Pages 3 to 33
“Quantum Chemistry”
Target Completion Date: October 1
About Slide Icons
Very Important Points
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!
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 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!
• 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.
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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.
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Conversions
• You must be able to do ALL standard metric
conversions, especially:
– Litres to millilitres, millilitres to litres
– Grams to kilograms, kilograms to grams
• Other conversions you will learn during the
course of the year:
– Temperature: degrees celcius (℃) to kelvins (K)
– Pressure: kilopascals (kPa) to millimetres (mmHg)
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
…
x10-5
micro (μ)
X10-6
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
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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:
or
m
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!
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
=50g
=?
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:
V
𝜌
=12cm x 5cm x 2 cm
= 120 cm3
= 50 g / 120 cm3
= 0.46 g/cm3
Always state your answer in a complete sentence, with appropriate units.
Answer: The density of the block is 0.46 g/cm3 (or 0.46 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”
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”.
Uncertainty
In math, numbers are considered pure, abstract
things. In math, 2.00, 2.0 and 2 are considered the
same, they all represent perfect number 2.
In science, numbers are considered to be
measurements, and all measurements have some
degree of uncertainty. They are almost never
considered perfect!
The absolute uncertainty of a measurement is
usually ½ of a measuring instruments smallest
gradation. If a graduated cylinder is marked in
millilitres, then each measurement taken with that
cylinder has a ±0.5 mL uncertainty.
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. 20.0 ±0.5 mL or 25.5 ±0.5 mL are
acceptable readings.
• With the same graduated cylinder, it would be wrong to
write 20 ±0.5 mL or 25 ±0.5 mL or even 20.00 ±0.5 mL
• In science 20 mL, 20.0 mL and 20.00 mL have
different meanings with respect to
.
Rules for Significant Figures
Interpreting Significant Digits
1. All 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.
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 any of the zeros at the end.
5000
has 1 or 4 significant digits
250
has 2 or 3 significant figures
123 000 000 has 3 or 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
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, but not one that adds to the precision of the answer.
Same Number, Different Precision
Try to avoid the “ambiguous”
situation in your answers. If an
200.000
6 significant digits
answer ends in zero, or worse, in
200.00
5 significant digits
several zeros, indicate whether it
200.0
4 significant digits
should be interpreted as “exactly”
200*
Ambiguous* 1 to 3 SD
or “approximately”.
2
2.00 x 10
3 significant digits
Better still, convert it to scientific
2
2.0 x 10
2 significant digits
notation, and leave only the zeros
2 x 102
1 significant digit
you know are accurate.
Eg. If your answer is 2500 mL, but
*Your textbook says to call this 3
you only measured to the nearest
significant figures. Traditional
measurement would call it 1
10mL, then write 2.50 x 103 mL.
significant figure. Written this way it
That way every one will know its
is ambiguous.
accurate to 3 significant figures
Avoid writing answers that end
Number
Precise to
with zeros and no decimal!
Math with Significant Figures
• The basic rule for math is that you do not
improve significance by multiplying or dividing
numbers:
53.81 m x 2.43 m = 131 m2 NOT 130.7583 m2 !!!
Why? Because the least precise measurement had
3 significant digits, so our answer should not have
more than 3 significant digits!
The technique for addition and subtraction is slightly
different (see p.396 ) but the concept is the same.
You cannot make your result better than your
measurements!
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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
One atom
• A molecule may contain one atom, or it may contain
several thousand atoms, or any number between.
– A molecule is represented by its formula
Ne
Ne
• Water molecules, for example, are represented by the
formula H2O, shown below:
Cl
S
several thousand atoms
DNA
2 atoms of
hydrogen
Cl
O
H
H2O
H
1 atom of
oxygen
SCl2
18
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 the same way as
regular atoms, 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 almost always form cations (+), non-metals
may form anions (-)
19
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.
Most 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.
20
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.
21
H
H
N
H
+
H
O 2–
O S O
O
Big Fat Ions
(Polyatomic 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 (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-
23
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Representation of Atoms
0.2.0
• Early Representations
– Democritus (c.450 BC) 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
24
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0.2.1
Page 5
1. 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.
25
Early
Rutherford
model
Revised
Bohr
model
Page 5
0.2.2
2. The Simplified Atomic Model
– 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
26
• 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
Shells
(1/1680)
28
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0.2.3
3. 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.
– 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
29
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 movement of dots, as
in the drawing of a nitrogen molecule (N2) shown here:
30
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.
31
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 .
32
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…
33
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
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.
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.
