Science 10
Unit 1
Chapter 3
Chemical Reactions
pp 82-137
1
A Chemical Engineer’s Recipe for Chocolate Chip Cookies
Objective: Household cooking involves chemistry. Here is a chemical experiment in which the product is edible.
Materials:
• 236.6 cm3 (500 mL) partially hydrogenated tallow triglyceride (butter or margarine) 177.45 cm3 (177
mL) crystalline sucrose (sugar) . 177.45 cm3 (177 mL) unrefined sucrose (brown sugar) 4.9 cm3 (5 mL) 4hydroxy- 3-methoxybenzaldehyde (vanilla) • 2 calcium carbonate-encapsulated avian albumen-coated
protein (eggs) • 532.35 cm3 (500 mL) gluten (flour) 4.9 cm3 (5 mL) sodium chloride (salt) • 4.9 cm3 (5
mL) sodium bicarbonate (baking soda) • 236.6 cm3 (250 mL) chopped de-encapsulated legume meats
(nuts) . 1 package of theobroma cacao (chocolate chips) • measuring spoons • measuring cups mixing
bowls • cookie sheets
Procedure:
1. Cream the partially hydrogenated tallow triglyceride, 4-hydroxy-3-methoxybenzaldehyde, crystalline
and unrefined sucrose in a bowl.
2. Add the calcium carbonate-encapsulated avian albumen-coated proteins and mix well.
3. In a separate bowl combine the gluten, sodium chloride, and sodium bicarbonate. Add to creamed
mixture.
4. Stir in theobroma cacao and de-encapsulated legume meats.
5. Place the final mixture piecemeal onto a cookie sheet.
6. Heat in a 463 K (190°C/ 375°F) oven for 8-10 minutes and allow the chemical reactions to take place.
7. Remove from oven and place on cooling rack.
Shopping Cart Solutions
The best way to tackle the safe disposal of
hazardous chemicals is to avoid buying them.
There are safe and effective alternatives to most
hazardous household chemicals. For example,
most of your household cleaning needs can be
met with six simple ingredients: soap, baking
soda, washing soda, vinegar, lemon juice, and
ammonia. Various combinations of these simple
substances can accomplish most household
cleaning jobs cheaply and safely.
Homemade Cleaners
As with commercial products, these solutions
should be stored in carefully labelled, wellclosed containers out of the reach of children.
Save your old cleaner containers for your
homemade solutions.
ALL-PURPOSE CLEANER
In a 4L container, pour
120 mL vinegar
60 mL baking soda
Fill container with hot water, label, and use for
most household cleaning jobs. For a stronger
cleaner or wax stripper, double the amount of
vinegar and baking soda, but not water. Do not
mix with other compounds, especially chlorine
bleach.
WINDOW CLEANER
Mix 60 mL of vinegar with 250 mL of water.
Label and use in a spray pump bottle.
SCOURING POWDER
Use baking soda and a damp cloth. For stubborn
stains, make a paste of baking soda and water
and leave on for a few minutes.
TOILET BOWL CLEANER
Baking soda is also an excellent toilet bowl
cleaner. For tough jobs, a paste made from borax
and lemon juice can be used to clean the entire
ring. Swish the toilet brush around to get the
bowl wet. Rub on the paste and leave it in place
for about two hours, scrub with brush, and flush.
Please remember to use borax sparingly.
TUB AND TILE CLEANER
Use baking soda and a damp cloth for routine
cleaning. For soap film, a paste of white vinegar
and washing soda works well.
OVEN CLEANER
Try to clean spills as they occur, sprinkling with
salt while still hot. The spill can be wiped off
with a sponge when cool. For baked-on spills,
sprinkle oven when cool with washing soda and
leave for about 15 ruin, then scrub with
dampened steel wool.
2
Chapter 3: Chemical Reactions pp 82-137
3.1
Recognizing and Describing Chemical Reactions p 84
A chemical reaction, or chemical change, has occurred
when ______________________________________
__________________________________________.
 The reactants : substances that _______________
_______________________________________
 The products are the _______________________
_______________________________________
Evidence of a Chemical Reaction: (see table 3.1 p84)
 _______________________________
 _______________________________
 _______________________________
 _______________________________
 Energy change such as:
 __________________________
__________________________
 __________________________
__________________________
*energy changes can also accompany physical changes, which
are simply changes in state
Evidence of a Physical Change:
 _______________________________________
 _______________________________________
 _______________________________________
Chemical Reaction Demos: “The Beverage Princess” & “Carbon Tower”
Do BLM 3-1 “Chemical or Physical Change”
Do Lab- Evidence of Chemical Reactions
3
BLM 3.1
Chemical or Physical Change?
Goal  Demonstrate your ability to distinguish between chemical and physical changes.
What to Do
Classify each change in the following table as either chemical or physical. Explain your
classification.
Change
Chemical or physical?
Explanation
A sheet of paper is crumpled into a
ball.
A sheet of paper is set on fire and
burns to ashes.
Steel wool is placed in a glass of
salty water. The steel wool rusts.
A sheet of flexible, colourless
plastic is left outside, in bright
sunlight, and becomes yellow and
brittle over time.
A teaspoon of white sugar (sucrose)
dissolves in a glass of warm water.
Vinegar is poured over a teaspoon
of baking soda. The white powdered
baking soda fizzes, and bubbles
form.
A red-hot nail is inserted into a large
block of ice. Steam is formed as the
nail contacts the block, and water
flows away from the nail.
4
Solubility and Chemical Reactions
 when ionic compounds are dissolved in water, the cations
