Chemical reactions
A chemical reaction occurs when one or more substances are changed into a
new substance/substances.
A chemical reaction can be represented by a chemical equation.
Reactants/reagents – substances that undergo a chemical change (left side,
usually).
Products – new substances formed (right side, usually).
The law of conservation of mass (matter) states that matter is neither created
nor destroyed in a physical process or a chemical reaction. (i.e. mass of
reactants = mass of products)
The law of conservation of energy states that energy can neither be created
nor destroyed in a physical or chemical process.
Chemical reactions can be described in different ways.
E.g. Rusting:
“Iron reacts with oxygen to produce iron (III) oxide (rust).”
Iron + oxygen  iron (III) oxide
A skeleton equation is a chemical equation that does not indicate the relative
amounts of the reactants and products.
Fe + O2  Fe2O3
Write a skeleton equation for each of these chemical reactions:
1. Aluminum metal reacts with oxygen in the air to form aluminum oxide.
2. When solid mercury (II) sulfide is heated with oxygen, liquid mercury metal
and gaseous sulfur dioxide are produced.
3. Oxygen gas and potassium chloride can be made by heating potassium
chlorate.
4. Glucose (C6H12O6) burns in the presence of oxygen gas to form carbon
dioxide and water (dihydrogen oxide).
A balanced equation has the same number of atoms of each element on each
side of the equation (necessary for the law of conservation of mass).
Rules for Balancing Equations:
1. Determine the correct formulas for all the reactants and products in the
reaction.
2. Write the formulas for the reactants on the left and the formulas for the
products on the right with an arrow in between. Separate two or more
reactants or products with plus signs.
3. Count the number of atoms of each element in the reactants and products.
A polyatomic ion appearing unchanged on both sides of the equation is
counted as a single unit.
4. Balance the elements one at a time by using coefficients. When no
coefficient is written, it is assumed to be 1. It is best to begin with an
element other than H or O. (Do not balance by changing the subscripts in
the chemical formulas).
5. Check each atom or polyatomic ion to be sure the equation is balanced.
6. Make sure all coefficients are in the lowest possible ratio.
Types of Chemical Reactions
Combination (Synthesis) Reactions – two or more substances react to form a
single substance.
The basic formula of a combination reaction is:
A
E.g.
+

B
Mg
+
AB
O2

MgO
Do # 5 (p. 152)
a) Ca +
b) Fe +
c) P +
d) N2O5
e) Na2O
f) Mg +
S
O2
O2
+
+
O2
H2O
H2O






calcium sulfide
iron (III) oxide
diphosphorous oxide
hydrogen nitrate
sodium hydroxide
magnesium oxide
Decomposition Reactions – a single compound is broken down into two or
more simpler products.
The basic formula of a decomposition reaction is:

AB
E.g.
CaCO3
A

+
B
CaO +
Do # 6 (p. 153)
a) nickel (II) carbonate  nickel (II) oxide + carbon dioxide
b) Ag2O 
c) Ammonium nitrate  dinitrogen monoxide + water
CO2
Single-Replacement Reactions – atoms of an element replace the atoms of a
second element in a compound. Whether one metal will replace another metal
can be determined by the relative reactivity of the two metals.
The activity series of metals lists metals in order of decreasing reactivity. A
reactive metal will replace any metal found below it in the activity series. (see
Table 7-2 Activity Series of Metals, p. 155)
The general formula for a single-replacement reaction is:
A
+
BC
A
+
BC
E.g.
Fe
+

or

AC
+
B
BA
+
C

CuSO4
FeSO4+
Cu
Do # 7 (p. 155)
Double-replacement reactions – involve an exchange of positive ions between
two compounds. One of the following are usually true about at least one of the
products in a double-replacement reaction: a) it is only slightly soluble and
precipitates from solution; b) it is a gas that bubbles out of the mixture; and c)
it is a molecular compound (two non-metals) such as water.
The general formula for a double-replacement reaction is:
TU
+
RS

