure your first and last name are on the sheet.
• Complete as much of the sheet as you can in the time
allotted. Use a pencil or dark colour pen.
• When the time is up, follow the teacher’s instructions
regarding corrections. Correct with a red pen.
• When the sheet has been corrected, put it into your
assignment folder or duotang. Keep it here until at
least the end of the current term.
Overview: Significant Figures
Knowing how much to round an answer.
In the sciences, we have an particular way of
determining how much precision we need in the
observations and answers we record. The
method of rounding is called significant digits or
significant figures. There is a detailed section in
the appendix to your textbook on pages 394 to
397. Unfortunately, a few of the details given
there are, well… I won’t say wrong, let’s just call
them “uncertain”.
App.5
Page 395
Significant Figures
(A.K.A. Significant Digits)
• In science, we use significant digits as a guide
to how precise an observation is, and as a
guide to how much we should round off the
results we obtain by doing math with those
observations.
App.5
Page 395
Rules for Significant Figures
Interpreting Significant Digits
1. Non-zero digits are ALWAYS significant
2. Zeros between significant digits are ALWAYS
significant.
3. Zeros at the beginning of a number are
NEVER significant.
4. Zeros at the end of a number MAY be
significant, but only if trusted.
5. Exponents, multiples, signs, absolute errors
etc. are NEVER significant.
Examples of Rule 1, 2 and 3
Rule 1. Non-zero digits are ALWAYS significant.
1.234
has 4 significant digits
145
has 3 significant digits
19567.2
has 6 significant digits
Rule 2. Zeros between significant digits ARE significant.
1001
has 4 significant digits
5007.4
has 5 significant digits
20000.6
has 6 significant digits
Rule 3. Zeros at the beginning are NEVER significant.
007
has 1 significant digit
0.0000005 has 1 significant digit
0.025
has 2 significant digits
Explaining Rule 4
Rule 4. Zeros at the end of a number MAY be significant.
Your textbook says that they are ALWAYS significant, but this is contrary to
what most textbooks say.
If there is a decimal point, there is no problem. All textbooks agree, the
zeros are ALL significant.
3.00000
5.10
10.00
has 6 significant digits
has 3 significant digits
has 4 significant digits
If there is NO decimal, the situation is ambiguous, and a bit of a JUDGEMENT
CALL. If you trust the source to be precise, then you count all the zeros at the
end. If you have reason to believe the person was estimating, then you don’t
count the zeros at the end.
5000
has 1 to 4 significant digits
250
has 2 or 3 significant figures
123 000 000 has 3 to 9 significant figures
Estimated source
In a test situation, assume the
numbers are precise, unless
something in the question
states otherwise.
Trusted precise source
Rule 5
Rule 5: Exponents and their bases, perfect multiples,
uncertainties (error values), signs etc. are NEVER
significant.
6.02 x 1023 has 3 significant digits
504.1 mL x 3 has 4 significant digits
5.3 ±0.5 mL has 2 significant digits
– 5.432 x 10-5 has 4 significant digits
π × 8.45
has 3 significant digits
In each case, the blue part is significant, the green part is
NOT significant.
Note: The term Significance in this usage is not the same as importance. A digit may
be “insignificant” but still very important. The significant digits guide you to the
correct way of rounding numbers to show precision. The insignificant digits may serve
as “placeholders”, making sure the decimal point is in the right place. An important
job indeed, but not one that adds to the precision of the answer.
Not in
TEXT!
Avoiding Ambiguity
• We mentioned before that measurements ending in zeros
with no decimal were ambiguous. Their accuracy depends
on how they were measured, and that doesn’t always
show up in the number.
• For example, if you measure 200 mL in a cylinder with
markings of 1 mL it will be more accurate than if you
measured it in one with markings 10 mL apart, and much
better than a beaker whose markings were 100 mL apart.
How can we show someone reading
our lab notes the number of what
our 200 mL really means?
One answer is Scientific Notation!
Same Number, Different Precision
Number
Precise to
200.000
6 significant digits
200.00
5 significant digits
200.0
4 significant digits
200 *
Ambiguous, 1 to 3 SD*
2.00 x 102
3 significant digits
2.0 x 102
2 significant digits
2 x 102
1 significant digit
*This could represent one significant digit,
or two significant digits, or three significant
digits depending on how precise the
measuring equipment was. If I am careless
enough to write a number like this on a test,
you should assume I mean 3 S.D., but you
have my permission to point out my
mistake!
Avoid using numbers like 200 mL.
Instead write them in scientific
notation.
2.0x102 mL means you measured it
to the nearest 10 mL (2 S.D.)
2.00 x 102 mL means you measured
it to the nearest 1 mL (3 S.D.)
2.0000 x102 mL means you
measured it to the nearest 0.01
mL… a very fine level of accuracy
indeed!
Another way of showing the difference is to
include the absolute uncertainty!
Math with Significant Figures
• Adding and Subtracting:
• All units must be the same (can’t add different units!)
• Line up all the measurements at their decimal points.
• Round off all numbers to match the shortest number of
decimals.
Decimals lined up
• Add or subtract as normal.
Round off
Example: add the following measurements.
This unit is not the
same as the
others!
(litres vs. millilitres)
5345.8
5345.7 6mL
mL
5.34576 L
55.1 43mL
547.1 mL
55.1 43mL
mL
547.1 mL
5948.0 mL
The answer is 5948.0 mL. Note that the answer has 5 sig. digits, even though
one of the measurements had only 4 sig. digits. This can happen with addition.
Math with Significant Figures
• Multiplication and Division:
Weakest
measurement
only 3 S.D.
• Different units may be multiplied or divided if there is a
formula to justify it.
• The main rule in multiplying and dividing is that you
cannot have an answer with more significant digits than
your “weakest” measurement (the one with fewest
significant digits)
• After doing the math, round off your answer to match
the weakest measurement.
You are the
Justification:
Multiply 2.53 g/mL by 75.35 mL
2.53 x 75.35 = 190.6355
Answer has
only 3 S.D.
=191 g
m=ρV
weakest link.
Goodbye!
About the unit:
𝑔
x mL = g
𝑚𝐿
Math with Significant Figures
• Perfect numbers
• Occasionally we consider a number to be perfect. For example, if
you are told to “double a quantity” the 2 you multiply by is
considered perfect. It does not affect the significant digits of your
answer, neither increasing or decreasing them. Mole ratios in
stoichiometry are also considered perfect, as are universal
constants like pi. Perfect numbers have no units.
• Other operations
• Generally, use the same rule as for multiplying for square roots,
exponents etc. That is, your answer can have no more significant
digits than your weakest measurement.
• Mixed operations
• When doing mixed operations in science, you will usually do the
additions or subtractions first (there should be brackets around
them), then the other operations.
Problems on Significant Figures
1. How many significant digits are in each measurement:
a) 123.45 mL
c) 007 spies
e) 0.0023 m
b) 4.500 x103 mL
d) times 5
f) 4000 kg
2. A Coulter counter is a device which counts the blood cells in a
sample as they pass through a beam of light. A laboratory
technician records 20000 wbc in a blood sample. At a
demonstration a reporter says there were 20000 protesters.
Both numbers are the same, which one has more significant
figures? Why?
3. Find the volume of a prism that measures 2.3 cm by 3.55
cm by 2.14159 cm.
4. Add the measurements: 2.500 kg, 354.2 g, 153.78 g
Also: Do the worksheet entitled “Significant Figures”
Review
1
Periodic Classification
Overview
The periodic table is a useful arrangement of the
elements, into regions, families and periods that
have important meanings. It is also a source of
much additional information about the elements.
With careful interpretation of the table, we can find
the number of protons an atom has, the
approximate number of neutrons, and the
arrangement of electrons in the atom and in its ions.
R
0.1.1
Page 4
Topic 1: Organization of Matter
• 0.1.1 Atoms and Molecules
O
– All matter is composed of atoms.
– The atoms that make up most matter are
assembled into molecules.
O
C
H
CO2
H
N
NH3
H
• A molecule may contain one atom, or it may contain
several thousand atoms, or any number between.
– A molecule is represented by its formula
One
atom
Ne
Ne
• Water molecules, for example, are represented by the
formula H2O, shown below:
• A few large molecules have abbreviations,Cl not
S
several thousand
atoms like DeoxyriboNucleicAcid
formulas,
DNA
2 atoms of
hydrogen
O
H
H2O
H
1 atom of
oxygen
Cl
SCl2
36
Page 4
cation
0.1.2
• Chemical Formulas and Ions
Na
Na+
Cl–
Cl
anion
– Some matter is formed from ions instead of
normal atoms or molecules.
• For the most part, we treat ions like regular atoms,
and ionic compounds like molecules but there are a
few very technical differences.
– Ions are atoms or clusters of atoms that have
become positively or negatively charged by losing
or gaining one or more electrons.
Notice the
slightly
stronger
wording with
respect to
metals than
nonmetals!
• Positive ions are called cations (ca+ions),
• Negative ions are called anions (aNions)
• Metals always form cations (+), non-metals usually
form anions (-)
37
Differences between ionic and
covalent compounds
Ionic Compounds
Covalent (molecular) Compounds
Ionic bonds “give” or “take” electrons
Covalent bonds “share” electrons
Ionic compounds don’t have distinct
molecules. Clusters of ions are
sometimes referred to as “formula units”
rather than “molecules”.
Covalent compounds have distinct,
strongly bonded molecules. This is why
some people call covalent compounds
“molecular” compounds.
Ionic compounds are solid at room
temperature.
Covalent compounds may be solid, liquid
or gas at room temperature.
Ionic compounds usually have a high
melting point. That’s why they are solid.
Covalent solids usually have a low melting
point.
Ionic solids are usually hard, but brittle
Covalent solids are usually softer
Ionic compounds are usually more soluble Covalent solids are usually less soluble in
in water, but less soluble in non-polar
water, but more likely to dissolve in nonsolvents like acetone.
polar solvents like acetone.
