Ch. 9 – Chemical Reactions
I.
Intro to Reactions
(p. 282 – 285)
I
II III IV V
Chemical Reaction

Chemical change

Atoms of one or more substances (reactants)
are rearranged into new substances (products)

Signs of a chemical reaction
 Change of temperature or light
 Formation of a gas (bubbles)
 Formation of a precipitate
 Odor
 Color change
 Change in volume
Law of Conservation of Mass

mass is neither created nor destroyed
in a chemical reaction
total mass stays the same
 atoms can only rearrange

4H
36 g
2O
4H
2O
4g
32 g
Chemical Equations
A+B  C+D
REACTANTS
PRODUCTS
Starting substances
Substances formed
Chemical Equations
p. 283
Writing Equations
2H2(g) + O2(g)  2H2O(g)

Identify the substances involved.

Use symbols to show:
 How many? - coefficient
 Of what? - chemical formula
 In what state? - physical state

Remember the diatomic elements!
Describing Equations

Describing Coefficients:
 individual atom = “atom”
 covalent substance = “molecule”
 ionic substance = “unit”
3CO2  3 molecules of carbon dioxide
2Mg
 2 atoms of magnesium
4MgO  4 units of magnesium oxide
Writing Equations
Two atoms of solid aluminum react
with three units of aqueous
copper(II) chloride to produce three
atoms of solid copper and two units
of aqueous aluminum chloride.
• How many?
• Of what?
• In what state?
2 Al (s) + 3 CuCl2 (aq)  3 Cu (s) + 2 AlCl3 (aq)
Describing Equations
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
• How many?
• Of what?
• In what state?
One atom of solid zinc react with two molecules
of aqueous hydrochloric acid to produce one unit
of aqueous zinc chloride and one molecule of
hydrogen gas.
Ch. 9 Chemical Reactions
II. Balancing Equations
I
II III IV V
Why is there a 2 after the oxygen?
__________________
Oxygen
is diatomic
4
1
12
2
10
4
2
8
Balancing Steps
1. Write the unbalanced equation.
2. Count atoms on each side.
3. Add coefficients to make #s equal.
Coefficient  subscript = # of atoms
4. Reduce coefficients to lowest
possible ratio, if necessary.
5. Double check atom balance!!!
Helpful Tips
Balance one element at a time.
 Update ALL atom counts after adding
a coefficient.
 If an element appears more than
once per side, balance it last.
 Balance polyatomic ions as single
units.
 “1 SO4” instead of “1 S” and “4 O”

Practice Balancing!
All the atoms on the reactants side
must equal all the atoms on the
products side
 First, you need to count all the atoms

H2
+
O2
2
2

H
O
H2O
2
1
How to Balance
Are they equal?
 Add coefficients to change the
number of atoms

Balanced!! Good job 
2 H2
+
O2
42
2
 2 H2O
H
O
2 4
1 2
Balancing Example
Aluminum and copper(II) chloride react
to form copper and aluminum chloride.
2 Al + 3 CuCl2  3 Cu + 2 AlCl3
2 1
Al
1 2
3 1
Cu
1 3
6 2
Cl
3 6
Practice…
Great!
Sn + 2 HF 
SnF2
1
Sn
1
2
1
H
2
2
1
F
2
+
H2
Cu + 2 AgNO3  Cu(NO3)2 + 2Ag
1
Cu 1
2
1
Ag 1
2
1 NO3 2
2
Good
Job!
Ch. 9 – Chemical Reactions
III. Types of
Chemical
Reactions
I
II III IV V
Combustion
The burning of any substance in O2 to
produce heat, usually a hydrocarbon
 Products are always carbon
dioxide and water vapor

CXHY + O2  CO2 + H2O
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Synthesis (composition)

the combination of 2 or more
elements or compounds to form a
more complex compound

Basic Form: A + X  AX
Examples of Synthesis Reactions
1) Metal + oxygen → metal oxide

EX. 2 Mg(s) + O2(g) →
2 MgO(s)
2) Nonmetal + oxygen → nonmetallic oxide
EX. C(s) + O2(g) → CO2(g)
3) Metal + nonmetal → salt
 EX. 2 Na(s) + Cl2(g) →
2 NaCl(s)
4) A few nonmetals combine with each other.
 EX. 2 P(s) + 3 Cl2(g) →
2 PCl3(g)

Decomposition

A single compound breaks down into
its component parts of simpler
compounds

Basic form: AX  A + X
Examples of Decomposition Rxns
1) Some oxides, when heated, decompose.

EX. 2 HgO(s) →
2 Hg(l) + O2(g)
2) Some are produced by electricity.

EX. 2 H2O(l) → 2 H2(g) + O2(g)
EX. 2 NaCl(l) → 2 Na(s) + Cl2(g)
Single Replacement
a more active element takes the place
of another element in a compound and
sets the less active one free
 Basic Form: A + BX  AX + B
or AX + Y  AY + X
Examples of replacement reactions:
1) Replacement of a metal in a compound
by a more active metal.
 Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Examples of Single Replacement Rxns
2) Replacement of hydrogen in water by an active
metal.

Note: it is helpful to think of water as HOH (H+OH-)
2 Na(s) + 2 H2O(l) → 2 NaOH(aq)
or
2 Na(s) + 2 HOH(l) → 2 NaOH(aq)
 Mg(s) + 2 H2O(g) → Mg(OH)2 (s)
or
Mg(s) + 2 HOH(g) → Mg(OH)2 (s)

+ H2(g)
+ H2(g)
+ H2(g)
+ H2(g)
Examples of Single Replacement Rxns
3) Replacement of hydrogen in acids by active
metals.
 EX. Zn(s) + 2 HCl(aq) →
ZnCl2(aq) + H2(g)
4) Replacement of nonmetals by more active
nonmetals.
 EX. Cl2(g) + 2 NaBr(aq) →
2 NaCl(aq) + Br2(l)
 If……

EX. Br2(g) + 2 NaCl(aq) →
No Reaction!
Single Replacement
NOTE: Refer to the activity series
for metals and nonmetals to predict
products of replacement reactions.
 If the free element is above the
element to be replaced in the
compound, then the reaction will
occur.
 If it is below, then no reaction (NR)
occurs.