In-Line Notation for a Carbon-14 atom
(carbon-14 is an isotope, or alternate form of normal carbon)
Isotope or Mass
number. Represents
the number of
nucleons in a
particular atom
Atomic number “Z”
represents the
number of protons in
this atom
Subtracting the Mass
# and the Atomic #
“Z” gives the number
of neutrons in the
atom
An average carbon atom weighs 12.01 amu
according to the periodic table. But no atom
of carbon has that exact weight. For every
thousand atoms that weigh exactly 12 amu,
a few weigh more. This one weighs 14 amu
14
4–
6
2
(6p+, 8n0, 10e-)
8
Configuration of this atom
Valence (4)
or
Ionic Charge (4–)
or
Oxidation # (–4)
Valence is the
number of bonds
the atom is likely
to form. Ionic
charge is the most
likely charge an
ion will have.
Number of atoms in
a molecule, such as
C2H4
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
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!)
38
The Periodic Table
with 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.
39
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
40
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.
41
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.
42
Use the Periodic Table
to Find the Electron Arrangement of an Atom
Eg. Find the electron arrangement of Iodine (I)
H
II
A
III
A
IV
A
V
A
VI
A
2
Li
Be
B
C
N
O
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
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
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!)
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.
– The 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.
45
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+ Cl-
Ca2+ O2O23+
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+
46
Non-metal ion endings
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
47
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 acceptable 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.
Monosodium glutamate is an organic compound that does not follow the rules.
Dealing with
• Some metal elements have more than one
possible valence. Copper, for example, can have
a valence of 1+ or 2+, depending on what
compound it is in (eg. CuCl or CuCl2). Since we
This
don’t use prefixes in naming ionic compounds,
copper
we shouldn’t use copper dichloride. We need a
ion has
a
new rule!
charge
of 1+
– If a metal is polyvalent, we include its current valence
in roman numerals inside parenthesis within an ionic
compound name, for example:
– CuCl = Copper (I) chloride (not copper monochloride)
– CuCl2 = Copper (II) chloride (not copper dichloride!)
This copper ion
must have a
charge of 2+
49
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
50
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
51
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-
Copper(I)oxide
Cuprous oxide
Cu+, O2-
Copper(II)oxide
Cupric oxide
Cu2+, O2-
Fe3O4
Cu2O
cuprite
CuO
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)
52
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
• 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
53
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.
54
Electronegativity
(how much an atom attracts electrons)
One of the uses of electronegativity is to decide which element goes first in a
formula or name. Usually the element with the lowest electronegativity goes
first. Therefore it is called carbon dioxide (CO2), NOT dioxygen carbide (O2C).
I
1 2.2
VII
I
II
III
IV
V
VI
VII
0.0
He
2.0
B
2.6
C
3.0
N
3.4
O
4.0
F
0.0
Ne
H
2
1.0
Li
1.6
Be
0.7 1.0
0.9 1.4
1.5
1.9
2.0
2.4
3
0.9
Na
1.3
Mg
III
B
IV
B
4
0.8
K
1.0
Ca
1.4
Sc
5
0.8
Rb
0.9
Sr
1.2
Y
6
0.8
Cs
7
0.7
Fr
2.5
2.9
3.0
3.1
V
B
VI
B
VII
B
1.5
Ti
1.6
V
1.7
Cr
1.6
Mn
1.8
Fe
1.9
Co
1.3
Zr
1.6
Nb
2.2
Mo
2.1
Tc
2.2
Ru
0.9
Ba
1.3
Hf
1.5
Ta
1.7
W
1.9
Re
0.9
Ra
Rf
Db
Sg
Bh
3.2
3.5
3.6
4.0
I
B
II
B
1.6
Al
1.9
Si
2.2
P
2.6
S
3.2
Cl
0.0
Ar
1.9
Ni
1.9
Cu
1.7
Zn
1.8
Ga
2.0
Ge
2.2
As
2.6
Se
3.0
Br
0.0
Kr
2.3
Rh
2.2
Pd
1.9
Ag
1.7
Cd
1.8
In
2.0
Sn
2.0
Sb
2.1
Te
2.7
I
0.0
Xe
2.2
Os
2.2
Ir
2.2
Pt
2.4
Au
1.9
Hg
1.8
Tl
1.8
Pb
1.9
Bi
2.0
Po
2.2
At
0.0
Rn
Hs
Mt
Ds
Rg
Cn
Uut Uuq Uup Uuh Uus Uuo
VIII
B
There are a few exceptions, like CH4 and NH3, where the more electronegative elements are written first. These formulas
have been used for years, and are based on organic chemistry concepts, so it’s unlikely we will change them.
