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Chemistry 30
Introduction and Review
Resources:
In-class notes and handouts
Objectives:
1.
Recognize and interpret domestic hazard symbols.
2.
Recognize and interpret WHMIS symbols.
3.
Recognize various pieces of lab equipment.
4.
Identify the location of safety devices.
5.
List safety procedures.
6.
List the important base and derived units of the metric system (2-2).
7.
List the important prefixes in the metric system (2-2).
8.
Perform metric conversions (2-3).
9.
Use and perform calculations in scientific notation (2-3).
10.
Discuss uncertainty (2-3).
11.
Compare and contrast between accuracy and precision (2-3).
12.
Use significant digits in calculations (2-3).
13.
Use both the Stock and Classical systems to name ionic compounds (chapter 7).
14.
Use the numbering system to name molecular compounds (chapter 7).
15.
Be able to name acids ( chapter 15).
16.
Be able to name simple organic compounds (chapter 20).
17.
Be able to write and balance chemical equations (chapter 8).
18.
Be able to identify types of chemical reactions and complete equations based on activity series (chapter 8).
Vocabulary:
accuracy
acute
biohazard
carcinogen
chronic
molecular
ionic bond
alkyne
cyclic
monoatomic
dependent variable
exponent
flammable
independent variable
oxidize
ionic
alkane
aromatic
hydrocarbon
diatomic
precision
radioactive
toxic
uncertainty
binary
covalent bond
alkene
aliphatic
synthesis
polyatomic
1
double replacement
single replacement
decomposition
combustion
activity series
organic
inorganic
Lab Safety
Hazardous Materials in the Home
Poison
- skull and crossbones symbol
- poisons act by entering the body in some way
- poisons can enter the body in one of three ways:
a) ingestion (eating)
b) inhalation (breathing)
c) absorption through the skin
- poisons can cause anything from mild illness to death, depending on the nature of the poison
Corrosive
-
Flammable
- symbolized by a flaming surface
- are substances which can burn easily or cause other materials to burn
Explosive
- symbolized by an exploding bomb
- can cause injury or death as a result a blast or because of the materials expelled by the blast (metal
shards)
- usually are pressurized aerosol containers which may explode when heated
Radiation
- symbolized by a 3-sided fan
- radioactive materials emit high energy atomic particles or high energy radiation (x-rays, gamma
rays), or both
- found in smoke detectors and involve no danger if kept at a safe distance
-
symbolized by a hand eaten away to the bones in a beaker of liquid
are chemicals which can act on clothing, skin, eyes or internally by drinking or eating
can cause symptoms ranging from mild rash to serious skin damage
can damage clothing
can cause blindness
can cause death if ingested
Octagon represents that the contents are dangerous.
Triangle represents that the container is dangerous.
WHMIS
Workplace
Hazardous
Materials
Information
System
WHMIS is a system of warning symbols and information sheets which detail the danger, safe handling and disposal
of a variety of chemical substances in Canada.
All chemicals handled in Canada must be labeled using the WHMIS system.
There are 8 classes of hazards, each with its own symbol.
Toxicity
2
Acute Toxicity -
refers to a substance which has immediate effects. If you were to eat breathe, or absorb the
substance through your skin, the effects would happen in short order. Effects can include
illness, organ damage or death.
Chronic Toxicity - refers to the effects of a substance through repeated exposure over a long period (weeks,
months or years). Effects may be similar to those of acute toxicity, like those which occur as
a result of long term alcohol or cigarette abuse. The effects can also include cancer, allergies
or chronic diseases (bronchitis, emphysema, cirrhosis of the liver, etc.)
Biohazard -
refers to an infectious agent (bacteria, virus or some other organism) which may spread
disease if improperly handled.
In Case of An Accident:
Inhaled Poison
Remove the patient to fresh air and apply artificial respiration if necessary. Keep the
victim warm with blankets.
Contact of Poison with
Skin or Eyes
Flood affected area with water, for at least 5 minutes. Remove contaminated clothing.
DO NOT attempt to use chemical antidote.
Swallowed Poison
If the person is conscious and able to swallow, immediately dilute the poison by giving
the victim 2 to 4 cups of milk or water.
Swallowed Corrosives
DO NOT INDUCE VOMITING. Give milk and water. If vomiting occurs naturally,
hold head below hips to avoid choking. (Note: Corrosives include drain cleaners, toilet
bowl cleaners, ammonia, oven cleaners, turpentine, kerosene, furniture polish, gasoline,
pine oil and bleaches.