TS
+
RU
Do #8 (p. 157)
Combustion Reactions
A combustion reaction occurs when oxygen reacts with another substance
(usually hydrocarbons, CxHy), often producing energy in the form of heat and
light.
In the combustion of hydrocarbons, the products are carbon dioxide and water.
When the combustion is incomplete, the products are carbon monoxide and
water.
E.g.
C6H6 + O2  CO2 +
Do #9 (p. 159)
H2O
Stoichiometry
Stoichiometry is the calculation of quantities in chemical equations by using the
known ratios of reactants and products in a balanced equation.
Do Problem 1 (p 172)
Mole to Mole Calculations
The moles of wanted quantity can be calculated by multiplying the moles of the
given quantity (G), whether reactant or product, by a mole ratio from the
balanced equation with the given in the denominator.
mol of G
x
b moles of W
= mol of W
a moles of G
Do Problem 2 (p 174)
Mass to Mass Calculations
1. The mass of given quantity must be changed to moles by using (dividing
by) molar mass of substance.
Mass of G x 1mol
=
mol G
g
Molar mass of G
2. Change the moles of G (given) to moles of W (wanted) by multiplying by
W/G ratio.
mol of G
=
mol of W
G
3. Change the moles of W to grams by multiplying by molar mass.
mol of W
x
x
W
mass of W =
1 mol
mass of W
Other Stoichiometric Calculations
In order to solve other stoichiometic problems always use this pattern of steps:
1. Determine what information is given (G) and what information is wanted
(W).
2. Find the W/G ratio for the compounds in the chemical equation.
3. Take the information of the given compound (# particles, mass or volume)
and convert it to moles of given compound by dividing by the following
conversion factors (Avogadro’s number, molar mass of substance, or
molar volume).
4. Change the given moles to wanted moles by multiplying by the W/G ratio.
5. Convert the wanted moles into wanted amounts (# particles, mass or
volume) by multiplying by the following conversion factors (Avogadro’s
number, molar mass of substance, or molar volume).
See & copy figure 8.8 (p. 177)
Do # 3 & #4 (p177) Do #6 (p179)
Limiting Reagents
Limiting Reagents are the compounds or molecules in a chemical reaction that
limits or determines the amount of product that can be formed in a reaction.
The excess reagent is the compound or molecule that there is more than enough
of in order to react with the limiting reagent.
E.g. 4 slices of bread + 1 jar of PB  2 PB sandwiches
LR
XS R
product
See Examples 9 & 10 (pp. 182-183)
Do # 7 & # 8 (pp. 182 – 184)
Percent Yield
Reactions do not always go to completion. Impure reactants and competing side
reactions cause other products to be formed.
When an equation is used to calculate the amount of product that will form during
a reaction, then a value for the theoretical yield is obtained.
The amount of product that forms when the reaction is carried out in the
laboratory is called the actual yield.
The percent yield is the ratio of the actual yield to the theoretical yield.
Percent yield
=
actual yield
theoretical yield
See example 11 (p. 184); do question 9 (p. 185)
x 100%
Energy Changes in Chemical Reactions
In many chemical reactions energy is needed to form new compounds. Energy is
stored in chemical bonds between atoms in molecules.
Exothermic reactions – release energy in the form of heat.
E.g. the combustion of coal.
C
+
O2
 CO2
+ 393.5 kJ
The energy stored in the bonds of the reactants may be released as heat, the
rest of the energy in the bonds of the reactants goes into the bonds of the
products
Endothermic reactions – reactions in which energy is absorbed.
E.g. The production of calcium oxide (lime)
CaCO3
+
176 kJ 
CaO +
CO2
The energy content of the products is higher than the energy content of the
reactants.
Equations that include the amount of heat produced or absorbed by a reaction
are thermochemical equations.See example 12 (p. 187); Do problems 10 & 11
(p. 187)
Heat of Reaction
Enthalpy (H)– is the amount of heat that a substance has at a given temperature
and pressure. H cannot be measured, but a change in enthalpy can be measured
and is symbolized ΔH.
Changes in enthalpy occur whenever heat is released or absorbed. The heat that
is absorbed or released in a reaction is called, the heat of reaction.
When one mole of a substance is completely burned, ΔH is called, the heat of
combustion.
ΔH is always positive for endothermic reactions (enthalpy of products > enthalpy
of reactants).
ΔH is always negative for exothermic reactions (enthalpy of products < enthalpy
of reactants).
Problem 12 (p. 189)
H2O(g)  H2O(l)
+
1
:
1
1 g H2O(g)
x
44.0 kJ
:
44.0 kJ
1 mol
18.0 g H2O
=
W=
G
0.0556 mol H2O
44.0 kJ
1 mol
0.0556 mol H2O(g) x
44.0 kJ
=
1 mol H2O
2.45 kJ
2.45 kJ x
1 kcal =
0.585 kcal
=
585 cal
4.18 kJ
The ΔH for a reaction in which 1 mol of a compound is formed from its elements
is the standard heat of formation, ΔH°f , of that compound.
See table 8.1 (p. 190) for ΔH°f of common substances. The more negative the
value the more stable the compound.
The ΔH of free elements in their standard form = 0.
E.g. O2(g), N2(g), H2(g), F2(g), Cl2(g), Br2(l), I2(s), or C(s)
ΔH for a reaction is the difference between the standard enthalpies of formation
of all reactants and all products.
ΔH
= ΔH°f (products) – ΔH°f (reactants)
See example 13 (p. 190); Do problem 13 (p. 190).
a) CH4(g)
ΔH
+
2O2(g) 
= ΔH°f (products)
CO2(g) +
–
2H2O(l)
ΔH°f (reactants)
(ΔH°f CO2(g) + 2(ΔH°f H2O(l)) – (ΔH°f CH4(g) + 2(ΔH°f O2(g))
(-393.5 + 2(-285.8)) – (-74.86 + 2(0.0)) = ΔH
ΔH = -890.2 kJ
b)
ΔH = -566 kJ