and anions __________________________________
 some ionic compounds ___________ dissolve well in water.
 if two ionic compounds are mixed, their free __________
____________________________________________
 if the new compound does not dissolve in water, a ____________ is
formed.
Predicting Solubility
To determine whether a compound is soluble in water or not, follow these
steps, using the solubility chart on your periodic table:
1.
_______________________________________________
_______________________________________________
2.
_______________________________________________
______________________.
3.
if the 2nd ion is found in the “high solubility” category, the
compound is soluble in water and should be given a __________
________________________.
OR
If the 2nd ion is found in the “low solubility” category, the
compound is NOT soluble in water and should be given the
__________________________.
It will form a_______________ in water.
Ions from Alum
 Alum is a _______________________ commonly added to water
when being treated.
 Alum removes ions of ____________________ found in raw,
untreated water by forming precipitates of these elements.
 In solution, alum forms ______, _______ and ________ ions.
Which of the following cations cannot be removed from the water by alum,
and why?
Ba2+
Pb2+
Ag+
Li+
5
Do Practice Problems 1-4 p90 and BLM 3-2 “Using a solubility Table”
Precipitate Demo:
KI(aq) & Pb(NO3)2(aq)
Do Investigation 3-A, p.86. (Use BLM 3-3 to record your
observations.)
Do questions 1-6, p.87.
STATES OF MATTER AND SOLUBILITY TABLE
Solubility —
1. When a compound dissolves in water it is said to be ________________________.
-we use (aq) to describe soluble compounds dissolved in water
2. When a compound does not dissolve in water it is said to be ___________________.
-we use (s) to describe the solid compound that does not dissolve in water
-if the solid compound is a product of a chemical reaction, we call it ______________.
We can determine the solubility of a compound by looking it up on a solubility table (see
the bottom of your periodic table or p. 466 in text).
To use the solubility table:
1. look for the anion (the negatively charged ion) across the top of the table
2. next, locate the cation (positively charged ion) in either the “soluble” or “insoluble”
rows
-if the cation is in the “soluble” row, write the empirical formula with (aq)
-if the cation is in the “insoluble” row, write the empirical formula with (s)
Rules:
1. All acids are soluble in water.
2. Molecular compounds can be either (s), (1) or (g).
-if there is only 1 element, use the periodic table
-for compounds with more than one element, use common sense
ie. CH4(g), H2O(l), CO2(g)
3. Ionic compounds are usually solids. However, if the compound is reacting, you would
have to add water
-soluble ionic compounds are (aq)
-insoluble ionic compounds are (s)
6
Solubility of Ionic Compounds in Water
Ion
BrICl-
S
High
solubility
Most
Group 1
Group2
Low
solubility
2-
Ag+
Pb2+
Cu+
Hg+
Most
OH
SO42-
CO32PO43SO32-
CH3COO-
NO3-
NH4+
Group 1
Sr2+
Ba2+
Most
Group 1
Most
All
All
Most
Ag+
Pb2+
Ca2+
Ba2+
Sr2+
Ra2+
Most
Ag+
2-
-
-
Group IA
ION
(aq)
SOLUBLE
s ≥ 0.1 mol/L
alkalis
all
H+
all
+
NH4
all
ClO3ClO4- CH3COONO3-
all
all
ClIBr-
SO42-
S
PO43SO32CO32-
OH
Group
Group
IA NH4+
IA &
most most
Sr2+
IIA
Ba2+
NH4+
(s)
LOW
SOLUBILITY
s < 0.1 mol/L
none
none none none
none
Ag+
Pb2+
Hg+
Cu+
Ag+
Pb2+
Ca2+
Ba2+
Sr2+
most
most
Group
IA
NH4+
most
7
BLM 3-2
Using a Solubility Table
SKILL BUILDER
1. Complete the following table.
Name
sodium chloride
Formula
NaCl
Cation
Anion
Na+
Cl–
High or low
solubility?
lithium iodide
Mg(ClO3)2
strontium hydroxide
BaCO3
2. Complete the following table.
Name
Formula
High or low
solubility?
Al(OH)3
ammonium chloride
K2S
molybdenum(V) chlorate
Pb(CH3COO)2
copper(II) iodide
FeCO3
calcium sulfite
Ba3(PO4)2 (s)
palladium(II) bromide
HgI
strontium sulfate
8
Chemical Reactions and Energy Changes
Energy changes accompany any chemical reaction.
Endothermic –____________________________________
Eg. _______________________
Exothermic – ___________________________
__________________________
Eg. ___________________
Do BLM 3-4
Law of Conservation of Energy
This law states that energy can be converted from one form into another,
but the
total energy of the universe remains constant.
Energy cannot be created or destroyed.
Energy is ______________ to break chemical bonds, and energy is
_______________ when chemical bonds are formed.
Endothermic reaction – __________________________________
____________________________________________________
____________________________________________________
Exothermic reaction – __________________________________
____________________________________________________
____________________________________________________
Do Check Your Understanding #1,3,4,5 p93
9
BLM 3-4
REINFORCEMENT
Endothermic and Exothermic Reactions
Goal  Demonstrate your understanding of endothermic and exothermic reactions.
What to Do
Read the following summary of exothermic and endothermic reactions. Then answer the
questions on the next page.
Summary
–
–
–
–
–
–
–
–
A chemical change is always accompanied by a change in energy because the atoms or ions
that make up the reactants are rearranged.
During a reaction, chemical bonds that hold the reactant atoms or ions together must be
broken and new chemical bonds must be formed within the product substance(s).
The breaking of chemical bonds requires the input of energy and is defined as an endothermic
process.
The formation of chemical bonds releases energy and is defined as an exothermic process.