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
• Many chemists simply say “hydrogen sulphide” instead of “dihydrogen
sulphide” for the compound H2S. Since H2S is the only common sulphide
of hydrogen, 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.
• Water can be called hydrogen oxide, but it is perfectly acceptable to use
“dihydrogen oxide” or even “dihydrogen monoxide”
57
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
58
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
59
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. 60
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.
61
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? 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 (…tone), 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
64
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
65
0.5.2
Molar Mass
• 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
66
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 and phosphorus, 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
67
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
68
Practice
• Page 14, #12, 13, 14
• Practice sheet:
• Moles and Molar mass
69
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.
70
0.6.1
Phase Change
• 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 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.
71
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
Melting (fusion)
Solid
Freezing (solidification)
Liquid
Liquid
72
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 particles
Vibration only
Rotation and
vibration
Rotation, vibration
and translation
73
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)
74
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.
75
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)
76
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!
77
• 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.
78
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
79
0.6.3
Concentration and Dilution
• Concentration is the ratio of dissolved solute to total
amount of solution.
• General formula for concentration is:
𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 =
𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
• 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.
80
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
81
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
82
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
83
Assignments
• Read pages 15 to 18
• Do page 16
• Questions 15 to 17
• Do page 19:
• Questions 18 to 23
84
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.
85
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) + Cl-(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
86
Assignments
• Read page 19
• Do page 20
• Questions 24-27
• Question 28
87
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
88
• 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.
89
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
90
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.
91
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.
92
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!)
93
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 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 8g hydrogen. To convert
it to moles we must divide by the molar mass
of hydrogen, 2.0; That gives us 4 moles of
hydrogen. Write this under the corresponding
mole ratio
2:1:2
4
Step 4: write an x
2:1:2
4 x
Step 5: cross multiply
2 = 1 so x = 2 mol
4 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.
94
Assignments
• Read pages 21-24
• Do Question 30 on page 23
• Do Questions 31 and 32 on page 24
95
0.8 Examples of Chemical Reactions
• There are many types of chemical reaction.
Among the most important types 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
96
–
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)
97
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.
• Precipitation is when a solid powder is formed
by the mixing of two solutions.
• Eg
98
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 +
99
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 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
100
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.
101
Assignments
• Page 25 #33
• Page 26 #35-36
• Page 27 #37-38
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
three most important are:
Not Studied
Studied
Studied
• Metallic:
• Ionic:
• Covalent:
metal to metal, found in alloys
metal to non-metal, found in salts
non-metal to non-metal, found in molecules.
Your textbook has little about metallic bonds, but since we don’t study alloys in detail, this is not a problem .
103
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
104
Electronegativities
(Supplemental)
• 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
> 50% ionic
Ionic
0.4 to 1.7
10% to 50% ionic
Polar-Covalent
0.0 to 0.4
<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.
106
Illustrating Covalent Bonds
With Rutherford-Bohr models:
With Lewis electron dot diagrams:
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
107
R
Energy
0.10.0
p. 30
• Energy is the ability to do work or make a change.
• There are many types of energy, a few of which are listed in
the table below:
Form of Energy
Associated with
Example
Kinetic Energy
An object’s movement
Car driving along a road
Thermal Energy
Agitation of particles
Boiling water
Radiant Energy
Electromagnetic waves
Light, microwaves, radio waves
Gravitational* Energy
Object’s position above ground
Water behind a dam
Elastic* Energy
Compressed/stretched materials
A spring that has been stretched
Electric* Energy
Force between electric charges
Charged particles in a storm cloud
Nuclear* Energy
Stored in the nucleus of atoms
Uranium in a reactor
Chemical* Energy
Stored in the bonds of molecules
Energy in gasoline or glucose
* The word “potential” is often inserted to indicate that these associated with potential energy.
108
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
109
0.10.2
p. 30
0.10.2 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)
110
R
• Gravitational Potential Energy is the product
of an object’s mass, its height above the
ground, and the gravitational acceleration.
0.10.2
p. 30
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)
• 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.
111
R
0.10.3
p. 30
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
112
R
0.10.4 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
𝑄 = 𝑚𝑐∆𝑇
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
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R
0.11 Fluids
• Compressible & Incompressible Fluids
• Substances that flow, like liquids and gases, are fluids
• Gases are compressible fluids
• Liquids are incompressible fluids
• 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.
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END of MODULE 1
• Prepare for the module 1 test.
– Read up to page 33 in your text book
– Prepare study notes.
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