Safety in the Chemistry Lab
You must know the location of safety equipment in the lab:
-
fire extinguisher
absorbent
-
eye wash
fire blanket
-
sinks
Lab Procedures and Rules
1.
No eating or drinking in the lab.
2.
Treat all chemicals as if they were hazardous:
-
never taste chemicals.
wash hands after chemicals have been handled.
wear eye protection when instructed.
3.
Never perform unauthorized experiments.
4.
Report all accidents immediately. Do not attempt to clean it up until checking with the teacher.
3
5.
If you get a chemical solution in your eye, do not wait for the teacher; go to the eyewash station immediately
and wash the eye for at least 5 minutes.
6.
If you get chemicals on your clothes, wash the clothes thoroughly.
7.
Do not wear loose clothing during a lab. Tie long hair back.
8.
Do not sit on the lab bench; you do not know how clean it is.
9.
Clean all equipment thoroughly and put it back where it belongs after a lab.
10.
Follow directions concerning the safe disposal of chemicals and solutions.
11.
Clean your lab station thoroughly after a lab.
Metric System
SI BASE UNITS
Quantity
length
mass
volume
temperature
time
amount of matter
electric current
Base Unit
Symbol
metre
gram
litre
kelvin
second
mole
ampere
m
g
L
K
s
mol
A
SI DERIVED UNITS
Quantity
Name of Unit
Symbol
density
kilogram per
cubic metre
kg · m-3
Expressed in Terms
of SI Base Units
kg · m-3
(kg/m3)
force
Newton
N
kg · m · s-2 (kg · m / s2)
pressure
Pascal
Pa
N · m-2
heat energy
Joule
J
N·m
4
(N / m2)
METRIC PREFIXES
This table lists the metric prefixes which are significant:
Prefix
Symbol
Factor by which Unit is Multiplied
Exponential Notation
1012
tera
T
1 000 000 000 000
giga
G
1 000 000 000
mega
M
kilo
k
1 000
103
hecto
h
100
102
deca
da
10
101
1
100
109
106
1 000 000
THE BASE UNIT
deci
d
0.1
10-1
centi
c
0.01
10-2
milli
m
0.001
10-3
micro

0.000 001
10-6
nano
n
0.000 000 001
10-9
pico
p
0.000 000 000 001
10-12
SCIENTIFIC NOTATION
The rules for putting numbers into scientific notation are simple. They work on the assumption that all numbers contain a
decimal point. For instance, the number 125 can be visualized as 125.000; 31 can be visualized as 31.000. There are two
basic rules:
1)
For numbers larger than 1, scientific notation is determined by counting the number of times the decimal place must
be moved to the left, leaving just one number to the left of the decimal. The number of times the decimal must be
moved is the power of 10 (the exponent).
Example:
3000 = 3000.0 = 3 x 103
454 000 = 454 000.0 = 4.54 x 105
3 860 000 = 3 860 000.0 = 3.86 x 106
602 000 000 000 000 000 000 000 = 6.02 x 10 23
5
2)
For numbers smaller than 1, scientific notation is determined by counting the number of times the decimal
place must be moved to the right, leaving just one number to the left of the decimal. The number of times
the decimal must be moved is the exponent, but it is a negative number.
0.068 = 6.8 x 10-2
0.000 049 3 = 4.93 x 10-5
0.000 000 002 41 = 2.41 x 10 -9
Example:
If the decimal does not have to be moved, the exponent is zero. For example, the number 1.23 written in
scientific notation is
1.23 x 100.
IF A NUMBER IS LARGER THAN 9999 OR SMALLER THAN 0.001 IT MUST BE WRITTEN IN
SCIENTIFIC NOTATION. Between these extremes you may use either decimal or scientific notation.
Metric Conversion
Using unitary rates
-
this method involves multiplication using conversion factors.
eg.
-
15 mm =
m
to solve, set up a ratio such that the units of the unknown are the units desired:
15 mm x m =
mm
m
(m/mm is the conversion factor; when the units are cancelled out only 'm' is left)
-
next fill in the units of the conversion factor. These units are unitary rates; they have a value of 1. In this case
1 m is equal to 1000 mm. You must be very familiar with the value of the prefixes to make this work.
15 mm x
-
1 m
=
1000 mm
m
now finish the calculation:
(15 x 1 )( mm x m ) = 0.015 m
100
mm
Make sure you record the examples in class.