Since any chemical change involves both the breaking and formation of chemical bonds,
there are two possible outcomes of any chemical change.
The overall change is exothermic if more energy is released to form the product chemical
bonds than is required to break the reactant chemical bonds.
The overall change is endothermic if less energy is released to form the product chemical
bonds than is required to break the reactant chemical bonds.
The energy that is released or absorbed is related to the external environment of the chemical
reaction. Endothermic reactions absorb thermal energy from the surrounding environment
and result in a decrease in temperature. Exothermic reactions, on the other hand, release
thermal energy to the surrounding environment and result in an increase in temperature.
Type of
reaction
exothermic
endothermic
Breaking chemical
bonds (reactants)
Forming chemical
bonds (products)
Overall energy
change
energy released
energy absorbed
10
1. Classify each reaction as either endothermic or exothermic, and briefly explain your answer.
Description of chemical reaction
Endothermic or
exothermic?
Explanation
A piece of paper is ignited and
burns with a bright flame.
Pentaborane (a colourless liquid),
B5H9, reacts violently with oxygen
gas to form solid diborane, B2O3,
and water, typically bursting into
flame and often exploding.
Pure iron metal is formed and
carbon dioxide is released when
iron(III) oxide ore is heated to a
very high temperature in the
presence of solid carbon.
Sodium hydroxide solution and
hydrochloric acid solution are
mixed. The temperature of the
mixture increases.
Mixing ammonium thiocyanate
and barium hydroxide octahydrate
in a beaker causes water on the
outside of the beaker to freeze.
The high temperature in an oven
causes baking soda (sodium
hydrogencarbonate) to break down
into carbon dioxide, water, and
sodium carbonate.
2. (a) What is the source of the energy that is released in an exothermic reaction? What absorbs
the energy that is
(b) What is the source of the energy that is absorbed in an endothermic reaction?
3. A student claims that the reaction of butane gas and oxygen gas must be endothermic since a
spark is needed to ignite the butane gas in a lighter. Do you agree or disagree with this claim?
Explain your answer.
11
rele
3.2
Representing Chemical Reactions p94
Closed system – __________________________________
______________________________________________
Example – ________________
Open system – ___________________________________
______________________________________________
Example – ___________________
Isolated system – ________________________________________
______________________________________________________
Example (theoretical) – ______________
Law of Conservation of Mass
Developed by Antoine Lavoisier (1743-1794)
This law states that during a chemical reaction, the total mass of the
reacting substances (reactants) is always equal to the total mass of the
resulting substances (products).
Conservation of Mass Demo: Alka Seltzer
Do Investigation 3-B, p.95 “Comparing Masses of Reactants and
Products” and p. 96 # 1-10
12
BALANCING CHEMICAL EQUATIONS
Purpose:
1. To introduce a theory of how chemical reactions occur.
2. To illustrate the conservation of atoms in balancing the following chemical reactions.
3. To illustrate the meaning of the coefficients vs. the subscripts in a balanced equation.
Prelab Exercise:
Count the total number of each kind of reactant and product atoms and record the
number in the space provided.
Observations
1. Rocket fuel may be produced from the decomposition of water.
2H2O(l) ----------------- > H2(g) + O2(g)
2. Rocket fuel is burned in a Saturn rocket.
2H2(g) + O2(g) ----------------- > 2H2O(g)
#H =
#O =
#H =
#O =
3. Natural gas (mostly methane) is burned as a heating fuel.
CH4(g) + 2O2(g) ----------------- > CO2(g) + 2H2O(g)
#C =
#H =
#O =
#C =
#H =
#O =
4. Hydrogen chloride gas is produced for the production of hydrochloric acid.
H2(g) + Cl2(g) ----------------- > 2HCl(g)
#H =
#Cl =
#H =
#Cl =
5. Ammonia for fertilizers is produced from nitrogen and hydrogen.
N2(g) + 3H2(g) ----------------- > 2NH3(g)
#N =
#H =
#N =
#H =
6. Ammonia dissolves in water to form some ammonium hydroxide.
NH3(g) + H2O(l) ----------------- > NH4+(aq) + OH-(aq)
#N =
#H =
#O =
#N =
#H =
#O =
13
7. In a lab hydrogen peroxide was decomposed.
2H2O2(l) ----------------- > H2O(l) + O2(g)
#H =
#O =
#H =
#O =
Questions:
1. Write out in words what the following equation represents.
N2(g) + 3H2(g) ----------------- > 2NH3(g)
What is wrong with the following answers?
1. 2NH3(g)
2. 2O2(g)
Writing Balanced Chemical Equations
Reaction Description:
________________________________________
Word equation:
_________________________
Skeleton equation:
________________________
Balanced equation:
_________________________
*shows atoms are ____________________.
*_______________ (numbers in front of chemical formulas) show how many of
each compound is there . (Note: there is 1 molecule of oxygen gas, but 2 atoms of
oxygen.)
Q: How many atoms of hydrogen are reacting?
Oxygen?
*_____________ (letters in brackets to the right of each compound) show what
state each compound is in.
14
How to balance equations:
1. Include all ___________ of matter for each reactant and product.
2. Balance the atom or polyatomic ion present in the _________________.
Find the lowest common multiple to obtain whole number coefficients to
balance.
3. ________________ to balance each remaining atom/ion.
HINTS:




O2(g) – when O2(g) is present in the reaction, balance it ____
When ________________ are present on one side of the
equations, balance those elements last (ie. Na, Mg)
As long as complex ions are found intact on both sides of
the equation, _________________________________
Check once again to make sure the equation is balanced.
Study Model Problem 1-3 p.99-100
Do Practice Problems 5-8 p.101
Do Check Your Understanding #2-5 p.102
Do Worksheet “Chemical Reactions Balancing Equations”
15
3.3
1.
Types of Chemical Reactions
p.103
_______________________:
Eg. ________________________________
See figure 3.15 and 3.16 p104.
2.
_______________________:
Eg. ____________________
See figure 3.17 and 3.18 p 104.
3.
_______________________:
Eg.1. ____________________
Eg.2. ____________________
4.
_______________________:
Eg.
___________________
*many of these equations result in a precipitate forming
16
5.
Reactions involving Carbon Compounds:
The study of carbon-containing compounds is called __________
chemistry. A ____________ is a compound that contains only
hydrogen and carbon atoms (eg. C2H6). (These compounds are retrieved
by refining crude oil and natural gas. Approximately 95% of these
hydrocarbons are burned as fuels in exothermic combustion reactions
to create thermal energy to warm buildings and provide energy for
transportation.)
eg.
Complete Combustion: a hydrocarbon reacts with
oxygen gas (or burns) to create _______________
______________________________________.
See figure 3.20 and 3.21 p.108.
Incomplete Combustion: when _______ is in poor
supply, the products are carbon dioxide gas, water
vapor, ________________________________________________
than a complete combustion reaction. (*carbon monoxide is a colorless,
odorless, highly toxic gas that, when breathed in, strongly binds to red blood
cells instead of oxygen and can lead to death.)
6. Other – anything that doesn’t fit into one of the above reaction types.
You will not be expected to determine the products of an “other” reaction.
Study the Tools of Science p.109 and learn the chemical tests for:
Hydrogen –
Oxygen –
Carbon dioxide –
Water Do Investigation 3-D Classifying Chemical Reactions p.110
Do Check Your Understanding #1-9 p114
17
Practise: Types of Chemical Reactions
1. Simple Composition:
2. Simple Decomposition:
3. Single Replacement:
4. Double Replacement:
5. Hydrocarbon Combustion:
6. Other:
18
Investigation 3-D
Classifying Chemical Reactions
Reactants
Part 1:
Demo
Mg(s) + O2(g)
Part 2:
Mg(s) + HCl(aq)
Part 3:
Demo
C20H42(s) + O2(g)
Part 4:
H2O2(l) + MnO2(s)
Note: manganese is a catalyst
Part 5:
CuCl2(aq) + Al(s)
Part 6:
H2SO4(aq) + NaOH(aq) +
phenolphthalein
Part 7:
NaOH(aq) + (NH4) 2SO4(aq)
Part 8:
AgNO3(aq) + NaCl(aq)
Part 9:
CuSO4(aq) + Zn(s)
ZnSO4(aq) + Cu(s)
Part 10:
HCl(aq) + CaCO3(s)
marble chips
Balanced Reaction
p.110
Reaction
Type
H2O2(l)  O2(g) + H2O(l)
NaOH(aq) + (NH4) 2SO4(aq) 
NH4OH(aq) + Na2SO4(aq)
(NH4OH (aq)  NH3(g) +
H2O(l)
HCl(aq) + CaCO3(s) 
CaCl2(aq) + H2CO3(aq)
H2CO3(aq)  H2O(l) + CO2(g)
Note: Go over burning splint test for H2 and glowing splint test for O2 before lab.
19
Science Focus 10 Unit 1 Chapter 3
Chemical Reactions and Chemical Equations
Evidence for Chemical Reactions
Some of the easy-to-observe clues for recognizing a chemical reaction are:
1. formation of a precipitate; a solid forms when two solutions are added to one another
2. formation of a gas; bubbles in a solution
3. color change
4. energy change; heat, light, electricity is given off or taken in.
Laws of Conservation
In all chemical reactions, the following are conserved:
1. The number of each kind of atom
2. The mass of products will equal the mass of reactants
3. Energy; energy given off by a reaction comes from chemical bonds and energy taken in by a
reaction forms new chemical bonds
Information Conveyed by a Balanced Chemical Equation
A balanced chemical equation gives the following information:
1. the chemical composition of reactants and products
2. the phase of the substance involved
3. the mole relationship of the substances involved
4. the reaction type
Classification of Chemical Equations According to the Reaction Type
The products of a chemical reaction can be predicted if the reactants are known according
to the following types of chemical reactions.
1. formation: element + element ---> compound
H2(g) + O2(g) ---> H2O(g)
2. decomposition: compound --- > element + element
NaCl(aq) ---> Na(aq) + Cl2(aq)
3. single replacement: element + compound ---> element + compound
Na(s) +
LiCl(s) ---> Li(s)
+ NaCl(s)
4. double replacement compound + compound ---> compound + compound
KOH(aq) + NaCl(aq) ---> KCl(aq) + NaOH(aq)
5. hydrcarbon combustion: hydrocarbon + oxygen -> carbon dioxide + water
C3H8(g) + O2(g) -> CO2(g) + H2O(g)
6. other:
Any reaction that does not fit into one of these categories cannot be predicted.
20
Balancing Reactions and Predicting Reactions
1. Classify the following reactions as to their type; sc (simple composition or
formation), sd (simple decomposition), sr (single replacement), dr(double
replacement),hc (hydrocarbon combustion), o (other).
2. Balance the chemical equations.
REACTION
TYPE
1. ___Al(s) + ___O2(g) -> ___Al2O3(s)
______________
2. ___HCl(aq) + ___Ca(OH)2(aq) -> ___HOH(l) + ___CaCl2(aq)
__________
3. ___CH4(g) + ___O2(g) -> ___CO2(g) + ___H2O(g)
__________
4. ___Zn(s) + ___Pb(CH3COO)2(aq) -> ___Pb(s) + ___ Zn(CH3COO)2(aq) __________
5. ___SO3(g) + ___H2O(l) -> ___H2SO4(aq)
__________
6. ___HgO(s) -> ___Hg(l) + ___O2(g)
__________
7. ___CaCO3(s) -> ___CaO(s) + ___CO2(g)
__________
8.
___NaI(aq) + ___Pb(NO3)2(aq) -> ___PbI2(s) + ___NaNO3(aq)
_______________
9. ___Cl2(aq) + ___NaI(aq) --> ___I2(aq) + ___NaCl(aq)
__________
10. ___Al2(SO4)3(aq) + ___Ca(OH)2(aq) -> ___Al(OH)3(s) + ___CaSO4(s)
__________
11. ___Al2(SO4)3(aq) + ___Ca(HCO3)2(aq)-> ___A1(OH)3(s) + ___CaSO4(S) + ___CO2(g)
_______________
12. ___C8Hl8(1) + ___O2(g) -> ___CO2(g) + ___H2O(g)
__________
13. ___H2O(1) -> ____H2(g) + ___O2(g)
_______________
14. ___Na(s) + ___Cl2(g) -> ___NaCl(s)
__________
15. ___Ca(s) + ___HOH(l) -> ___Ca(OH)2(s) + ___H2(g)
_______________
Balancing Formation Reaction Equations
1. Balance each equation for a formation reaction.
(a) K(s)
+
O2(g)