38
Sample Ions
Ion names
alternate names
Sodium
Na
Na+
Sodium ion
Calcium
Ca
Ca2+
Calcium ion
Aluminum
Al
Al3+
Aluminum ion
Tin
Sn
Sn4+,Sn2+
Tin(IV) ion, Tin(II) ion
Stannous, Stannic
Copper
Cu
Cu2+, Cu+
Copper(II) ion, Copper(I) ion
Cuprous, Cupric
Iron
Fe
Fe3+, Fe2+
Iron(III) ion, Iron(II) ions
Ferrous, Ferric
Carbon
C
C2+, C4+, C4-
Carbon(II), Carbon(IV), Carbide
Carbon can form both anions and
cations as well as covalent bonds
Nitrogen
N
N3-
Nitride ion,
Phosphorus
P
P3-
Phosphide ion,
Oxygen
O
O2-
Oxide ion,
Sulphur
S
S2-
Sulphide ion
Fluorine
F
F1-
Fluoride ion
Chlorine
Cl
Cl1-
Chloride ion
Sulfide ion
Metal Ions (+)
ions
Non-Metal Ions (-)
– Anions
+ Cations
Element
Notice that some elements can form more than one type of ion. Compounds of the same element can differ quite a
bit, for example, red iron oxide (rust) has Fe3+ ions, black iron oxide (wustite) contains Fe2+ ions. Note also, that most
negative ions have the name ending changed to –ide.
39
H
H
N
H
+
H
O 2–
O S O
O
Big Fat Ions
Cl O
–
3–
O
O
P
O
O
• Polyatomic ions are ions that are composed of
a cluster of atoms, instead of a single atom.
• For example, the nitrate ion (NO3–) looks like
O
this:
NO O
• But it acts like a single, negatively charged
O
Na
particle in reactions.
Na + NO  NaNO
N O
O
• Polyatomic ions are sometimes called radicals.
• They are not the same as molecules.
+
-
+
3
-
3
Common Polyatomic Ions (see p.422)
3-
2-
Formula
Name (ionic charge)
Formula Name (ionic charge)
PO4 3-
Phosphate ion (3-)
NO3 -
Nitrate (1-)
PO3 3-
Phosphite ion (3-)
NO2-
Nitrite (1-)
SO4 2-
Sulphate ion (2-)
ClO4 -
Perchlorate (1-)
SO3 2-
Sulphite ion (2-)
ClO3 -
Chlorate (1-)
CO3 2-
Carbonate (2-)
ClO2 -
Chlorite (1-)
CrO4 2-
Chromate (2-)
ClO -
Hypochlorite (1-)
2-
Oxalate (2-)
MnO4
SiO3 2-
Silicate (2-)
H2PO4 -
Dihydrogen phosphate (1-)
HPO4 2-
Hydrogen phosphate (2-) .
H2PO3-
Dihydrogen phosphite (1-)
HPO3 2-
Hydrogen phosphite (2-)
HSO4-
Hydrogen sulphate (AKA: bisulphate) (1-)
Cr2O7 2-
Dichromate (2-)
HSO3 -
Hydrogen sulphite (AKA: busulphite) (1-)
C2H3O2 -
Acetate (AKA: ethanoate) (1-)
HCO3 -
Hydrogen carbonate (AKA: bicarbonate) (1-)
OH-
Hydroxide (1-)
NH4 +
Ammonium (1+)
CN -
cyanide (1-)
H3O+
Hydronium (also written as H+) (1+)
C2O4
1-
-
Permanganate (1-)
This information is important when naming ternary ionic compounds. Click to skip ahead to Ionic Naming Rules
1-
41
Review
2
Representation of Atoms
Overview
Since the time of classical Greece, humans have
tried to represent what matter was made of.
Because the particles of matter are too small to see,
we have used models to represent our concepts of
atoms and molecules.
R
Representation of Atoms
0.2.0
• Early Representations
– Democritus (c.450 BCE)
• first suggested that matter was made of particles.
– John Dalton (1800)
• represented the atoms as spheres (like microscopic
bowling balls)
– J.J. Thomson represented the atom as a “plum
pudding” of positive charge with negative charged
electrons scattered inside “like rasins”
-
-
+
-
-
– You studied the historic importance of these models last year,
so you will not be tested on them this year. We will
concentrate on the three most widely used representations on
the slides that follow.
H
C
N
O
P
S
Cl
Dalton
models
Original
and
Modern
-
43
R
0.2.1
Page 5
Page 5
Rutherford-Bohr Model
– Rutherford discovered that the atom has a
dense nucleus containing positively charged
protons.
– Negatively charged electrons move around
this nucleus in paths that resemble an orbit.
– Later, Bohr calculated that there were
different orbital energy levels or “shells” that
could hold different numbers of electrons.
– A pattern of “Bohr numbers” corresponds to the
formula 2n2 where n is a whole number.
• Bohr numbers: 2,8, 18, 32, 50…
44
Early
Rutherford
model
Revised
Bohr
model
Page 5
The Simplified Atomic Model
0.2.2
Page 6
– The simplified atomic model that we often use
today adds neutrons (discovered by James Chadwick
after the Bohr-Rutherford model had been proposed) to
the protons in the nucleus.
– We often draw this in a simplified way, showing
the nucleus as a full circle, and the electron
“shells” as half-circles.
Symbol: The symbol of the element
Na
Nucleus: If asked for a
complete simplified
model, give the #protons
and #neutrons (if known)
in the nucleus. Otherwise,
just draw a full circle.
Electrons: 2 in first shell, 8 in 2nd 1 in 3rd
11p+
12n0
Z=11,
2e-
8e-
1e-
configuration: 2,8,1
The Atomic Number, Z, is the
number of protons in the element.
The configuration is the
arrangement of the electrons in the
shells
45
• Be careful how you draw them!
• The diagram must show the nucleus!
Nucleus is not shown.
Nucleus shown as solid circle.
Labelled with element symbol beside.
ACCEPTABLE
Nucleus is confused with 1st shell
Nucleus shown as full circle.
Labelled with #protons and neutrons.
ACCEPTABLE
Page 6
The Sub-atomic Particles
Particle
nucleons
Symbol
Charge
Actual Mass
(g)
Rounded mass
(amu)
Location in
atom
Proton
p+
1+
1.672x10 -24
≈1 u (1679/1680)
Nucleus
Neutron
n0
0
1.674x10 -24
≈1 u (1680/1680)
Nucleus
Electron
e-
1-
9.109x10 -28
≈0 u
(1/1680)
Shells
47
R
0.2.3
Page 7
Lewis Model: (AKA Lewis electron dot notation)
– Lewis notation is a way of drawing a representation of
the valence electrons of an atom
– When sketching an atom, write the symbol, and then
arrange dots around it to represent its valence
electrons.
2 paired electrons
– Example: N has 5 valence electrons
N 3 “odd” unpaired electrons
– The “odd” or unpaired electrons are available for the
purpose of bonding.
– Because there are 3 electrons available for bonding,
we say nitrogen has a valence of 3.
– When bonding, atoms gain, lose or share electrons in
order to get a total of 8* electrons around each atom.
1 5
4
2
3
48
The preferred way of drawing Lewis diagrams of the first ten elements is shown below:
However, the dots may be moved around to show different arrangements. All of the
drawings of Beryllium shown below might be correct in some circumstances.
Sometimes electrons are removed from one atom to others in order to get 8
Sometimes showing the bonding between atoms requires clever sharing of dots, as in
the drawing of a nitrogen molecule (N2) shown here:
49
The Modern Model
(Optional Enrichment)
• The Modern Model of the Atom
– Of course, the Rutherford-Bohr model and the
Simplified Model do not perfectly represent what
happens inside the atom. No model can!
– A more complete model, The Modern or ElectronCloud model exists, but is more complicated and
extremely difficult to draw.
– The Modern Model more accurately explains the
relationship between the atom and the periodic table,
and allows you to produce simplified models of
elements in the transition area of the periodic table.
50
The Modern Model
(Optional Enrichment)
• The 2-8-8 vs. 2-8-18 problem.
– You have probably been taught how to draw
Simplified Models for the first 20 elements
– If so, you have noticed that for the elements
potassium and calcium, the third shell only holds 8
electrons—but Bohr said it should hold up to 18!
– The models you have been taught can’t explain
why, but the modern model includes a concept
called “orbitals” or subshells, and a filling pattern
called the “aufbau diagram” that explain this .
51
The Modern Model
(Optional Enrichment)
• You are not required to learn the Aufbau
diagram or the modern electron cloud model,
but if time permits, I will show you how it
works near the end of the review section. In
the meantime:
You must know that the third shell CAN hold up to 18
electrons, but often doesn’t.
And you must learn how the periodic table can be
used to figure out the electron arrangement of many
elements past the first 20.
But that is part of the next lesson…
52
Atomic Model Exercises
1. Draw Simplified Models of the first 20 elements.
2. Draw Lewis Models of the first 20 elements.
3. Convert the following:
a) 125 cm to m
b) 280 g to kg
c) 4.63 L to mL
d) 320 mL to L
e) 45000 mm to km
f) 5.52 kg to g
g) 750 mL to L
h) 0.0035km to cm
i) 0.45L to mL
Review
3
Periodic Classification
Overview
The periodic table is a useful arrangement of the
elements, into regions, families and periods that
have important meanings. It is also a source of
much additional information about the elements.
With careful interpretation of the table, we can find
the number of protons an atom has, the
approximate number of neutrons, and the
arrangement of electrons in the atom and in its ions.
Information in your Periodic Table
Atomic number (Z)
8
The number of protons 2-
Ionic charge
Electronegativity
3.44
0.65
Atomic Radius
Ionization Energy
Melting Point (°C)
1314
1.43
Density (g/L gas)
Boiling Point (°C)
Electronegativity is a
rating of how well the
atom attracts
electrons, on a scale
from 0 to 4
Ionization Energy is
how much energy it
takes to remove an
electron (kj/mol)
-218.3
-182.9
O
Oxygen
(g/mL solid/liquid)
Symbol
Name
The English name of the element
15.999
Atomic weight (amu)
(or g/mol)
Also the molar mass in g/mol
The symbol is a 1 or 2 letter abbreviation of the element’s name, or sometimes its Latin or
German name. The first letter is always uppercase. If there is a second letter it MUST be
written in lowercase. (eg. For sodium, Na is correct, na or NA are absolutely unacceptable!)
55
In-line Notation of Element
Information
• An alternative to the periodic table is in-line
notation of elements and isotopes. Note that the
arrangement of information in this notation
system is not the same as the arrangement in
most periodic tables.
• Examples of inline notation:
7
1+
1
+1
𝐿𝑖
𝐿𝑖
𝐿𝑖
𝐿𝑖
𝐿𝑖2 𝑂
3
• In-line notation is designed to be more compact,
but less complete presentation of the information
in a full periodic table.