Assignment:
1) 16 kg = ? g
2) 632 cm = ? km
3) 2.18 x 105 μN = ? mN
4)
5)
6)
0.036 s = ? ns
7120 Mg = ? Tg
8.88 x 10-10 kL = ? mL
6
UNCERTAINTY
This is a fundamental issue in science. A scientist knows nothing for certain; there are no laws in science. A theory can
never be proven; it may only be disproven.
Every measurement done in science has some amount of uncertainty. For instance, if you take the mass of a substance the
scale may read 46.58 g. The last digit is rounded off, so the mass could be as low as 46.575 g and as high as 46.584 g. A
more sensitive scale could be used to reduce the uncertainty, but it will always be there.
Two terms associated with uncertainty are accuracy and precision. They are defined as follows:
Accuracy - how close a measured or calculated value is to a known or real value.
Precision - how close many repeated measurements are to each other.
Scientists repeat measurements as a way to reduce uncertainty. If a number of measurements are very close to one another,
they have good precision and the scientist can be assured that the average of the measurements is likely close to the actual
value (accurate)
SIGNIFICANT DIGITS
Refers to the certain digits in any measurement. This is related to issues of uncertainty; the results of a calculation can be no
more precise than the least precise measurement which goes into it.
To determine the number of significant digits you must be able to handle zeros and their relationship to the non-zero digits in
a number. Note the following rules:
1.
All non-zero numbers are considered significant; that is, they are counted
123 has 3 significant digits; 1267 has 4 s.d.
2.
There are two situations where zeros are significant:
i)
Zeros between two non-zero numbers are considered significant.
102 has 3 s.d.; 10203 has 5 s.d.; 1002 has 4 s.d.
ii)
A zero at the end of a decimal number is significant.
12.00 has 4 s.d.; 0.010 has 2 s.d.; 1200.000 has 7 decimal places
Note: in the last example, all the zeros are significant because they are between s.d.
3.
In any other situation zeros are not considered significant:
i)
For a number larger than 1, a zero between the decimal and the first non-zero number is not significant.
120 has 2 s.d.; 10200 has 3 s.d.; 130 000 000 has 2 s.d.
ii).
For a number smaller than 1, a zero between the decimal and the first non-zero number is not significant.
0.0012 has 2 s.d.; 0.02102 has 4 s.d.; 0.000 000 001 has 1 s.d.
7
Exact numbers - are numbers that are defined (conversion factors in the metric system) or numbers which result from
counting objects (like the coefficients used to balance chemical equations). Exact numbers are said to have an infinite
number of significant digits.
Rounding off - is necessary when a number from a calculation must have the number of significant digits reduced. The
rules for rounding are as follows:
o
if the digit following the last digit to be kept is greater than 5, the last digit is increased by 1
e.g.
o
123.5
123.44
rounded to 4 s.d. is now
123.4
if the digit following the last digit to be kept is equal to 5, followed by a nonzero digit, the last digit is
increased by 1
e.g.
o
rounded to 4 s.d. is now
if the digit following the last digit to be kept is less than 5, the last digit stays the same
e.g.
o
123.46
123.452
rounded to 4 s.d. is now 123.5
if the digit following the last digit to be kept is equal to 5, and not followed by a nonzero digit , the last digit is
increased by 1 only if it produces an even number
123.45
123.55
rounded to 4 s.d. is now
rounded to 4 s.d. is now
123.4
123.6
Rule for addition and subtraction
Add or subtract and then round-off so that the answer is no more precise than the least precise number in the
calculation..
Rule for multiplication and division
Multiply or divide and then round-off so that the answer has no more significant digits than the number with the
fewest significant digits in the calculation. Remember that any exact numbers do not enter into the determination of
least significant digits.
For long, multi-step calculations:
Do not round off the number at each step in your calculator; keep the entire number, with all its decimal
places and use it in the next step. The danger is that you introduce ROUNDING ERROR.
Record the examples in class.