K2O(s)
(b)
P4(s)
+
Cl2(g)

PCl5(s)
(c)
Cu(s)
+
S8(s)

CuS(s)
(d)
Mg(s)
+
O2(g)

MgO(s)
(e)
Fe(s)
+
O2(g)

Fe2O3(s)
(f)
P4(s)
+
S8(g)

P2S5(s)
(g)
C(s)
+
O2(g)

CO(g)
21
(h)
N2(g)
+
O2(g)

NO2(g)
(i)
Li(s)
+
N2(g)

Li3N(s)
(j)
S8(s)
+
O2(g)

SO2(g)
2. Write a balanced chemical equation to represent each reaction described below.
(a) Solid aluminum metal reacts with oxygen gas to form solid aluminum oxide.
_____________________________________________________________
_____________________________________________________________
(b) Metallic zinc combines with elemental sulfur to form zinc sulfide.
_____________________________________________________________
_____________________________________________________________
Balancing Decomposition Reaction Equations
3. Balance each equation for a decomposition reaction.
(a)
NaCl(s)

Na(s)
+
Cl2(g)
(b)
CaBr2(s)

Ca(s)
+
Br2(l)
(c)
CCl4(l)

C(s)
+
Cl2(g)
(d)
NCl3 (g)

N2(g)
+
Cl2(g)
(e)
P4O10(s)

P4(s)
+
O2(g)
(f)
Ag2O(s)

Ag(s)
+
O2(g)
(g)
HCl(aq)

H2(g)
+
Cl2(g)
(h)
KI(s)

K(s)
+
I2(s)
(i)
AlCl3(s)

Al(s)
+
Cl2(g)
(j)
CuO(s)

Cu(s)
+
O2(g)
4. Write a balanced chemical equation to represent each reaction described below.
(a) Rubidium oxide decomposes into its elements.
(b) Calcium chloride decomposes into its elements.
22
Balancing Single Replacement Reaction Equations
5. Balance each equation for a single replacement reaction.
(a)
K(s)
+
H3PO4(aq)

K3PO4(aq)
+
H2(g)
(b)
Fe(s)
+
H2S(aq)

Fe2S3(s)
+
H2(g)
(c)
Cl2(g)
+
MgBr2(aq)

MgCl2(aq)
+
Br2(aq)
(d)
Cu(s)
+
Ag2CO3(s)

CuCO3(s)
+
Ag(s)
(e)
Br2(g)
+
KI(aq)

I2(aq)
+
KBr(aq)
(f)
Mg(s)
+
Zn3(PO4)2(s)

Mg3(PO4)2(s)
+
Zn(s)
(g)
K(s)
+
Al(NO3)3(aq)

Al(s)
+
KNO3(aq)
(h)
Ca(s)
+
H2O(l)

Ca(OH)2(s)
+
H2(g)
(i)
Na(s)
+
H2SO4(s)

Na2SO4(aq)
+
H2(g)
(j)
K(s)
+
H2O(l)

KOH(aq)
+
H2(g)
6. Write a balanced chemical equation to represent each reaction described below.
(a) Silver reacts with gold(III) nitrate.
(b) Copper reacts with lead(II) sulfate.
Balancing Double Replacement Reaction Equations
7. Balance each equation for a double replacement reaction.
(a)
Na2SO4(aq)
+
BaCl2(aq)

BaSO4(s)
+
NaCl(aq)
(b)
HNO3(aq)
+
Ba(OH)2(aq)

H2O(l)
+
Ba(NO3)2(aq)
(c)
Na2CO3(aq)
+
Fe(NO3)3(aq)

Fe2(CO3)3(s)
+
NaNO3(aq)
(d)
CaCl2(aq)
+
K3PO4(aq)

Ca3(PO4)2(s)
+
KCl(aq)
(e)
Al2(SO4)3(aq)
+
Ba(OH)2(aq)

Al(OH)3(s)
+
BaSO4(s)
(f)
NaOH(aq)
+
H2SO4(aq)

H2O(l)
+
Na2SO4(aq)
(g)
Na3PO4(aq)
+
Ag2SO4 (s)