Inline Isotope, Ion and Molecule Notation
Mass Number
(AKA. Isotope Number)
Represents the
number of nucleons
in THIS particular
atom
Total Nucleons 
Atomic number “Z”
represents the
number of protons
in this atom
Minus Protons 
Valence Number
(used in bonding)
14
4+
+4
4
Ionic Charge
(used for ions)
Oxidation Number
–
6
8 neutrons
(used in electrochemistry)
2 (s)
Phase Marker
(solid, liquid, gas
aqueous)
Number of atoms in
a molecule
Looking at the Examples Again
• Examples of inline notation:
7
1+
1
+1
𝐿𝑖
𝐿𝑖
𝐿𝑖
𝐿𝑖
3
7
3𝐿𝑖
1+
𝐿𝑖
𝐿𝑖1
𝐿𝑖 +1
𝐿𝑖2 𝑂
𝐿𝑖2 𝑂
means an isotope of Lithium with 3 protons and 4 neutrons
means a lithium ion with a charge of 1+
means that the normal valence of lithium is 1 (forms 1 bond)
means that in a particular compound, lithium has oxidation #+1
means a compound that contains 2 atoms of lithium and 1 atom of
oxygen
Protons
Neutrons
Electrons
Charge
Oxidation
or Valence
Avg. mass
7
3
4
3
Charge= 0
6.99
≈7
3
≈4
2
Charge= 1+
6.99
14
6
C
14
6
8
6
Charge = 0
12.0
12C4+
12
6
6
2
Charge= 4+
12.0
C -4
≈ 12
6
≈6
≈10
Oxid. #= -4
12.0
18O2
18
8
10
8
Valence= 2
16.0
Mass #
(total nucleons)
FAQ
• What is the difference between mass number and
atomic mass (AKA atomic weight) ?
– Mass number is the actual number of nucleons in a
particular atom (or isotope). It is mostly used in
nuclear chemistry to distinguish isotopes of the same
element (eg Uranium-238 238U vs. Uranium-235: 235U).
It is always a whole number.
– Atomic mass is the average mass of all the atoms of an
element, as it is recorded in periodic tables. This
average was found by taking the average mass of many
samples using a mass spectrometer. It is normally a
decimal number.
FAQ
• What is the difference between valence, charge and
oxidation number ?
– Valence is the number of bonds an atom is likely to form.
Specifically, it is the number of “odd” electrons in an atom
available for forming bonds.
– Charge is the actual electric charge an atom will get when it
forms an IONIC bond. Some elements have more than one
possible charge. When representing charge, the sign is
written after the number. Eg 2+, 3‒ etc.
– Oxidation number is the charge assigned to an atom inside
a compound, determined by a set of arbitrary rules. This is
used mainly in electrochemistry to determine if and
oxidation/reduction reaction can occur. Oxidation signs are
written before the numbers. Eg. -2, +3 etc.
The Periodic Table
with Primary Regions shaded
1
2
3
4
I
5
6
Solid
1
H
II
2
Li
Be
3
Na
Mg
III
B
IV
B
V
B
4
K
Ca
Sc
Ti
5
Rb
Sr
Y
6
Cs
7
Fr
7
8
9
10 11 12 13 14 15 16 17 18
Gas
Liquid
VIII
Synthetic
metal
Metaloid
Nonmetal
VI
B
VII
B
VIII
B
V
Cr
Mn
Fe
Co
Zr
Nb
Mo
Tc
Ru
Ba
Hf
Ta
W
Re
Ra
Rf
Db
Sg
Bh
III
IV
V
VI
VII
He
B
C
N
O
F
Ne
I
B
II
B
Al
Si
P
S
Cl
Ar
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
Hs
Mt
Ds
Rg
Cn
Uut Uuq Uup Uuh Uus Uuo
↑ The properties and region associations of these 10 elements are hypothetical ↑
6
La
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
7
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
The heavy “staircase” line was the traditional separation between metals & non-metals but we now
know it is not a sharp division.
62
1
H
2
Li
Be
3
Na
Mg
III
B
IV
B
V
B
4
K
Ca
Sc
Ti
V Iron
CrTriad
Mn
5
Rb
Sr
Y
Zr
Nb
Mo
6
Cs
Ba
Hf
Ta
7
Fr
Ra
Rf
Db
II
8
9
VII
B
B
C
N
O
F
Ne
10 11 12 13 14 15 16 17 18
Transition Elements
VI
B
VIIIA: Noble Gases
7
VIIA: Halogens
6
VI: Oxygen Family
5
V: Nitrogen Family
4
IVA: Carbon Family
3
IIIA: Boron Family
I
2
with Families Shaded
IB: Coin Metals
1
IIA: Alkaline Earths
IA: Alkali Metals
The Periodic Table
VIII
B
III
IV
V
VI
VII
VIII
He
I
B
II
B
Al
Si
P
S
Cl
Ar
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
Sg
Bh
Hs
Mt
Ds
Rg
Cn Uut
Uuq Uup Uuh Uus Uuo
↑ The properties and family associations of most elements in period 7 are hypothetical↑
Lanthanides
6
La
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
Actinides
7
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
2nd Transition or “Rare Earth” Elements
63
The Periodic Table
Li
Be
THREE (III)
FOUR (IV)
FIVE (V)
SIX (VI)
SEVEN (VII)
3
Na
Mg
III
B
IV
B
V
B
VI
B
VII
B
4
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
5
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
6
Cs
Ba
Hf
Ta
W
Re
7
Fr
Ra
Rf
Db
Sg
Bh
8
9
10 11 12 13 14 15 16 17 18
III
Transition Elements
IV
V
VI
VII
TWO
7
ONE
6
EIGHT (VIII)
2
5
SEVEN (VII)
II
4
SIX (VI)
H
3
FIVE (V)
TWO (II)
1
2
FOUR (IV)
I
1
THREE (III)
ONE (I)
and Valence Electrons (electrons in outermost shell)
VIII
He
B
C
N
O
F
Ne
I
B
II
B
Al
Si
P
S
Cl
Ar
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
Hs
Mt
Ds
Rg
Cn Uut
VIII
B
Uuq Uup Uuh Uus Uuo
↑ The properties and family associations of these synthetic elements are hypothetical ↑
6
La
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
7
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
If the square is the same colour as the arrow above, it obeys its family with respect to valence. If the square
is rainbow shaded, it is polyvalent, and not obeying its family rules. If the square is partly shaded, then it
obeys its family rules most of the time.
64
The Periodic Table
with Periods shaded
1
2
3
4
5
6
7
8
9
10 11 12 13 14 15 16 17 18
I
VIII
1
H
II
1st Period = 1 shells
III
IV
V
VI
VII
He
2
Li
Be
2nd Period = 2 shells
B
C
N
O
F
Ne
3
Na
Mg
3rd Period = 3 shells
Al
Si
P
S
Cl
Ar
4
K
Ca
Sc
Ti
V
Cr4th Mn
Fe = 4Coshells
Ni
Period
Cu
Zn
Ga
Ge
As
Se
Br
Kr
5
Rb
Sr
Y
Zr
Nb
Mo5th Period
Tc Ru = 5Rh
Pd
shells
Ag
Cd
In
Sn
Sb
Te
I
Xe
6
Cs
Ba
Hf
Ta
W6th Re
Os = 6Irshells
Pt
Period
Au
Hg
Tl
Pb
Bi
Po
At
Rn
7
Fr
Ra
Rf
Db
Sg7th Period
Bh Hs= 7Mt
Ds
shells
Rg
Cn
Uut Uuq Uup Uuh Uus Uuo
↑ The properties and family associations of these 10 elements are hypothetical ↑
6
La
Ce
Pr
Nd
Pm
Sm Eu= 6Gd
Tb
6th Period
shells
Dy
Ho
Er
Tm
Yb
Lu
7
Ac
Th
Pa
U
Np
PuPeriod
Am =
Cm7 shells
Bk Cf
7th
Es
Fm
Md
No
Lr
The periods of the table show how many shells of electrons an element normally has.
65
How to Use the Periodic Table
to Find the Electron Arrangement of an Atom
Eg. Find the electron arrangement of Iodine (I)
H
II
A
2
Li
Be
3
Na
Mg
III
B
IV
B
V
B
VI
B
VII
B
4
K
Ca
Sc
Ti
V
Cr
Mn
5
Rb
Sr
Y
Zr
6
Cs
Ba
Hf
Transition Elements
III
A
IV
A
V
A
VI
A
B
C
N
O
VIII
A
He
F
Ne
Cl
Ar
SEVEN (VII)
1
I
B
II
B
Al
Si
P
S
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Nb
Mo 5thTcPeriod
Ru =Rh
Pd Ag
5 shells
Cd
In
Sn
Sb
Te 53I
Xe
Ta
W
Hg
Tl
Pb
Bi
Po
Rn
Re
VIII
B
Fe
Os
Co
Ir
Ni
Pt
Au
At
Iodine is at the intersection of Period 5 and Family VII. Its number is 53. It has a
total of five shells, 7 electrons in the outermost shell, and will have 53p+, and
normally 53 e-. From this we can USUALLY figure out the electron arrangement.
Five shells
Note:
The inner shells usually
contain Bohr numbers
53p+
2 8 18 18
7e-
Total 53, So far: 35, left: 18
Periodic Table Exercises
• Write the name and symbol of each of the first
20 elements. (bragging rights if you can do it
without looking!)
Review
4
The Naming of Compounds
Overview
In this section you must learn to name compounds
based on their formulas, and also to find their
formulas based on their names. This is a bit trickier
than you might think, since different naming
methods are used for different types of compounds.
You will need to distinguish between ionic, covalent
and organic compounds, and also learn some
common polyatomic ions.
Some elements with FORMULAS as well as SYMBOLS
Element
(systematic name)
Formula
Common name
Other names
dioxygen
O2
Oxygen gas
Diatomic oxygen
Trioxygen
O3
Ozone
Triatomic oxygen
dihydrogen
H2
Hydrogen gas
Diatomic hydrogen
dinitrogen
N2
Nitrogen gas
Diatomic nitrogen
dibromine
Br2
Bromine
Diatomic bromine
diiodine
I2
Iodine
Diatomic iodine
dichlorine
Cl2
Chlorine gas
Diatomic chlorine
difluorine
F2
Fluorine gas
Diatomic fluorine
tetraphosporus
P4
White phosphorus
Yellow phosphorus
amorphous phosphorus
Pn
Red phosphorus
hexasulphur
S6
Orange sulphur
hexathiane
octasulphur
S8
Yellow sulphur
brimstone
carbon (diamond)
CcF8 or C8n
diamond
amorphous carbon
Cn or C
Soot, charcoal
carbon (graphite)
CmP6 or C6n Graphite
Pencil “lead”
buckminster fullerine
C60
“Buckyballs”
Fullerine C-60
Carbon black
Naming Compounds
• There are four sets of rules for naming
compounds:
– The binary ionic rules:
• For compounds containing only two elements, joined by an
ionic bond (usually a metal and non-metal).