8
Multiplication in Scientific Notation
The rule is that the integers (the first number in each scientific notation) in each number are multiplied, with the
resulting number becoming the integer in the answer. The exponents (the powers of 10) are added; the answer
becomes the exponent in the answer. For example:
(3.4 x 102)(2.0 x 103)
Problem:
Solution:
1. Multiply the integers
3.4 x 2.0 = 6.8
2. Add the exponents
102 x 103 = 102+3 = 105
3. Combine
6.8 x 105
If the integer in the solution has more than one digit to the left of the decimal point the scientific notation must be
corrected:
(6.0 x 103)(2.5 x 107)
15.0 x 1010 = 1.50 x 1011
Problem:
Solution:
Division in Scientific Notation
Division in s/n is similar to multiplication, except that in the case of division the integers are divided, and the
exponent of the second number is subtracted from the first. For example:
Problem:
Solution:
(3.4 x 102) / (2.0 x 103)
1. Divide the integers
3.4 / 2.0 = 1.7
2. Subtract the exponents
102 / 103 = 102-3 = 10-1
3. Combine
1.7 x 10-1
If the integer in the solution has less than one digit to the left of the decimal point the scientific notation must
be corrected:
Problem:
Solution:
(6.0 x 103)(8.0 x 107)
0.75 x 10-4 = 7.50 x 10-5
9
Addition and Subtraction in Scientific Notation
These two operations are slightly different from multiplication and division. Numbers that are expressed in s/n can
only be added or subtracted if the exponents are the same. If the exponents are the same the integers are added or
subtracted and the exponent is carried into the solution. For example:
Addition:
(2.07 x 106) + (1.30 x 106) = 3.37 x 106
Subtraction:
(2.07 x 106) - (1.30 x 106)
= 0.77 x 106 = 7.70 x 105
If the exponents are different there are two methods which may be followed to solve the problems:
1)
Move the decimal of one number to change the exponent.
Example:
=
=
2)
(2.75 x 103) + (3.20 x 102)
(2.75 x 103) + (0.32 x 103)
3.07 x 103
Convert s/n numbers back to normal notation. Add or subtract as required, then convert the answer back to s/n
format.
Example:
=
=
=
(5.75 x 104) - (2.37 x 103)
57 500 - 2 370
55 130
5.51 x 104
NOMENCLATURE
There are two types of compounds; inorganic and organic. There are two types of inorganic compounds; molecular and
ionic. Molecular compounds are bound together almost exclusively with covalent bonds and are compounds composed of
non-metallic elements. Ionic compounds are bound together by ionic bonds and are compounds composed of metallic and
non-metallic ions. Organic substances are comprised almost exclusively of hydrogen and carbon (hydrocarbons).
Further, most compounds we will deal with are binary compounds, compounds made up of two parts. Thus when naming a
compound one must consider the two parts involved, as well as whether the compound is molecular or ionic. The naming
system used hinges on whether the compound is molecular or ionic.
Finally, naming is done using two principles:
1)
2)
The name must be as simple as possible; nothing unnecessary is added to a name.
The name must be unique to a substance; two different structures cannot have the same name.
Elements
Are substances made up of one kind of atom. Most elements on the Periodic Table are monoatomic. Eight elements are
diatomic (H2, N2, O2, F2, Cl2, Br2, I2 and At2). Two elements are polyatomic (S8, P4).
10
Common Substances
These are compounds that are know by names other than their systematic names. You must be familiar with their formulas,
systematic anmes and common names. Here is the list:
Formula
Chemical Name
Common Name
H2 O
hydrogen oxide
water
NaCl
sodium chloride
table salt
HCl
hydrogen chloride
hydrochloric acid
HNO3
hydrogen nitrate
nitric acid
H2SO4
hydrogen sulfate
sulfuric acid
H3PO4
hydrogen phosphate
phosphoric acid
CH3COOH
acetic acid
vinegar
CaSO4
calcium sulfate
gypsum (dry wall)
NH3
nitrogen trihydride
ammonia
H2 O2
hydrogen peroxide
hydrogen peroxide
C2H5OH
ethanol
drinking alcohol
CH3OH
methanol
wood alcohol
CH4
methane
natural gas
O3
ozone
ozone
C12H22O11
sucrose
table sugar
KCl
potassium chloride
potash
NaOH
sodium hydroxide
lye, caustic soda
CaO
calcium oxide
lime
If you are asked to name one of these formulas, give the name written here in bold.
11
Naming of ionic compounds
Ionic compounds can be named using either of two similar methods. In order to name ionic compounds one must first
understand how ionic compounds come to be. Ionic compounds are formed as a result of cations and anions which are
attracted to each other because of their opposite charge and join in ratios which result in a net charge of zero. Thus Na 1+ and
Cl1- form the ionic molecule NaCl. Ca2+ and F1- form the ionic molecule CaF2. Mg2+ and O2- form the ionic molecule
MgO.