Na2SO4(aq)
+
Ag3PO4(s)
23
(h)
Na2CrO4(aq)
+
Cu(NO3)2(aq)

NaNO3(aq)
+
CuCrO4(aq)
(i)
H3PO4(aq)
+
KOH(aq)

H2O(l)
+
K3PO4(aq)
(j)
Na2CO3(aq)
+
HNO3(aq)

H2CO3(aq)
+
NaNO3(aq)
8. Write a balanced chemical equation to represent each reaction described below.
(a) Solutions of sodium hydroxide and hydrochloric acid react.
(b) A silver nitrate solution reacts with a sodium chloride solution.
Balancing Combustion Reaction Equations
9. Balance each equation for a combustion reaction.
+
O2(g)

CO2(g)
+
H2O(g)
(b) C3H8(g)
+
O2(g)

CO2(g)
+
H2O(g)
(c) C6H14(g)
+
O2(g)

CO2(g)
+
H2O(g)
(d) C8H18(g)
+
O2(g)

CO2(g)
+
H2O(g)
(e) C2H2(g)
+
O2(g)

CO2(g)
+
H2O(g)
(f)
+
O2(g)

CO2(g)
+
H2O(g)
+
O2(g)

ZnO(s)
+
SO2(g)
(h) CH3NO2(l) +
O2(g)

CO2(g)
+ H2O(g) + NO2(g)
(i)
NH3(g)
+
O2(g)

NO2(g)
+
(j)
C2H5SH(g) +
O2(g)

CO2(g)
+ H2O(g) + SO2(g)
(a)
C2H6(g)
C2H4(g)
(g) ZnS(s)
H2O(g)
10. Write a balanced chemical equation to represent each reaction described below.
(a) Candle wax, C25H52, is burned to produce carbon dioxide and water.
(b) Sucrose, C12H22O11, is burned to produce carbon dioxide and water.
24
Classifying and Balancing Equations
11. Classify each reaction as a formation (F), decomposition (D), single replacement (SR),
double replacement (DR), or combustion (C) reaction. Then balance each equation.
Reaction
Classification
Li(s)
+
AlCl3(aq)

Al (s)
NH3(g)

N2(g)
+
H2(g)
K(s)
+
Br2(l)

KBr(s)
C10H22(l)
+
O2(g)

H2CO3(aq)
NH4OH (aq) +
+
LiCl(aq)
CO2(g)
+
H2O(g)

H2O(l)
+
(NH4)2CO3(aq)
+
ZnF2(aq)
H2O(l)

H2(g)
+
O2(g)
Al(s)
+
Cl2(g)

AlCl3(s)
Zn(s)
+
SnF4(aq)

Sn(s)
Completing and Balancing Reactions
For each of the following state the type of reaction and write the balanced equation. Give
the phase of each substance in the even numbered questions.
1. sodium chlorate + iron(III) dihydrogen phosphate
2. antimony(III) bisulfite + iron(II) glutamate
3. sodium burns in oxygen
4. copper(I) bromate + antimony(V) iodate
5. ethane burns (C6H6)
6. iron(II) bicarbonate + sodium chloride
7. silver carbonate + strontium
25
8. magnesium + phosphorus
9. aluminum borate + chromium(II) dihydrogen phosphate
10. aluminum sulfide decomposes
11. hydrochloric acid + calcium hydroxide
12. lithium astatide + cobalt(II) nitride
13. tin(IV) fluoride + beryllium permanganate
14. octane, C8H18 burns
15. iron(II) oxide + potassium permanganate
15. phosphoric acid + magnesium hydroxide
17. chlorine + aluminum iodide
18. stearic acid + strontium hydrogen carbonate
19. calcium + copper(I) sulfite
20. nickel(II) tetraborate + potassium dihydrogen phosphate
21. iron(II) nitrite + antimony(V) dichromate
22. tin(IV) sulfide + chromium(II) carbonate
23. silver hydrogen sulfite + calcium thiosulfate
26
THE MOLE
THE MOLE, MOLAR MASS AND MASS - MOLE RELATIONSHIPS
The Mole
The mole was introduced in Unit C as a convenient number of atoms, ions or molecules to work with in the
laboratory. This convenient number (Avogadro's number, 6.02 x 10 23) also has significance in terms of the atomic
mass of elements. The mole is defined as the number of atoms in exactly 12 g of a particular isotope of carbon. The
particular 1 carbon isotope is the most common isotope of carbon -- the carbon-12 isotope with 6 protons and 6
neutrons.
The Green Pea Analogy
If you selected a hundred (102) average-sized peas, you would find that they occupy roughly a volume of 20 cm3. A
million (106) peas are just enough to fill an ordinary household refrigerator and a billion (10 9) peas will fill a three
bedroom house from cellar to attic. A trillion (10 12) peas will fill a thousand houses, the number you might find in a
medium-sized town. A quadrillion (1015) peas will fill all the buildings in one of our larger cities such as Calgary or
Edmonton.
Obviously you will run out of buildings very soon. Let us try a larger measure, for instance the province of Alberta
Suppose that there is a blizzard over Alberta, but instead of snowing snow, it snows peas. Alberta is covered with a
blanket of peas about one metre deep all the way from Saskatchewan out to British Columbia and all the way from
the United States to the Northwest Territories. This blanket of peas drifts over the roads and banks up against the
sides of the houses, and covers all the fields and forests. Think of flying across the province with the blanket of peas
extending out as far as you can see. This gives you an idea of our next number. There will be in this blanket about a
quintillion (1018) peas.
Imagine that this blizzard of peas falls over the entire land of the globe - North America, Africa, South America,
Europe, Australia, and Asia. All of the continents are covered with peas one metre deep. This global blanket will
contain sextillion (1021) peas. Then imagine that the oceans are frozen over and the blanket of peas covers the entire
land and sea area of Earth. Go out among the neighboring stars and collect 250 planets the size of Earth and cover
each of these with a blanket of peas one metre deep. Then you have a mole of peas.
Furthermore, go out into the farthest reaches of the Milky Way, and collect 250 000 planets, each the size of Earth
Cover each one with a blanket of peas one metre deep. You now have cotillion (10 27) - a number corresponding to
the number of atoms in your body.
Molar Mass
One mole is defined as the number of atoms of carbon-12 in exactly twelve grams. The mass of one mole of all other
atoms is determined relative to the mass of one mole of carbon-12. The average mass of one mole of atoms of an
element is given to the nearest hundredth of a gram on the ALCHEM periodic table. For example the mass of one
mole (molar mass) of chlorine atoms is 35.45 g/mol. This molar mass in ;in average value which takes into account
that a sample of chlorine is composed of several isotopes of chlorine.
The molar mass (mass of one mole) of compounds may be determined from the molar masses of their component
atomic elements. Examples of how to determine these molar masses (always in grams per mole, g/mol) are provided
on the pages to follow.
Molar mass is a general term which may refer to the mass of one mole of atoms, molecules, formula units, etc. In
order to avoid confusion the term atomic molar mass should be used to refer to the mass of one mole of atoms
(versus molecules or formula units).
27
3.4 The Mole p116
Since atoms and many compounds are very tiny, scientists group them into an
extremely large number called a mole.
The mole:
 One mole of substance that contains _________________ particles in it.
 In 12 g of carbon-12 (the most common isotope of
carbon) there are 6.02 x 1023 atoms.
 6.02 x 1023 = Avogadro’s number
The unit mol is short for the German word “Molekulargewicht”, which is
literally translated to, you guessed it, molecular weight!
In Equations:
The coefficients make the molar ratio
ex. 2 H2O(l)  2 H2(g) + 1 O2(g)
…means that ___ moles of water makes ___ moles of hydrogen and ___ mole of
oxygen
The molar ratio is ___: ___: ___
Molar Mass (M)
 ____________________________________________________
______________________________
 this is a …
 eg.
Molar Mass