– Ternary ionic rules:
• For compounds containing 3 or more elements, including a
polyatomic ion.
– The covalent rules:
• For two elements joined by covalent bonds (usually two
non-metals)
– The organic rules:
• Used for compounds that contain carbon atoms bonded to
each other covalently.
70
Compound Types
Ionic
Binary
Ionic
Covalent
Ternary
Ionic
Exceptions
H2O, H2O2, CH4, NH3
Covalent/Organic
Organic
Is it a common exception?
Does it end with COOH?
Does it end with CH2OH?
Does it end with OH?
Does it end with a radical?
Does it start with H?
Does it begin with a metal?
Does it begin with NH4?
Does it contain >1 carbons?
Did you answer NO to all?
See “exceptions”
Organic (acid)
Organic (alcohol)
Ternary Ionic (base)
Ternary Ionic
Binary Ionic (acid)
Binary Ionic (salt)
Ternary Ionic (ammonium)
Organic
Covalent
The Exceptions
• H2O is usually called WATER (common name)
• Other acceptable names are dihydrogen monoxide, dihydrogen
oxide, hydrogen monoxide, hydrogen hydroxide and oxidane.
• H2O2 is usually called hydrogen peroxide (ionic rule)
• Other names are dihydrogen dioxide, dioxidane, and oxidanyl.
• CH4 is usually called methane (organic rule)
• Pronounced either Meth-ane (Amer.) or me-thane (Brit.)
• Other names include: carbon tetrahydride, tetrahydriocarbon.
• NH3 is usually called ammonia (organic functional rule)
• Other names can include azane, hydrogen nitride, trihydrogen
nitride and nitrogen trihydride,
This list is by no means exclusive. Many other exceptions exist, but they are unlikely to
be encountered in a high school chemistry course.
The Binary Ionic Rules
– First name the element on the left side of the
compound’s formula.
– Then name the element on the right hand side of
the compound’s formula, but change the suffix
to “ide”
1+
• For example:
1–
Na+ ClCa2+ O2-
O23+
Al O2- Al 3+
O2-
NaCl  sodium chloride
CaO  calcium oxide
Al2O3 aluminum oxide
BaCl2  barium chloride
K2S potassium sulphide
Ca2C  calcium carbide
Cl- Ba2+ ClK+ S2- K+
Ca2+ C4- Ca2+
We don’t need prefixes because ionic
compounds ALWAYS follow the crossover rule
73
Common Non-metal Ion Names
Element (symbol)
Negative Ion (charge)
Element (formula)
Negative Ion (charge)
Boron (B)
Boride (B5-)
Phosphorus (P4)
Phosphide (P3-)
Carbon (C)
Carbide (C4-)
Sulphur (S8)
Sulphide (S2-)
Silicon (Si)
Silicide (Si4-)
Fluorine (F2)
Fluoride (F–)
Arsenic (As)
Arsenide (As3-)
Hydrogen (H2)
Hydride (H–)
Selenium (Se)
Selenide (Se2-)
Chlorine (Cl2)
Chloride (Cl–)
Bromine (Br2)
Bromide (Br–)
Iodine (I2)
Iodide (I–)
Nitrogen (N2)
Nitride (N3-)
Oxygen (O2)
Oxide (O2-)
Other monatomic negative ions occur
rarely. If you encounter one, use the
atomic name, with the last syllable
altered to ide as sounds best. Eg.
Antinide or Polonide
74
Ionic Rules No No!
• When naming an ionic compound (and that
includes most compounds that contain a metal)
YOU SHOULD NOT USE A PREFIX!
• Do NOT say: calcium difluoride for CaF2
• It’s Wrong. The correct name is just calcium fluoride.
• Do NOT say: dialuminum trioxide for Al2O3
• It’s Wrong. The correct name is aluminum oxide.
There are, or rather there USED to be, a few exceptions to this. Chromium dioxide
was an accepted name for CrO2, and is still used occasionally. Now the name
chromium(IV)oxide is preferred for the compound, since it obeys the ionic rules.
Ionic Rules:
Dealing with
This
copper
ion has
a
charge
of 1+
• Some metal elements have more than one
possible valence. Copper, for example, can have
a charge of 1+ or 2+, depending on which
compound it is in (eg. CuCl or CuCl2). Since we
don’t use prefixes in naming ionic compounds,
we shouldn’t use copper dichloride. We need a
new rule!
• If a metal is polyvalent, we include its current
valence in roman numerals inside parenthesis
within an ionic compound name, for example:
This
copper
ion must
have a
charge of
2+
– CuCl = Copper (I) chloride (not copper monochloride)
– CuCl2 = Copper (II) chloride (not copper dichloride!)
76
Polyvalent Elements
The elements with flashing circles have more than one positive valence.
1+ 2+ 3+ 4+ 5+ 6+ 7+
4-
3-
2-
1-
I
0
VIII
1
H
II
III
IV
V
VI
VII
He
2
Li
Be
B
CC
N
O
F
Ne
3
Na
Mg
III
B
IV
B
V
B
VI
B
VII
B
II
B
Al
Si
P
S
Cl
Ar
4
K
Ca
Sc
Ti
Ti
V
Cr
Cr
Mn
Mn
Co Ni Cu
Cu
Fe Co
Zn
Ga
Ge
As
Se
Br
Kr
5
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
R
u
Rh
Pd
Pd
Cd
In
Sn Sb
Sb
Sn
Te
I
Xe
6
Cs
Ba
Hf
Ta
W
Re
Os
Ir
Pt Au
Au
Pt
Hg
Hg
Tl Pb
Pb Bi
Bi Po
Po
Tl
At
Rn
7
Fr
Ra
U
n
c
e
r
Nonmetal
VIII
B
I
B
t
Ag
a
i
n
6
La
Ce
Pr
Nd
Pm
Sm
Sm
Eu
Eu Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
7
Ac
Th
Pa
UU
Np
Pu
Am
Bk
Cf
Es
Fm
Md
No
Lr
Cm
77
The Common Polyvalent Ions
Formula
Charge
Stock Name (new name)
Classical Name (old name)
Cu+
1+
Copper (I) ion
Cuprous ion
Cu2+
2+
Copper (II) ion
Cupric ion
Fe2+
2+
Iron (II) ion
Ferrous ion
Fe3+
3+
Iron (III) ion
Ferric ion
Sn2+
2+
Tin (II) ion
Stannous ion
Sn4+
4+
Tin (IV) ion
Stannic ion
Pb2+
2+
Lead (II) ion
Plumbous ion
Pb4+
4+
Lead (IV) ion
Plumbic ion
Mn2+
2+
Manganese (II) ion
Manganous ion
Mn3+
3+
Manganese (III) ion
Manganic ion
Cr2+
2+
Chromium (II) ion
Chromous ion
Cr3+
3+
Chromium (III) ion
Chromic ion
Hg+, Hg22+
1+
Mercury (I) ion
Mercurous ion
Hg2+
2+
Mercury (II) ion
Mercuric ion
78
Examples of Ionic Compounds with
Polyvalent Elements
Formula
Common
Name
Stock (new) Name
Classical (old) Name or
Incorrect name
ions
FeO
Wustite
Iron(II)oxide
Ferrous oxide
Fe2+, O2-
Fe2O3
rust
Iron(III)oxide
Ferric oxide
Fe3+, O2-
Iron(II,III)oxide
Ferosso Ferric oxide*
Fe2+, Fe3+, O2-
Fe3O4
Cu2O
Cuprite (red)
Copper(I)oxide
Cuprous oxide
Cu+, O2-
CuO
(Black green)
Copper(II)oxide
Cupric oxide
Cu2+, O2-
CrO
Chrome black
Chromium(II)oxide
Chromous oxide
Cr2+, O2-
Cr2O3
Chrome green
Chromium(III)oxide
Chromic oxide
Cr3+, O2-
CrO2
Crolyn
Chromium(IV)oxide
Chromium dioxide
Cr4+, O2-
CrO3
Chromic acid
Chromium(VI)oxide
Chromium trioxide
Cr6+, O2-
PbCl2
cotunnite
Lead(II)chloride
Plumbous chloride
Pb2+, Cl-
PbO2
platternite
Lead(IV)oxide
Plumbic oxide
Pb4+, O2-
*Ferrosso ferric oxide is a unique combination of Iron(II)oxide and Iron(III)oxide together in a crystalline ionic structure
Its formula can also be given as (FeO∙Fe2O3)
79
The Ternary Ionic Rules
– First name the metallic element (or ammonium
ion) on the left of the formula.
– Then name the polyatomic ion on the right side of
the formula.
Polyatomic ions: See Table 8.10 on p. 422 or click here
• If the compound is an ammonium salt, then name the
non-metal ion, changing it to end in “ide”
• Examples:
– NaNO3sodium nitrate
– K2SO4potassium sulphate
– Al2(CrO4)3aluminum chromate
CaCO3calcium carbonate
Ba(CN)2barium cyanide
NH4Cl  ammonium chloride
80
Covalent Rules
– Name the less electronegative element on the left.
– Name the more electronegative element on the
right, changing its suffix to “ide”
– Add prefixes to each element to indicate the
number of atoms in the formula:
• Mono*=1, di=2, tri=3, tetra*=4, penta*=5, hexa*=6
• Examples:
– CCl4 carbon tetrachloride**
– PF3  phosphorus trifluoride**
– CO2  carbon dioxide**
N2H4  dinitrogen tetrahydride
P2O5 diphosphorus pentoxide
CO  carbon monoxide**
* The last “o” in mono or the “a” in tetra, penta, or hexa is usually dropped before “oxide” to sound better. (eg.
“Carbon monoxide”, not “carbon monooxide”)
** The “mono” prefix is usually dropped from the first element of the compound, except when that would
cause confusion between two similar compounds.
81
Simplified System
Bonds
(how much an atom attracts electrons)
I
I
1 2.2
Solid
Solid
Another use of
electronegativity is
to find how ionic or
covalent a bond is.