Naming of ionic compounds is a way of clearly stating their chemical formula. The name which a compound receives
gives enough information to determine exactly its chemical formula. The method is simple; the name of the cation comes
first, the anion is second. Tables 4 and 5 of the package of tables gives the name of common cations and anions, including
polyatomic cations and anions. When giving a chemical formula, it is important to identify both the ions involved and their
charge.
Knowing the charge of the ions is very important. For instance, in the compound FeCl 2 the iron is iron(II), or ferrous. Thus
the compound name is either iron(II) chloride or ferrous chloride. There is a difference between iron(II) chloride and
iron(III) chloride.
You are required to use the Stock system in this class, but you should be able to recognize Classical nomenclature when it
comes around.
Stock Method
-
Classical Method -
for naming ionic compounds (metal + non-metal)
uses the oxidation number to identify the charge on the cation (iron II, iron III) for those cations
which has more than one possible charge
for naming ionic compounds
uses archaic names to identify the charge on the cation (ferrous = Fe2+, ferric = Fe3+)
To use either method for naming compounds:
1.
Identify the cation and the anion.
eg.
2.
Find the name of the cation and the anion.
eg.
3.
Na3PO4 = Na1+ , PO43-
Na1+ = sodium, PO43- = phosphate
Put the two together.
eg.
sodium phosphate
To use either method to determine the formula of a compound from the given name:
1.
Use the tables to find out the formulas for the respective ions.
eg. ammonium nitrate = NH41+ , NO31-
2.
Combine the ions in a ratio which will result in a net charge of zero.
eg. 1+, 1-; combine in a 1:1 ratio; NH4NO3
12
Naming of molecular compounds
Molecular compounds use a prefix system of naming. Prefixes are used which indicate the number of atoms of each element
present in the compound. The prefixes are as follows:
mono = 1
di = 2
tri = 3
tetra = 4
penta = 5
hexa = 6
hepta = 7
octa = 8
nona = 9
deca = 10
Thus the compound N2O4 is named dinitrogen tetraoxide. The only exception to the numbering is when the first
element in the compound is present as a single atom, such as CO2. The prefix "mono" is omitted and the compound
name is carbon dioxide. This exception is more visible in CO, or carbon monoxide. "Mono" is omitted on the
carbon, but appears on the oxygen.
To use the method for naming compounds:
1.
Determine the identity of the elements and the number of each element.
eg.
2.
Find the appropriate prefixes.
eg.
3.
N2O5 = 2 nitrogen, 5 oxygen
2 = di-, 5 = penta-
Put the name together with the appropriate prefixes.
eg.
dinitrogen pentaoxide
To use the method to determine the formula of the compound from the name:
1.
Identify the prefixes and their meaning.
eg.
2.
tetraphosphorus decaoxide = 4 P, 10 O
Put the formula together.
eg.
P4O10
Waters of hydration
Some molecules exist in nature in relationship with water. Water just seems to be attracted to the molecules and it is possible
to determine the number of water molecules relative to the number of molecules of the substance. These water molecules are
called waters of hydration. For instance CaSO 4 · 2 H2O is the formula for gypsum, used in dry wall. There are two waters
of hydration for every gypsum molecule. The proper name for this molecule is calcium sulfate dihydrate. The number of
water molecules is indicated using the numbering system from molecular compounds.
13
Organic Nomenclature
Organic chemistry refers primarily to a class of compounds called hydrocarbons. These compounds contain two
elements, hydrogen and carbon. Other elements, such as oxygen, nitrogen, sulfur, can also be included, but these
are minor constituents, if they are present at all.
The naming of organic compounds is based on three considerations:
1.
1.
Length of the carbon chain
2.
Type of bonds which join the carbons
3.
Functional groups attached to the carbon chain
Length of the carbon chain
When naming a carbon compound it is necessary to find the longest, continuous carbon chain in the
compound. Part of the name comes as a result of a prefix which indicates the number of carbons in the chain.
Following is a list of the prefixes:
2.
Number of carbons
Prefix
1
2
3
4
5
6
7
8
9
10
methethpropbutpenthexheptoctnondec-
Type of bonds which join the carbons.
There are three main groups of hydrocarbons, based on the type of bonds which join the carbons:
a)
Alkanes
- the carbon chain is joined exclusively by single bonds.
- an alkane is indicated in the name by the suffix -ane.