of a Compound (M)
___________________________________
units are in ______________
eg. CO2
Do p.120 #9 and 10
28
Formula for Moles (n)
n = _m_
M
___ = number of moles of substance
___ = mass of substance ( )
___ = molar mass of substance ( )
1. Find the number of moles in 25.0 g of CuSO4:
2. Find the mass of 2.5 mol of MnCl4.
3. How many moles in 50.0 g of sodium carbonate?
4. How much would 0.0500 moles of sodium carbonate weigh?
The Mole and the Law of Conservation of Mass
Study the chemical reaction at the top of p124.
The coefficients explain the ___________________________________
_________________________________________________________.
Notice how molecules, moles and mass are all related in the chemical equation.
- the ________________________________________
____________________________________________.
Do
Do
Do
Do
Do
Practice Problems p.122 and 123
p.122 #11-22
Check Your Understanding p.125 #2-6
Chapter 3 Review p.128 #1-5,7,9,10,12(a-b),13(a-c),14-19
Unit 1 Review p.134 #1-38,39-43(a-c for each),46-48
29
THE MOLE
SIGNIFICANT DIGITS
There is some degree of uncertainty in every measurement. When scientists record and communicate data
it is important that the degree of uncertainty be shown. One method of indicating uncertainty is by the
number of significant digits recorded.
Definition
Significant digits are those digits obtained from a properly taken measurement. Significant digits as
obtained from a measurement are all of the certain digits from a measurement plus one uncertain
(estimated) digit. Generalized to all situations (i.e., values from measurement or calculation) significant
digits are those digits which are certain plus one uncertain digit. Only significant digits are reported.
Counting Significant Digits
1. Count all digits from 1 to 9 plus zeros in between and following other digits.
2. Do not count zeros in front of a value because they only serve to set the decimal place
(i.e., 21.5 g and 0.0215 kg are both the same value and both have three significant digits).
Exact Numbers
Exact numbers are not uncertain and are said to have an infinite number of significant digits.
1. Numbers that are defined.
For example 1000 kg = 1 t exactly.
2. Numbers that result from counting objects.
For example: 32 students, 158 beakers, $4.95 (exactly 495 cents).
Table D2
Examples of Counting Significant Digits
Measured Value
# sig. digs.
Measured Value
156 g
3
120.50 L
0.2602 m
4
0.050 02 s
6.02 x 1023 molecules
3
7.2 ºC
The italics digits above are uncertain (estimated).
# sig. digs.
5
4
2
Rounding Off
1. When the first digit after those retained is less than 5, all digits retained remain the same.
(eg., 2.249 g = 2.2 g).
2. When the first digit after those retained is 5 or greater, the last digit retained is increased by
one (eg., 12.654 cm3 = 12.7 cm3 ).
e.g.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
Measured Value
#s.d.
12.42 g
0.1407 m
10.0 mL
1000 ºC
0.060 h
126 km
15.00 t
0.0004 kPa
40 s
0.0100 L
100 cm/m
4
Calculated value
e.g.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
0.1495 m2
29.95 m/s
139.49 cm3
10.54 mol
100.4 ºC
9.998 g
80.46 km/h
197.042 L
0.0462 m3
8.29 g/mol
5.49 mm
Rounded-Off
(#s.d.)
0.150 m2
(3)
(3)
(3)
(3)
(3)
(2)
(2)
(4)
(2)
(2)
(1)
30
The Mole
Molar Mass
Determine the molar mass(mass of one mole) of each of the following substances. Show all work as in
the example.
A number with two decimal places multiplied by an exact number has two decimal places in the answer.
See notes below.
3. sodium carbonate decahydrate
(washing soda)
(List water of hydration as
10H2O = 10 x 18.02 = 180.20)
4. MgSiO3
(asbestos)
5. Sodium hypochlorite
(laundry bleach)
6. Al(OH)3
(water clarifier)
7. Sodium chloride
(table salt)
8. Calcium carbonate
(limestone)
9. dinitrogen oxide
(anesthetic)
10. Na2S2O3.5H2O
(photographic hypo)
11. NH4H2PO4
(fertilizer)
Notes:
1. The rule for multiplication and division and the rule for addition and subtraction are followed in this