Click the arrow
above to skip to the
section on bonds
V
VII
IVII
I
0.0
H
He
IV
VI
VII
II
III
2.2
0.0
1
1.0
1.0
1.6
0.7
1.5
2.0
2.5
3.0
3.2
3.6
2.0
2.6
3.0
3.4
4.0
0.0
2
H
He
Li
Be
0.9 1.4 1.9 2.4 2.9 3.1 3.5 4.0
B
C
N
O
F
Ne
2.0 2.6 3.0 3.4 4.0 0.0
2 1.0 1.6
IV
V
VI
VII
VIII
I
II
1.6
3 0.9Li 1.3Be IIIB
B 1.9C 2.2N 2.6O 3.2F 0.0Ne
B
B
B
B
B
B
B
Na Mg
Al
Si
P
S
Cl
Ar
IV
V
VI
VII
VIII
I
II
1.6
1.9
2.2
2.6
3.2
0.0
3 0.9 1.3 IIIB
4 0.8Na 1.0Mg 1.4 1.5B 1.6B 1.7B 1.6B 1.8 1.9B 1.9 1.9B 1.7B 1.8Al 2.0Si 2.2P 2.6S 3.0Cl 0.0Ar
K
Ca Sc
Ti
V
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
0.8
1.0
1.4
1.5
1.6
1.7 1.6 1.8 1.9 1.9 1.9 1.7 1.8 2.0 2.2 2.6 3.0 0.0
4
0.8
0.9
1.2
1.3
1.6
2.2
5
K
Ca Sc
Ti
V
Cr 2.1
Mn 2.2Fe 2.3Co 2.2Ni 1.9Cu 1.7Zn 1.8Ga 2.0Ge 2.0As 2.1Se 2.7Br 0.0Kr
Rb Sr
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I
Xe
0.8
0.9
1.2
1.3
1.6
2.2
2.1
2.2
2.3
2.2
1.9
1.7
1.8
2.0
2.0
2.1
2.7
0.0
5
0.8
0.9
1.3
1.5
1.7
1.9
2.2
2.2
2.2
2.4
1.9
1.8
1.8
1.9
2.0
2.2
0.0
6
Rb Sr
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I
Xe
Cs Ba
Hf Ta W Re Os Ir
Pt Au Hg Tl Pb Bi Po At Rn
1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2 0.0
6 0.8 0.9
0.7
0.9
Rf
7
Cs Ba
Hf DbTa SgW BhRe HsOs MtIr DsPt RgAu CnHg UutTl UuqPb UupBi UuhPo UusAt UuoRn
II
III
IV
V
VI
VII
Fr Ra
0.7 0.9
Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo
There are a few
like CH4 and NH3, where the more electronegative elements are written first. These formulas
Fr exceptions,
Ra
have been used for years, and are based on organic chemistry concepts, so it’s unlikely we will change them.
7
DO
use prefixes with covalent compounds
# atoms
Covalent prefix
Examples
1
Mono…, mon…
carbon monoxide (CO), mononitrogen monoxide(NO)
2
Di…
carbon dioxide (CO2), dihydrogen dioxide* (H2O2)
3
Tri…
nitrogen trichloride (NCl3)
4
Tetra…, tetr…
carbon tetrachloride (CCl4), tetramethyl lead ((CH3)4Pb)
5
Penta…, pent…
diphophorus pentoxide (P2O5), nitrogen pentafluoride (NF5)
6
Hexa…, hex…
sulphur hexafluoride (SF6)
7
Hepta…, hept…
bromine heptafluoride (BrF7), heptose (C7H14O7)
8
Octo…, oct…
diphosporus octafluoride (P2F8) , octane (C8H18)
9
Nona…, non…
nonane (C9H20)
10
Deca…, dec…
Decane (C10H22)
*commonly called hydrogen peroxide.
Simplification of Covalent Names
• IUPAC (The International Union of Physicists and Chemists)
which oversees naming conventions, allows some
simplifications to the systematic names of covalent
compounds.
– The “mono” prefix may be dropped from an element, unless doing
so could result in confusion.
• We usually say “carbon dioxide” rather than “monocarbon dioxide”
• However, we always say “carbon monoxide” for CO, since there are two
common oxides of carbon (CO2 and CO)
– A prefix may be dropped from a formula if there is no ambiguity in
the formula, and if the formula obeys the crossover rule.
• Many chemists simply say “hydrogen sulphide” instead of “dihydrogen
sulphide” for the compound H2S. Since H2S is the only common sulphide
of hydrogen and obeys the crossover rule, this doesn’t cause confusion.
– Knowing when simplification is allowed is a matter of experience.
Until you become familiar with the conventions, it is safer to use
all the prefixes. It’s not wrong to include them all (in covalent compounds).
• Water can be called “dihydrogen monoxide”, but it is acceptable to use
“hydrogen oxide” since water obeys crossover.
84
Finding Formulas
from Compound Names
• For covalent compounds, the name usually tells you
the formula:
• For example: dinitrogen pentoxide = N2O5
• However:
• If the name has been simplified by dropping a prefix you may have
to use the crossover rule, discussed later.
• For example: “sulphur fluoride” has had a prefix dropped, so
S(valence=2) F(valence=1) crossover SF2
• “Sulphur fluoride” is the short name for the compound more accurately called sulphur difluoride.
• For ionic compounds, the name never tells you the
formula.
• You always use the crossover rule to find the formula.
• Example: Sodium oxide is Na1 and O2crossoverNa2O
85
The Crossover Rule
and simple ionic compounds
• The crossover rule is used to find the formula
of a compound when the name has no
prefixes (ie. all ionic compounds and some covalent
compounds that have had a prefix removed)
•
•
•
•
•
•
Example 1: What is the formula of aluminum sulphide?
Aluminum sulphide :
Al
S
Ions:
Al3+
S2Valences (remove signs):
Al3
S2
Cross over:
Al2S3
The formula of aluminum sulphide is Al2S3
86
The Crossover Rule
and covalent compounds
• The crossover rule can also be used for covalent
compounds if prefixes have been dropped from a
name. When a covalent compound’s name has
no prefixes at all, check it with the crossover rule.
•
•
•
•
•
•
Notes:
Example 1: What is the formula of “sulphur chloride”?
Sulphur chloride:
S
Cl
Oxidation numbers:
S2ClValences (remove signs):
S2
Cl1
Cross over:
S1Cl2 or SCl2
The formula of “sulphur chloride” is Al2S3
1) The compound “sulphur chloride” should properly be called sulphur dichloride
2) The prefixes trump the crossover rule. If any prefixes were used in the name, then
they take precedence over whatever formula the crossover rule would give you. 87
The Crossover Rule
simplifying ionic compounds
• Ionic compounds can often be simplified
• Example 1: What is the formula of the compound
made from Barium ions (Ba2+) and Carbide ions (C4-)?
• Ions:
Ba2+ C4• Remove the signs
Ba2
C4
• Cross over:
Ba4C2
• Cancel (divide both by 2) Ba2C
• The formula of barium carbide is Ba2C
Note: Do not simplify covalent compounds by cancellation. Covalent compound
formulas must reflect the compound names that include prefixes.
88
Reverse Crossover Rule
for finding the valence of uncertain ions
• Sometimes we can use the crossover rule in reverse
to find the valence or ionic charge of an ion we are
not certain of, such as an ion of polyvalent metal.
• For example, what is the name of Fe2O3? of FeO?
– They are both iron oxide, but which iron oxide (there are
several types!)
– Fe2 O 3
Fe has a valence of 3, so the
name of the compound is:
Iron(III)oxide
Fe12O12
There’s a problem here!
Oxygen hardly ever has a
valence of 1. Let’s double
both valences.
Fe’s proper valence here is 2
Iron(II)oxide
 The Organic Rules
(not studied this year)
A system of names for organic compound exists that is based on the
number of carbon atoms they have (as a prefix), and the type of
compound they are (as a suffix): alkane (…ane), alkene (…ene) alcohol
(…ol), aldehyde (…hyde), ketone (…one), organic acids, etc.
# carbons
Prefixes
examples
1
Methyl, Formyl
Methane, methanol, formaldehyde, formic acid
2
Ethyl, Acetyl
Ethane, ethanol, acetaldehyde, acetone, acetic acid
3
Propyl,
Propane, propanol, propanoic acid
4
Butyl
Butane, butanol, butanoic acid
5
Pentyl
Pentane, pentanol, pentanoic acid
After this the prefixes resemble those for inorganic compounds, 6=hex, 7=hept, 8=oct, etc.
As you may notice, the common names of some chemicals come from the organic system,
such as methane, the common name of carbon tetrahydride (CH4) . For more information
on organic nomenclature, see the wikipedia article.
Practice
• Page 12, Question #9
• Practice sheets:
• Naming ionic compounds
• Naming covalent compounds
• Naming mixed compounds
91
Review
5
Enumeration of Matter
Overview
In this section we will review the concept of the
mole and the mole formula.
We will also see how the mole concept applies to
molecules and to diatomic and polyatomic
elements.
This is a required concept for stoichiometry, which
we will cover later.
The Mole Concept
0.5.1
and the Enumeration of Matter
• The Mole: The mole is a unit used to count
atoms, ions, molecules, and other
fundamental particles.
• A mole corresponds to Avogadro’s Number of
particles: 6.02 x 1023 particles.
NA
=6.02 x 1023
= 602 000 000 000 000 000 000 000
= six hundred and two sextillion
= the particles in a mole.
93
Molar Mass
0.5.2
• Molar mass is the mass of one mole of atoms
or molecules.
• The symbol for molar mass is M
(not MM!)
• For elements, molar mass corresponds to the atomic
mass found in the periodic table, but expressed in
grams/mol rather than amu. For example, the molar
mass of carbon, M(C )= 12.011 g/mol, (frequently rounded
to 12.0 g/mol)
• For compounds, M is the sum of the masses of all the
atoms in the molecule or all the ions in the formula. For
example, the molar mass of carbon dioxide molecules is:
M(CO ) =44.009 g/mol,
2
(frequently rounded to 44.0 g/mol)
• that is: 2M(C)+2M(O) or 12.001 +2(15.999) g/mol
94
Alternate Method of Calculating Molar Mass
(very useful if you don’t have a scientific calculator)
• Find the molar mass of Na2 CO
CO3:
Na = 2 atoms x 23.0 amu=
46.0
C
= 1 atom x 12.0 amu =
12.0
O
= 3 atoms x 16.0 amu = + 48.0
106.0 g/mol
Multiply the number of
each atom (as given by
the formula) , by the
atomic mass of the atom
(as given in periodic table)
Then add the total
masses together
Change the unit from
amu  g/mol
Diatomic and Polyatomic Elements
• Diatomic elements: There are seven elements
whose molecules normally contain two atoms:
I2, H2, N2, Br2, O2, Cl2 and F2.
• If finding the molar mass of these elements, remember
to double the mass of one atom.
• M (I ) = 253.808 g/mol (not 126.904 g/mol!)