- the general formula for an alkane is CnH2n+2
b)
Alkenes
- the carbon chain contains at least one double bond.
- an alkene is indicated in the name by the suffix -ene.
- the general formula for an alkene is CnH2n
c)
Alkynes
- the carbon chain contains at least one triple bond.
- an alkyne is indicated in the name by the suffix -yne.
- the general formula for an alkyne is CnH2n-2
14
In the case of alkenes and alkynes, it is necessary to indicate where on the carbon chain the double or triple bond is
located. This is done by numbering the carbons on the chain, starting from the end closest to the double or triple
bond.
3.
Functional groups attached to the carbon chain
A functional group is any atom or group of atoms other than hydrogen attached to the main carbon chain.
Functional groups are identified according to their content, as well as their placement on the carbon chain. Like
alkenes and alkynes, it is necessary to indicate where on the carbon chain the functional group is located. This is
done by numbering the carbons on the chain, starting from the end closest to the functional group. In a case when
there is both a double or triple bond and a functional group, the numbering of the carbon chain is done in favor of
the bond, rather than the group.
There are three main types of functional groups it is necessary to recognize:
a)
Hydrocarbon chains
If a chain of carbons is off the main chain, the presence of the functional group is named using the prefixes
from 1. above, followed by the suffix -yl remembering also to indicated where the functional group is
located. For instance, a butane with a two-carbon chain on the number-2-carbon is called 2-ethylbutane.
b)
Halogens
If one of the halogens is attached to the carbon chain, their name is changed slightly:
Fluorine
Chlorine
Bromine
Iodine
c)
becomes
becomes
becomes
becomes
FluoroChloroBromoIodo-
Hydroxyl groups
If an hydroxyl (-OH) group is attached to the carbon chain, the hydrocarbon becomes a new class of
organic compounds, called alcohols.
The location of the hydroxyl group is indicated by carbon number at the beginning of the name, and the
extra suffix -ol is added at the end of the name. For instance, ethane with an hydroxyl or alcohol group
attached is now called ethanol.
Note: Numbering of the location of special features like multiple bonds or functional groups is not necessary in
cases where there is no question where the bond or functional group must be. This applies mainly to small
molecules of 1, 2, or 3 carbons.
15
Nomenclature Flowchart
Is the substance an element?

Yes

Name as an element
Yes

Name as a common substance
Yes

It is a molecular compound.
Use the numbering system
Yes 
It is an ionic compound.
Use Stock system

No

Is the substance a common substance? 

No

Does the substance contain non-metals only?


No

Does the substance contain a metal and a non-metal? 

No

It is likely an organic compound. Use organic nomenclature
16
Chemical Formulas and Equations
A chemical equation is a form of shorthand which gives an outline of the progress of a chemical reaction:
H2O → H2 + O2
REACTANT
PRODUCT
One very important principle of chemistry affects chemical equations. That principle is the law of conservation of
mass. Very simply, it states that matter is neither gained nor lost in a chemical reaction. In a chemical reaction you
end up with the same number and type of each atom that you started with. This applies to mass and to atoms. This
principle must be reflected in chemical equations.
Balancing equations is really very simple. It is similar to the method used to determine the chemical formulas for
ionic compounds, as discussed in Unit 3. For instance, using the equation for the decomposition of water, first tally
up the number of each kind of atom on each side of the reaction equation:
1 H2O
→
H=2
O=1
1 H2 + 1 O2
H=2
O=2
Once this is done, the method is very straightforward. The idea is to have the same number of atoms on each side of
the reaction equation. To do this, it is necessary to determine if there are fewer atoms on one side of equation,
compared with the other side. For instance, in this case the product side has one more oxygen than the reactant side.
In order to balance the equation for oxygen, the number of oxygen atoms on the reactant side must be increased; the
only way to do that is to increase the number of water molecules on the reactant side:
2 H2O
→
H=4
O=2
1 H2 + 1 O2
H=2
O=2
As you can see, increasing the number of water molecules balances the equation in terms of oxygen atoms, but now
the reactant side has too many hydrogen atoms; we must now balance the equation in terms of hydrogen. The only
way to do this is to increase the number of hydrogen atoms on the product side, by increasing the number of
hydrogen molecules:
2 H2O
H=4
O=2
→
2 H2 + 1 O2
H=4
O=2
Now the equation is balanced. The total number of each type of atom on the reactant side is equal to that on the
product side, and the law of conservation of mass is satisfied.