example and Chemistry 20 and 30. These rules are used in the key for answers and on exams. If
these rules are followed everyone can expect to get the same answer.
2. For hydrates the molar mass of water (18.02 g/mol) should be memorized.
31
THE MOLE
MASS TO MOLES CALCULATIONS
Determine the number of moles in 1000 g (1.000 kg) of each of the following. Show all work
exactly as in the example. Use correct SI symbols and significant digits
7. The container with the greatest number of moles in 1000 g of compound is ______________.
32
THE MOLE
MOLES TO MASS CALCULATIONS
Follow the example to show all work and calculate the mass of each sample. Use correct SI
symbols and significant digits
In order to determine the mass of a number of
identical things, multiply the number of things
by the mass of one thing.
Example:
mass = number of things x mass of one thing
or mass = # dozen x mass of one dozen
or mass = # moles x molar mass
1 Na = 1 x 22.99 = 22.99
1 Cl = 1 x 35.45 = 35.45
58.44 g/mol
0.21 mol of table salt
m
or if m = mass, n = # moles and M = molar
mass
then m = nM
=
NaCl =
nM
0.21 mol x 58.44 g/mol
= 12 g
1.
0.100 mol of cream of tartar (KHC4H4O6)
2. 1.2 mol of detergent filler
(Na2SO410H2O)
3.
0.15 mol of white phosphorus
4.
55.56 mol of water
5.
0.025 mol of tin(II) fluoride
6.
0.400 mol of gypsum (CaSO42H2O)
33
Molar Mass Exercise
Find the molar mass of the following compounds.
1.
AgNO3
M = _______________
2.
CaCl2
M = __________________
3.
Ca(NO3)2
M = _______________
4.
CO2
M = __________________
5.
CH4
M = _______________
6.
Cu(NO3)2
M = __________________
7.
Fe2O3
M = _______________
8.
FeCl3
M = __________________
9.
HNO3
M = _______________
10. HCl
M = __________________
11. Hg(OH)2
M = _______________
12. K2SO4
M = __________________
13. NaClO3
M = _______________
14. NH4OH
M = __________________
15. NH4ClO3
M = _______________
16. SO2
M = __________________
17. ZnSO4
M = _______________
18. Zn3(PO4)2
M = __________________
Moles and Mass
1. What is the mass of a mole of the following elements?
a) Cu ________
b) F2 _________
c) Fe _________
d) H2 _________
e) Hg ________
f) I2 _________
g) K _________
h) P4 _________
2. Find the molar mass of the following compounds. Show the calculations and molar mass of
each element present.
a) C4H10 _________
b) H2S
_________
c) Hg(NO3)2 _________
a) Fe2S3 _________
b) Zn(OH)2 _________
c) Pb(ClO3)2 _________
Find the mass of the following.
1. 0.234 mol of sodium sulfate
formula ______________(
)
M = ______
______ g
2. 0.14 mol of potassium nitrate
formula ______________(
)
M = ______
______ g
3. 1.2 mol of magnesium chloride
formula ______________(
)
M = ______
______ g
34
4. 0.0850 mol of aluminum nitrite
formula ______________(
)
M = ______
______ g
5. 0.984 mol of strontium oxide
formula ______________(
)
M = ______
______ g
Find the number of moles in the following.
1. 136 g of magnesium sulfate
formula ______________(
)
M = ______
______ mol
2. 3.60 g of water
formula ______________(
)
M = ______
______ mol
3. 83.1 g of potassium chloride
formula ______________(
)
M = ______
______ mol
4. 4.00 g of ethanol
formula ______________(
)
M = ______
______ mol
5. 95.60 g of barium bromide
formula ______________(
)
M = ______
______ mol
35
In each sample, determine what amount (in mol) of the compound is present. Show your work on
a separate piece of paper.
(a) 8.40 g of NaOH
(b) 4.2 kg of water
(c) 0.0240 kg of Na2SO4
(d) 1.77 g of CuSO3
(e) 1.00 kg of methane, CH4
(f) 45.0 g of ammonia, NH3
Balance the chemical equation in each table, and use the mole numbers to complete the table.
For example:
2H2O
→
2H2
+
O2
5.00 mol
5.00 mol
2.50 mol
1.20 mol
1.20 mol
0.600 mol
3.00 mol
3.00 mol
1.50 mol
(a)
_____ Na
+
_____ H2O
→
_____ H2
+
_____NaOH
8.0 mol
0.20 mol
4.80 mol
16.0 mol
(b)
_____ Al2(SO4)3
+
_____ NH4OH
→
_____ Al(OH)3
+
_____(NH4)2SO4
3.0 mol
2.0 mol
1.00 mol
36