2
• Polyatomic elements: a few elements, such as
sulphur, phosphorus, and sometimes carbon
occur in larger molecules (eg. S8 or P4)
• If a formula like this has been used in a balanced
equation, remember to multiply the atomic mass by
the appropriate amount (eg. M(S8)=256.52 g/mol)
How to Remember the Diatomic Elements: I Have No Bright Or Clever Friends
96
The Mole Formula
The mole formula is used to convert from grams to moles and vice-versa
𝑚
𝑛=
𝑀
m
Actual mass
(g)
𝑎𝑐𝑡𝑢𝑎𝑙 𝑚𝑎𝑠𝑠
# moles = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠
𝑚
𝑀=
𝑛
Molar mass =
n
# moles
(mol)
𝑎𝑐𝑡𝑢𝑎𝑙 𝑚𝑎𝑠𝑠
# 𝑚𝑜𝑙𝑒𝑠
M
Molar mass
(g/mol)
𝑚 = 𝑛𝑀
Actual mass = # moles x molar mass
97
Practice
• Page 14, #12, 13, 14
• Practice sheet:
• Moles and Molar mass
98
Review
6
Physical Changes
Overview
We will look at types of physical change that are
important in chemistry, including
Change of form (deformation)
Change of phase (change of state)
Change of mixture (dissolution)
Physical Changes
• A physical change occurs when a substance
undergoes a modification in its appearance or
form, but does not alter its nature or
characteristic properties.
• In a physical change the molecules or ionic
formula of the substance do not change.
• There are 3 main categories of physical change
• Change of form, caused by crushing, cutting, grinding,
bending, denting, etc.
• Change of phase or state, caused by melting, boiling,
freezing, evaporation, condensation, sublimation, etc.
• Change of mixture, caused by dissolving (dissolution without
reaction), blending, stirring together dry ingredients, mixing
paints, etc.
100
Change of Form
(AKA. Deformation)
• Only one type of deformation is of much
importance in chemistry, and that is when we
grind up reactants into fine powder so that
they will react more quickly.
• We will go into this in greater detail later in the course,
when we study the rates of reaction
• For now, let’s just say that ground up materials usually
react faster.
• Remember, a powdered solid is still a solid
0.6.1
Phase Change
(AKA. Change of State)
• As a pure substance is heated, its particles move
faster. It changes from a solid state to a liquid
state and then to a gaseous state. Your textbook
refers to this as “phase change”
• Change of “phase” is a physical change, since the
particles of the pure substance do not (usually)
change.
Picky note: What your textbook calls “phase change” should more properly be called
“change of state”. Although “phase” and “state” are frequently used as synonyms, the
word phase has a broader meaning in chemistry. There are three main states of pure
matter (solid, liquid, and gas) , “phase” includes these three, but may also apply to many
other possible phases of matter– including aqueous (a solid dissolved in water), gel (a
jelly-like colloidal mixture) etc. In addition, phase can refer to a boundary between two
similar phases that don’t mix, for example, a liquid mixture could have an oily phase and
a watery phase that contact each other but do not mix.
102
Sublimation occurs
when a material
“evaporates” from a
solid straight to a gas,
like dry ice or iodine.
Exothermic
Process
Gas
Terminology associated with
Rapid vaporization is
called “boiling”,
Slow vaporization is
“evaporation”
Endothermic
Process
Change of Phase
or State
Melting (fusion)
Solid
Freezing (solidification)
Liquid
Liquid
103
Comparison of the States of Matter
Solid
Liquid
Gas
Shape
Definite
Variable
Variable
Volume
Definite
Definite
Variable
Compressibility
Incompressible
Incompressible
Compressible
Fluidity
Not Fluid
Fluid (flows)
Fluid (flows)
Particle separation
Close together
Close together
Far apart
Motion of particles1
VIBRATION only
ROTATION2 and
vibration
TRANSLATION,
Rotation, and
vibration
Notes:
1. The most characteristic (ie.abundant) motion of the phase is underlined.
2. Rotation implies some tumbling motion, so molecules can move around a little.
104
Phase Markers
• During the course of the year, you will often
notice small letters in parenthesis added formulas
in equation. These “phase markers” are inserted
whenever it is important to know what state or
phase the reactants or products are.
• The most important phase markers are:
•
•
•
•
(s) = solid: the substance is a solid or a powder
(l) = liquid: the substance is a pure liquid
(g) = gaseous: the substance is a gas
(aq) = aqueous: the substance is dissolved in water
Eg:
NaCl(s) H2O(l) NH3(g) NaCl(aq)
105
0.6.2
Dissolution and Solubility
• In dissolution, one or more solutes are mixed
into a solvent to create a solution.
• During dissolution:
• The mass of the substances does not change.
• The total volume is usually slightly less than the sum of
the volumes of the components (since some particles
pass into the spaces between other particles)
• When the solvent cannot dissolve any more of the
solute, the solution is saturated.
106
Dissolving = Physical Change
• Remember that dissolution is normally
considered a physical change, not a chemical one.
The material mixes with the solvent, but is not
significantly altered by it
• In a few cases a material will react with the solvent, rather than
just dissolve. For example, trying to dissolve sodium in water, or
baking soda in vinegar will produce a reaction. In this case a
chemical change has occurred as well.
eg: Na(s) + H2O(l)  NaOH(aq) + H2(g)
• Ionic compounds may “dissociate” while dissolving, that is, their
ions may separate by some distance. While this may seem like a
chemical change, it is not a permanent condition, and is
considered to be a physical change.
eg: NaCl(aq)  Na+(aq) + Cl-(aq) (dissociation of salt)
107
Dissolution of Ammonia Gas in Water
(an extreme case of solubility at 25°C)
• 100g of water + 50g of ammonia  150g of ammonia solution
100
g
+
50g of NH3(g)

150
g
• 100 mL of water + 72058 mL of ammonia  101 mL of NH3(aq) solution
100
mL
+
72.058 litres NH3(g)

101
mL
• If you try to dissolve more than 50g of ammonia in 100 mL of water,
you won’t be able to. There will be leftover ammonia!
100
g
+
55 g of NH3

150
g
+
5g
Ammonia is a great example, because water can absorb what seems like a huge amount of ammonia gas before it
becomes saturated. Mass-wise, its actually half the weight of the water, but volume-wise its over 720 times greater!
108
• Solubility indicates the maximum amount of
solute that can dissolve in a given volume of
solvent at a given temperature.
• Solubility is usually expressed as grams of solute per
100 mL of solvent (g/100mL).
• A substance’s solubility can vary with
temperature:
• Solubility of solids usually increases with temperature
• Solubility of gases usually decreases with temperature
• Solubility of gases can also be affected by pressure.
109
Solubility Curves
(Graphs of Solubility vs. Temperature. See page 16)
• Notice how most of the
solids become more soluble
at higher temperatures
– KNO3, for example, starts at a
mere 10 g/100 mL at 0°C, but
goes right off the top of the chart
by 70°C
• Notice that most of the
gases become less soluble
at high temperatures
– NH3 goes from 90 g/100mL at 0°C
to less than 10 g/100 mL at 100°C
110
0.6.3
Concentration and Dilution
• Concentration is the ratio of dissolved solute to total
amount of solution.
• General formula for concentration is:
• But concentration can be expressed in many different
units, including:
• g/L (grams per Litre)
• % (by volume)
• ppm (parts per million)
g/mL (grams per millilitre)
% (by mass)
mol/L (molar concentration)
• Molar concentration is the most important.
111
Molar Concentration
(molarity)
• The molar concentration is the number of moles
of solute that is dissolved in one mole of the
solution.
• Molar concentration can be represented by the
letter C, or by square brackets [] or occasionally
by a capital M used as a unit (molarity). Any of
the following notations could represent a 2.0
mol/L solution of hydrochloric acid:
CHCl = 2.0 mol/L
[HCl] = 2.0 mol/L
CHCl = 2.0 M
The correct unit for molar concentration
is mol/L, although this is sometimes
abbreviated with a capital M for molarity
112
R
Molar Concentration Formula
Molar Concentration =
𝒏𝒖𝒎𝒃𝒆𝒓 𝒐𝒇 𝒎𝒐𝒍𝒆𝒔
𝑽𝒐𝒍𝒖𝒎𝒆 𝒊𝒏 𝒍𝒊𝒕𝒓𝒆𝒔
n
C=
𝒏
𝑽
# moles
(mol)
V=
𝒏
𝑪
V
C
Concentration
(mol/L)
Volume
(L)
n = CV
113
R
Dilution
• Dilution is a physical change that lowers the
concentration of a solution by adding more
solvent.
• The dilution formula is:
C1V1 = C2V2
Where: C1 is the concentration before dilution,
V1 is the volume before dilution
C2 is the concentration after dilution
V2 is the volume after dilution
114
Assignments
• Read pages 15 to 18
• Do page 16
• Questions 15 to 17
• Do page 19:
• Questions 18 to 23
115
Electrolytes
Overview
Our bodies are full of electrolytes: salts, acids and
bases dissolved in our blood and tissues that are
very important for the function of muscles and
nerves.
Most people think that water conducts electricity,
but in fact, pure water is a poor conductor. It is the
presence of dissolved electrolytes that gives the
water in lakes, rivers and us the ability to conduct
electricity.
R
0.6.4
Electrolytes
• Electrolytes are substances which, when
dissolved in water, allow the solution to
conduct electricity.
• Electrolytes are usually ionic compounds.
• Electrolytes “dissociate” into positive and
negative ions when they dissolve.
• There are three main types of electrolytes:
Acids, Bases, and Salts.
• Most solid electrolytes do not conduct
electricity until they are dissolved.
117
Electrolyte Characteristics
Acids
Bases
Salts
Ions:
Release H+ ions
Release OH-ions
Metal and non-metal ions
neutralization
Neutralize bases
Neutralize acids
Products of neutralization
pH
pH is less than 7
pH is greater than 7
pH variable, close to 7*
Litmus
Turn litmus paper red Turn litmus paper blue
Don’t change litmus*
Phenolphthalein
Stays clear
Turns red/purple
Stays clear*
Formula
H + non-metal
Metal + OH
Metal + non-metal
Dissociation eg:
HCl(g)  H+(aq) + Cl-(aq)
NaOH(s)Na+(aq)+OH-(aq) NaCl(s)Na+(aq) + Cl-(aq)
pH Scale
The pH (positive Hydrogen potential) scale is used to measure the relative acidity or
alkalinity of a solution. It is in theory open-ended, but in practice runs from 0 to 14.