Now for another example we use a reaction between lead nitrate (Pb(NO 3)2) and sodium iodide (NaI), both of
which produce clear solutions. When these solution react a solid yellow compound is produced called lead iodide
(PbI2), as well as another clear solution which contains sodium nitrate (NaNO 3). The reaction equation follows:
Pb(NO3)2 + NaI → PbI2 + NaNO3
17
Step 1: Tally up the number of each atom:
Pb(NO3)2 + NaI → PbI2 + NaNO3
Pb = 1
Na = 1
I=1
N=2
O=6
Pb = 1
Na = 1
I=2
N=1
O=3
Step 2: Begin balancing for each atom:
This equation is already balanced for lead and sodium; the next on the list is iodine.
Pb(NO3)2 + 2 NaI → PbI2 + NaNO3
Pb = 1
Na = 2
I=2
N=2
O=6
Pb = 1
Na = 1
I=2
N=1
O=3
Note that the only way to increase the number of iodine atoms on the reactant side is to add another NaI molecule,
thus increasing both the number of sodium and iodine atoms. This causes the number of sodium atoms to become
unbalanced, so the number of sodium atoms on the product side must be increased, by adding another sodium nitrate
molecule:
Pb(NO3)2 + 2 NaI → PbI2 + 2 NaNO3
Pb = 1
Na = 2
I=2
N=2
O=6
Pb = 1
Na = 2
I=2
N=2
O=6
Other information can also be gained or given from a reaction. It is sometimes useful to know the state a compound
is in for a chemical reaction. By "state" we mean whether the compound is a liquid, solid, or a gas. This
information is given by the use of subscripts following each compound. The list of subscripts is as follows:
(s) (l) (g) (aq) (ppt) -
solid
liquid
gas
aqueous (meaning that the compound is dissolved in water.)
precipitate (meaning that the reaction produces a solid which falls out of solution.)
Using the two previous examples, these subscripts can be quite valuable:
2 H2O(l)  2 H2(g) + 1 O2(g)
Pb(NO3)2(aq) + 2 NaI(aq)  PbI2(ppt) + 2 NaNO3(aq)
18
Types of Chemical Reactions
There are five basic types of chemical reactions:
1)
Synthesis - the combination of two or more substances to form a compound.
A + B → AB
another way to look at it is:
element + element  compound
2)
Decomposition - One substance breaks down to form two or more simpler substances.
AB → A + B
another way to look at it is:
compound  element + element
3)
Combustion -
involves the burning of a hydrocarbon in the presence of oxygen to form carbon dioxide
and water.
CxHy + O2 → CO2 + H2O
4)
Single Replacement - reactions occur when one element is replaced by another in a compound
A + BC → AC + B
another way to look at it is:
element + compound  element + compound
5) Double Replacement -
reactions occur when the elements in a solution of reacting compounds exchange
places, or replace each other.
AB + XY →
AY + XB
another way to look at it is:
compound + compound  compound + compound
.
19
20
Chemistry 30 - Significant Digits
1.
2.
For the following count the significant digits:
a)
18.56 m
i)
1500 C
b)
0.5306 kg
j)
0.0062 L
c)
0.0128 km
k)
2.300 kPa
d)
20 apples
l)
8.0 J
e)
1.03 x 104 N
m)
15 000 000 A
f)
406.010 mol
n)
120 mm
g)
0.00920 g
o)
500 students
h)
90 502 cm
p)
100 000 t
Round off the following numbers to the given number of significant digits:
a)
6.249 mm, 2 s.d.
b)
10.98 g, 3 s.d.
c)
0.0573 mol, 2 s.d.
d)
69.95 km/h, 2 s.d.
e)
298.036 cm3, 4 s.d.
f)
349.9 A, 3 s.d.
g)
9.100 g, 2 s.d.
h)
56087250 N, 4 s.d.
i)
21.35 m, 3 s.d.
j)
450.5 kL, 3 s.d.
k)
67.77 mg, 1 s.d.
l)
2880 L, 4 s.d.
m)
675 J, 2 s.d.
20
21
3.