Strong Acids
0
1
2
Weak Acids
3
4
5
Neutral
6
7
Weak Base
8
9
10
Strong Base
11
12
13
* Some salts are slightly acidic (aluminum salts) or slightly basic (carbonates)
14
118
Assignments
• Read page 19
• Do page 20
• Questions 24-27
• Question 28
119
Review
7
Chemical Change & Stoichiometry
Overview
Chemical changes are the root of chemistry, and the
chemical equation is the fundamental tool we use to
understand chemical changes.
Once you understand the chemical equation, you
can use the techniques of stoichiometry to find the
amounts of materials needed or used in chemical
changes.
R
0.7 Chemical Changes
• Chemical changes occur when substances
(reactants) react to form new substances
(products).
• The products differ from the reactants:
• They have different characteristic properties.
• They have different molecular or ionic arrangements.
Reactants on the
Left side of
equation
Reactants  Products
Products on the
Right side of
equation
121
R
• Indications that a chemical change has taken
place include:
•
•
•
•
Release of a gas (effervescence)
Significant change in colour
Formation of a precipitate (solid from two solutions)
Change of energy in the form of heat, light or explosion.
• Parts of a chemical equation:
Chemical equation
Reactants
Change to
4 Fe (s) + 3 O2 (g) 
Coefficients
4:3:2
(used for balancing)
Indexes* 2,2,3
Number of atoms in
the molecules
Product
2 Fe2O3 (s)
Phases
(s) Solid (l) liquid (g) gas
(aq) dissolved in water
*Yes, I am fully aware that the dictionary says that the correct plural of index is indices, but for clarity I am using the term the text uses.
122
0.7.1
Conservation of Mass
• During a chemical reaction, mass is neither
lost nor gained
• The total mass of all the reactants is equal to
the total mass of the products.
• This is because no atoms are created or destroyed
during the reactions. The atoms are just rearranged.
H
HH
H
H
2H2
OO
+ O2

2H2O
• The balancing of chemical equations is based on the
law of conservation of mass.
m reactants = m products
123
0.7.2
p. 22
Balancing Equations
• Balancing means adding coefficients in front of
the formulas of an equation so that it will
conform to the law of conservation of mass
• A word equation names the reactants and products
• A skeleton equation is an unbalanced equation
• A balanced equation respects conservation of mass.
• Rules for balancing equations:
• Only coefficients may be added or changed. The indexes in
formulas must not be changed.
• You do not need to write the coefficient 1. It is understood.
• Balanced equations should be reduced to the lowest terms.
• When an equation is properly balanced, the total number of
atoms of each element on the left and right sides will be
equal.
124
R
0.7.3
p. 23
Stoichiometry
• Stoichiometry is the study of the relationships
between the amounts of substances
(reactants and products) that take part in a
chemical reaction.
• Stoichiometry can be used to:
• Calculate the amount of reactants need for a reaction
• Calculate the expected amount of product from a
reaction.
125
Steps for Stoichiometry
1. Balance the equation, or verify that the equation you
have been given is properly balanced.
2. Use the coefficients to find the mole ratios
3. Write the amount in moles of the known reactant
under the corresponding mole ratio number.
•
If the amount is given in grams, convert it to moles using
the mole formula.
4. Write an x under the mole ratio of the substance you
are looking for. Ignore the other substances for now
5. Change the : to =; Solve for x by cross multiplying.
6. The result is the answer in moles.
•
If you need an answer in grams, convert using the mole
formula (with the proper molar mass!)
126
Example: moles to moles
Problem: 8.0 grams of hydrogen are burned with oxygen to make
water. How much oxygen was used?
2 H2
Mole Ratio Row:
+
2
O2
 2 H2O
1
Actual Mass Row: m= 8.0 g
m=36.0 g
Molar Mass Row: M=2.0 g/mol
M=18.0 g/mol
Moles row:
n= 2.0 mol
n =4.0 mol
2
1

4.0 mol x mol
2
Example: gram to gram
Problem: 8.0 grams of hydrogen are burned with oxygen to make
water. How much oxygen was used?
2 H2
Mole Ratio Row:
+
2
O2
 2 H2O
1
Actual Mass Row: m= 8.0 g
m=36.0 g
Molar Mass Row: M=2.0 g/mol
M=18.0 g/mol
Moles row:
n= 2.0 mol
n =4.0 mol
2
1

4.0 mol x mol
2
Example
1.
Balance the equation, or
verify that the equation you
have been given is properly
balanced.
Use the coefficients to find
the mole ratios
Write the amount in moles of
the known reactant under the
corresponding mole ratio
number.
2.
3.
•
4.
If the amount is given in grams,
convert it to moles using the
mole formula.
Write an x under the mole
ratio of the substance you are
looking for. Ignore others
Change : to =. Solve for x by
cross multiplying.
The result is the answer in
moles.
5.
6.
•
If you need an answer in grams,
convert using the mole formula
(with the proper molar mass!)
Problem: 8.0 grams of hydrogen are burned
with oxygen to make water. How much
oxygen was used?
Step 1:
H2 + O2  H2O (skeleton)
2H2 + O2  2H2O (balanced)
Step 2: mole ratios
2:1:2
Step 3: known reactant is 8.0 g hydrogen. To
convert it to moles we must divide by the
molar mass of hydrogen, 2.0; That gives us 4.0
moles of hydrogen. Write this under the
corresponding mole ratio 2 : 1 : 2
4.0
Step 4: write an x under
2_ : 1 : 2
4.0 x
Step 5: cross multiply
2 = 1 so x = 2.0 mol
4.0 x
Step 6: to get the answer in grams, multiply the 2
mol by the molar mass of oxygen (32 g/mol) to
give us the answer 64 g of oxygen is used.
129
Limiting and Excess Reagents
• Sometimes the gram amounts of both reactants
are given to you in a stoichiometry problem.
• It is possible that the given amounts could
produce different amounts of product.
• You must do the stoichiometry twice... Once for
each reactant. The actual amount of product will
be the lower of the two results.
• The reactant that produced the lower result is
called the limiting reagent. The one that
produced the greater result is called the excess
reagent. Not all of the excess reagent will be
used up.
Example of Limiting and Excess
• 10 grams of hydrogen are burned with 64 grams of
oxygen. Which is the limiting reagent? How much of
the excess reagent will be left over?
•
•
•
•
•
•
•
2 : 1 : 2
X
Step 1. Balance equation:
2 H2 + O2  2 H2O
10 g = 5 mol
x
Step 2. Molar ratios
64 g = 2 mol x
Step 3. Known reactant 1
2mol 2mol

Step 4. Ignore reactant 2
5mol x mol
Step 5. Cross multiply
1mol 2mol
Step 6. Find answer

Repeat for reactant 2
2mol x mol
Answer 2 is the correct one.
Oxygen is the limiting reactant
There will be an excess of one mole of
hydrogen which will weigh 2 grams.
Answer 1 = 5 mol water or 90 g
Answer 2 = 4 mol water or 72 g
Assignments
• Read pages 21-24
• Do Question 30 on page 23
• Do Questions 31 and 32 on page 24
132
Review
8
Examples of Chemical Reactions
Overview
Neutralization, Synthesis, Decomposition,
Precipitation, Oxidation, Combustion.
These are all terms associated with specific chemical
reactions.
R
0.8 Examples of Chemical Reactions
• There are many types of chemical reaction.
Among the most important categories are:
•
•
•
•
•
Acid-base reactions
Synthesis, Decomposition and Precipitation Reactions
Endothermic and Exothermic Reactions
Oxidation and Combustion
Photosynthesis and Respiration
• These are just a few of the types. Some your
textbook does not mention include:
• Single Replacement
• Double
Replacement
NOTE: A particular reaction can belong to several categories simultaneously. The
burning of hydrogen is a synthesis, an exothermic reaction and an oxidation reaction.
134
R
–
0.8.1
Reactions
• When an acid and a base are mixed:
The words
Neutralization
and
Titration
are also
associated
with this
process
• The H+ ions from the acid join the OH– ions from the
base to make H2O, that is water.
• The other ions, usually a metal and a non-metal ion,
join to form a salt whose nature depends on the
reagents.
• If the original solutions contained equal amounts of H+
and OH-, then the mixed solution will be neutral.
• If there was a surplus of H+ or OH- ions, then the
resulting solution will be slightly acid or slightly basic.
In general: ACID(aq) + BASE(aq)  WATER(l) + A SALT(aq)
Example: HNO3(aq) + KOH(aq)  H2O(l) + KNO3(aq)
135
0.8.2
Synthesis, Decomposition,
Precipitation
• Synthesis is when two or more reactants
combine to form a single product.
• Eg.
2Na(s) + Cl2(g)  2 NaCl (s)
• Decomposition is when a single reactant
breaks into two or more products.
• Eg. 2 H2O(l)  2H2(g) + O2(g)
• Precipitation is when a solid powder is formed
by the mixing of two solutions.
• Eg
136
R
and
0.8.3
• Endothermic reactions are chemical reactions
that absorb energy. Endothermic reactions
usually make their immediate surroundings
cooler.
• Reactants +
 Products
• Exothermic reactions release heat. They often
make their surroundings warmer.
• Reactants  Products +
137
R
0.8.4
Oxidation and
• Oxidation is process where a substance combines with
an oxidizer (usually O2, but O3, F2, Cl2, N2O and other
substances work as oxidizers too).
• Your textbook incorrectly states that an oxide is always formed,
but sometimes chlorides or fluorides are formed by oxidation.
• Slow oxidation takes time to happen.
• Eg. The rusting of iron: 4 Fe + 3 O2  2 Fe2O3
• Combustion is rapid oxidation that produces heat and
flames.
• Eg. Combustion of gasoline: 2 C8H18 + 25 O2  16 CO2 + 18H2O
138
R
0.8.5
p. 27
and
• Life on Earth depends on two related chemical
processes:
• Photosynthesis is the chemical reaction in which
organisms, such as plants, transform radiant energy from
sunlight into stored chemical energy.
6 CO2 + 6 H2O + energy  C6H12O6 + 6 O2
• Respiration is the process by which organisms release
stored chemical energy in sugars and other organic
compounds in living cells.
C6H12O6 + 6 O2  6 CO2 + 6 H2O + energy
Error in textbook: On p. 27, respiration is referred to as a “combustion” reaction.
What the textbook means, of course, is that is an “oxidation” reaction.
139
Replacement
Single and DoubleDisplacement
• Replacement (or displacement) occurs when
atoms of an element detach from one compound
and attach to a new compound.