Perform the indicated operations. Give the answer to the correct number of
significant digits:
a) Add
c) Add
8.55 mL
11.6 mL
20.0 mL
480
km
24.07 km
d) Subtract
136
g
3.49 g
e) Subtract
16.56 mL
6.3 mL
f) Subtract
51.08 mol
9.9 mol
g)
i)
k)
4.
b) Add
9.54 g
6.578 g
10.02 g
18.4 g/mL x 5.5 mL =
1.0058 t x 1000 kg =
1t
358.6 g
2.02 mol
=
h)
21.4 g x
j)
6.0 g
24.3 g/mol
l)
2.64 g
=
5.38 g/mL
1 kg
1000 g
=
=
Perform the following calculations. Give your answers to the proper number of significant digits. If the numbers are in
scientific notation, give the answer in scientific notation.
a)
9.25 + 4.10 - 2.05 =
b)
134.8 + 2.05 - 13 =
c)
14.896 - 2.42 + 4.60 =
d)
(3.45 x 10-1) - (4.789 x 10-3) =
e)
(7.95 x 10-2) + (2.05 x 10-1) =
f)
4.18 x 0.051 960 =
g)
0.50 ÷ 4.12 =
h)
(93.30 x 10-2) x (4.612 x 101) =
i)
(1.981 x 101) ÷ (2.5 x 10-2) =
j)
((4.68 x 10-4) x (8.743 x 105)) ÷ (1.04 x 10-2) =
21
22
Chemistry 30
Using Ion Charge to Predict Formulas
Determine the most likely chemical formulas for the following combinations of ions:
a) K1+, Br1-
k) Hg2+, O2-
b) Ca2+, Cl1-
l) K1+ , PO43-
c) Li1+, H1-
m) Pb4+, O2-
d) Fe3+, OH1-
n) NH41+ , SO42-
e) Ca2+, OH1-
o) Na1+, PO43-
f) Zn2+, O2-
p) Mn7+, O2-
g) Mg2+, NO31-
q) Na1+, SeO32-
h) Fe2+, O2-
r) Na1+, SeO42-
i) Fe3+, O2-
s) Al3+, SO42-
j) Sn4+, F1-
t) H1+, S2-
22
23
Chemistry 30 - Nomenclature
Name the following substances:
1) NH3
26)
CO2
2) Fe(NO3)2
27)
SO3
3) Fe(NO3)3
28)
PbO
4) CuBr2
29)
BaS
5) SnCl4
30)
PbO2
6) CuSO3
31)
SO2
7) CuSO4
32)
BaH2
8) C2H5OH
33)
CaSO44 H2O
9) Li2CrO4
34)
O2
10) CuBr
35)
Al2O3
11) CuCr2O7
36)
PCl3
12) NH4NO3
37)
N2
13) NH4NO2
38)
Fe2O3
14) SnCO3
39)
CrCl3
15) Sn(CO3)2
40)
Pb(NO3)2
16) Na2SO4
41)
Al2(SO4)3
17) H2
42)
Cu(NO3)2
18) MgSO3
43)
(NH4)3PO4
19) KH
44)
Mg(OH)2
20) MgBr2
45)
Ca(HCO3)2
21) CO
46)
H2S
22) Li2O
47)
CaH2
23) P2O56 H2O
48)
O3
24) I2
49)
XeF8
25) CH4
50)
S8
23
24
Write molecular formulas for each of the following:
1) sodium fluoride
26) calcium carbonate
2) potassium carbonate
27) ammonium sulfite
3) aluminum sulfide
28) iron (II) hydroxide
4) calcium bromide
29) copper (II) nitrate
5) aluminum chloride
30) oxygen
6) silver oxide
31) lithium dichromate
7) ammonium sulfide
32) hydrogen nitrate
8) barium hydroxide
33) barium bicarbonate
9) iron (III) iodide
34) nitrogen dioxide
10) sodium sulfate
35) carbon monoxide
11) tin (II) nitrate
36) ammonium carbonate
12) potassium sulfite
37) ammonium dichromate
13) magnesium sulfate
38) aluminum hydroxide
14) lithium phosphate
39) tin (II) nitrate
15) bromine
40) tin (IV) carbonate
16) plumbous oxide
41) mercurous chloride
17) hydrogen
42) mercuric chloride
18) sulfur dioxide
43) zinc perchlorate
19) methanol
44) chromium (III) stearate
20) carbon dioxide
45) chromium (II) acetate
21) sodium chromate
46) barium nitrite dihydrate
22) sulfuric acid
47) iron (II) phosphate
23) sodium hydride
48) iron (III) chlorate
24) stannous chloride
49) cuprous bromide
25) hydrogen selenide
50) ammonium oxalate
24