• Single replacement pattern A+BC  AB+C
Cu 4  Cu + FeSO4
example: Fe + CuSO
• All simple metal with acid reactions are single replacements
• Double replacement, AB + CD  AD + CB
example: AgNO3(aq) + HCl(aq)  AgCl(s)↓ + HNO3(aq)
• Many precipitation reactions are double replacements.
Assignments
• Page 25 #33
• Page 26 #35-36
• Page 27 #37-38
Review
9
Chemical Bonds
Overview
Electrical forces hold matter together. These forces
come from the attraction between negative
electrons and positive protons in atoms.
We call these forces chemical bonds, and there are
several types of them.
Ultimately, it is chemical bonds that prevent matter
and objects from spontaneously disintegrating .
0.9.0
Chemical Bonds
Sugar, covalent molecule
Salt, ionic crystal lattice
• Chemical bonds are the forces that bind atoms
together into larger structures, such as molecules
or crystal lattices.
• Chemical bonds are the result of exchange or
sharing of electrons between two atoms, which
causes the formation of a compound or diatomic
or polyatomic element.
• There are many types of chemical bond. The four
most important are:
Not Studied
Studied
Studied
Not Studied
•
•
•
•
Metallic:
Ionic:
Covalent:
Weak bonds:
metal to metal, found in alloys
metal to non-metal, found in salts
non-metal to non-metal, found in molecules.
Van der Waals force, Hydrogen bonds
Your textbook has little about metallic bonds, but since we don’t study alloys in detail, this is not a problem .
143
Metallic Bonds and Weak Bonds
(optional)
• A metallic bond consists of trillions of positive metal
ions sharing a vast pool of negative electrons.
• Atoms of several different types of metal can share the same pool
of electrons. That’s why metals can form an infinite number of
alloys instead of specific compounds.
• If a piece of metal is dented or deformed, the disrupted
pool of electrons instantly reforms bonds.
• This accounts for the great malleability and ductility of metals.
• The pool of electrons allows easy passage of other
electrons through the material
• That is why metals are such good conductors
• Weak bonds:
• Hydrogen Bonds: weak forces between polar-covalent molecules.
These account for some crystal structures, like ice.
• Van der Waals Forces: weak forces between particles of solids
Ionic Bonding
0.9.1
p. 28
• An ionic bond forms when electrons are
exchanged between two atoms.
Sodium has an “extra”
electron in its outer shell
cation
Na
Na+
Cl
Cl–
anion
Chlorine “needs” another
electron in its outer shell
X.= 3.16
X.= 0.93
ΔX = 2.23
• This type of bond forms when one of the elements has a
much higher electronegativity (X) than the other. This usually
happens between a metal atom and a non-metal atom.
• Ionic bonds are between negative and positive ions
• Ionic bonds do not form strong, distinct molecules. In most
ionic solids, the ions form a crystal lattice of alternating
positive and negative particles. Some chemists prefer the
term “formula units” to “molecules” when talking about
ionic compounds.
Cl
Alternating particles
A crystal lattice
structure with
alternating ions
do not overlap.
Na+
Cl–
A sodium chloride formula unit
N Cl
Cl A covalent molecule
145
Electro-negativity
and Bond Type
Chart of Electronegativity
• The electronegativity (X )(Greek letter chi or curly x) of an element
can be found from the periodic table in front of your textbook.
• It indicates how much an element attracts electrons.
• The greater the electronegativity difference between two elements,
the more likely they will form an ionic bond.
• No bond is 100% ionic or 100% covalent, but we treat them that
way for simplicity.
• The character of a bond is based on several things, in addition to
electronegativity, so the chart below is an approximation.
ΔX
Character of bond
Name of Bond type
1.7 to 3.9
Over 50% ionic
Ionic
0.4 to 1.7
10% to 50% ionic
Polar-Covalent
0.0 to 0.4
Less than 10% ionic
Covalent
0.9.2
p. 29
Covalent Bonding
• A covalent bond forms when electrons are
shared between two atoms.
• This type of bond forms when two elements have
similar electronegativity. This usually happens between
two identical atoms, or between two non-metal atoms.
• Covalent bonds can be single (sharing one pair of
electrons), double (sharing two pairs) or triple (sharing
three pairs)
O C O
• Covalent compounds form true, strong molecules.
Shared electrons in They are sometimes referred to as molecular
overlapping shells
compounds.
147
0.9.2
p. 29
Illustrating Covalent Bonds
With Rutherford-Bohr models:
With Lewis electron dot diagrams:
Electrons shared between Carbon
and one Hydrogen atom.
Or you can just circle an electron
pair to show they are shared.
Another
Electrons shared between Carbon
and another Hydrogen atom.
Another
In either case, we draw the atoms to show a stable number of electrons (usually 8) in
the outer shell of each atom involved in the covalent bond.
Cl
N Cl
Cl
 Another way to illustrate covalent bonds is with overlapping circles
148
Assignments
• Page 25 #33
• Page 26 #35-36
• Page 27 #37-38
Energy
Overview
Chemical reactions involve energy. Some reactions
absorb energy, others release it.
There are dozens of specific types of energy, but we
can group them all into two main categories: Kinetic
energy, and Potential Energy.
R
Energy
0.10.0
p. 30
• Energy is the ability to do work or make a change.
• Energy is classed into Kinetic and Potential
• There are many sub-types of energy, a few examples of
which are listed in the table below:
Form of Energy
Associated with
Example
Mechanical Kinetic Energy
An object’s movement
Car driving along a road
Thermal Kinetic Energy
Agitation of particles
Boiling water
Radiant Kinetic Energy
Electromagnetic waves
Light, microwaves, radio waves
Gravitational Potential Energy
Object’s position above ground
Water behind a dam
Elastic Potential Energy
Compressed/stretched materials
A spring that has been stretched
Electric Potential Energy
Force between electric charges
Charged particles in a storm cloud
Nuclear Potential Energy
Stored in the nucleus of atoms
Uranium in a reactor
Chemical Potential Energy
Stored in the bonds of molecules
Energy in gasoline or glucose
* Potential energy can rapidly change into kinetic, for example, electrical potential energy can
create electricity, a form of kinetic energy. Nuclear energy can create thermal energy.
151
Units of Energy
• In chemistry, energy is measured in joules
• if there are more than a thousand joules then we can measure
them in kilojoules. (1 kJ = 1000 J)
• A joule is the amount of energy required to
accomplish any of the following:
•
•
•
•
•
Moving a force of one newton a distance of one metre.(N∙m)
Passing a 1 A current through a 1Ω resistor for 1 s (A2∙Ω∙s)
Using 1 Watt of power for 1 second (W∙s)
Moving a coulomb of electrons through 1 volt (C∙V)
Keep a human body alive for 1/100th of a second.
• In other fields energy can be measured in units such
as thermal calories (cal), food Calories (kcal or Cal),
ergs, British Thermal Units (BTU), Kilowatt-hours
(kW∙ h), Tonnes of TNT (or kilo or megatonnes).
Energy Comparisons
• Energy units smaller than a joule
• 1 joule
• 1 joule
= 6.24 x 10 18 electron Volts (eV)
= 10 000 000 ergs (erg)
• Energy units greater than a joule
•
•
•
•
4.184x10 9 joules
3 600 000 joules
4184 joules
1055 joules
= 1 ton of TNT
= 1 kilowatt-hour (kW∙h)
= 1 kilocalorie (kcal)*
= 1 British Thermal Unit (BTU)
• A joule is:
• The energy required to lift a small apple (102g) 1 metre
• The energy a 1 watt LED uses every second.
* A kilocalorie is also known as a food calorie. It is the same unit identified as a
calorie (Cal) on food packages. A thermal calorie (cal) is 1/1000 of a food calorie, or
only 4.18 joules.
0.10.1
p. 30
Kinetic Energy
• Kinetic energy is the energy associated with
the movement of an object, or with the
movement of its particles (molecules).
• Kinetic energy depends on the mass of the
object and the velocity of its motion.
Where: Ek= kinetic energy
m= mass of the object
v= velocity of the object
154
0.10.2
p. 30
Potential Energy
• Potential Energy is energy stored in a body
that can be transformed into another form of
energy.
• Potential energy is sometimes referred to as “hidden
energy”, since it is difficult to observe and measure.
• There are several types of potential energy,
including:
• Gravitational Potential Energy (important in physics)
• Chemical Potential Energy (important in chemistry)
155
R
• Gravitational Potential Energy is the product
of an object’s mass, its height above the
ground, and the gravitational acceleration.
Where: Ep = Gravitational Potential Energy in joules
m = mass of the object in kilograms
g = gravitational acceleration (9.8 m/s2 on Earth)
h = height of the object above a reference point (such as the ground)
0.10.2
p. 30
• Chemical Potential Energy (Enthalpy) is
associated with the energy in the bonds
between the particles of a material.
We will devote a section later in the course to calculating enthalpy.
156
R
0.10.3
p. 32
Conservation of Energy
• The law of conservation of energy states that
energy cannot be created or destroyed in
chemical reactions, but it can be changed
from one form to another.
• Potential energy can change to kinetic and vice versa
• Mechanical Energy is the total energy of an
isolated system.
Where: Em = total Mechanical Energy
Ep = Potential Energy
Ek = Kinetic Energy
157
R
0.10.4
p. 32
Thermal Energy & Temperature
• Thermal energy or “heat” is a form of energy
possessed by a substance due to the agitation
of its particles. It depends on:
• The mass of the substance
• The temperature of the substance
• The specific heat capacity of the substance
Where: Q = amount of heat energy in joules
m = mass of the substance heated in grams (usually the water in a calorimeter)
c = specific heat capacity of the substance heated, in j/g∙°C
ΔT = the change in temperature in °C
158
R
Fluids
0.11.1
p. 33
• Compressible & Incompressible Fluids
• Substances that flow, like liquids and gases, are fluids
• Gases are compressible fluids
• Liquids are incompressible fluids
Gas
Liquid
• Pressure
• Pressure is the force exerted on a surface.
• The standard unit of pressure is the kilopascal (kPa)
• Formula for pressure: Pressure = Force divided by Area.
Standard Atmospheric Pressure:
Ps = 101.3 kPa
159
END of MODULE 1
• Prepare for the module 1 test.
– Make sure that you have read and understood all
pages up to page 33 in your text book.
– Make sure you have downloaded and read the
study notes called “Chemistry Unit 0” from my
web site: chem534.wikispaces.com
– Prepare your own study notes by highlighting all
the sections you think are important, and then
recopying the highlights into your notebook.
– practice conversions and significant figures
160