Double Replacement (Ionic)




ions in two compounds “change partners” ;
cation of one compound combines with anion of
the other
Occurs between ions in aqueous solution. A
reaction will occur when a pair of ions come
together to produce at least one of the following:
 a precipitate (s)
 water (l)
 a gas (g)
Basic form: AX + BY → AY + BX
Complex form:
ABX + CDY → ADY + CBX
Double Replacement


NOTE: Use the solubility rules to decide
whether a product of an ionic reaction is
insoluble in water and will thus form a
precipitate. If a compound is soluble in
water then it should be shown as being in
aqueous solution, or left as separate ions.
If it products are both in solution (aqueous)
then
 No Net Ionic Reaction occurs!
Aqueous Rxns: Double Replacement

Ex: in the reaction involving the ionic compounds
silver nitrate and potassium chloride we have:

AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)

The driving force for double replacement
reactions is the removal of ions from solution
Precipitation Reactions




Double replacement rxns that result in an insoluble
precipitate
Solubility: amount of a substance that can be
dissolved in a given quantity of water
The solubility of an ionic compound determines
whether it will precipitate or not.
The reaction of KI and Pb(NO3)2:
2KI(aq) +Pb(NO3)2(aq)  PbI2(s) + 2KNO3 (aq)
 The PbI2 is an insoluble ionic compound that
will precipitate out of the solution
Precipitation Reactions




Can we predict whether an ionic compound will be
soluble or not?
If an ionic compound is insoluble it means that
neighboring ions have an attraction for each other
that is greater than the attraction of water for the
ions
Unfortunately, there are no clear rules for solubility
based on physical properties of ions.
General behaviors of certain ions are observed
by consulting A Solubility Table!
Neutral Molecular Compounds
Even though the neutral molecular
compound may be soluble in aqueous
solution, its formation is essentially irreversible.
 Thus, ions are effectively removed from
solution by this irreversible process
 The neutralization reaction of HCl and NaOH:
HCl (aq) + NaOH(aq)  H2O(l) + NaCl(aq)
 The formation of the covalent compound (H2O)
from the proton and hydroxide ions is
essentially irreversible and drives the double
replacement reaction (even though we would
consider H2O to be "highly soluble" in H2O).

Gas Formation in Double Replacement

When a double replacement reaction involves the
formation of a gas (& the gas is not soluble in H2O)
the loss of the gas can drive the double
replacement reaction (i.e. the gas is lost - therefore,
it is an irreversible process)

Gasses that can form from ionic compounds include:
 CO2 (carbon dioxide)
 H2S (hydrogen sulfide - smells like rotten eggs)
 NH3 (ammonia).
Formation of CO2 from carbonic acid
The bicarbonate ion (HCO3-) can combine with a
proton to produce carbonic acid (H2CO3):
HCl (aq) + NaHCO3 (aq)  NaCl (aq) + H2CO3 (aq)
 Carbonic acid in water is unstable & spontaneously
decomposes to form water and carbon dioxide gas:
H2CO3 (aq)  H2O (l) + CO2 (g)
 The carbon dioxide gas is lost, & thus the formation
of carbonic acid is irreversible and drives the double
replacement reaction. The reaction is thus:
HCl (aq) + NaHCO3 (aq)  NaCl (aq) + H2O (l) + CO2 (g)

Gases produced in Spontaneous Rxns

The three typical gasses produced take
place with the following spontaneous
reactions (only when they appear on the
product side):
KNOW THEM
 H2CO3(aq) → CO2(g) + H2O(l)
 NH4OH(aq) → NH3(g) + H2O(l)
 H2SO3(aq) → SO2(g) + H2O(l)
Net Ionic Reactions - DR
Complete ionic equation = equation
that shows all dissolved ionic
compounds as dissociated free ions
 Net ionic equation = Equation for a
reaction in solution that shows only
those particles that are directly
involved in the chemical change
 Spectator Ion = an ion that appears
on both sides of an equation and is
not directly involved in the reaction

Neutralization of Nitric Acid & KOH




The molecular equation would be:
HNO3 (aq) + KOH (aq)  KNO3 (aq) + H2O (l)
The complete ionic equation would be:
H+1 (aq) + NO3 -1 (aq) + K+1 (aq) + OH -1 (aq) 

K+1 (aq) + NO3 -1 (aq) + H2O (l)
The net ionic equation would therefore be:
H+1 (aq) + OH -1 (aq)  H2O (l)

Spectator Ions: NO3 -1 (aq) + K+1 (aq)

Net Ionic Reactions! - DR
Molecular (or Full) Equation:
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
CIE:
Ag+(aq) + NO3– (aq) + K+(aq) + Cl–(aq)  AgCl(s) + K+(aq) +NO3–(aq)
NIE:
SI:
Ag+(aq) + Cl–(aq)  AgCl(s)
K+(aq) and NO3–(aq)
Net Ionic Reactions! - DR
Molecular (or Full) Equation:
Pb(NO3) 2(aq) +2LiCl(aq)  PbCl2 (s) + 2LiNO3(aq)
CIE:
Pb+2(aq)+2NO3–(aq)+2 Li+(aq)+2Cl–(aq)PbCl2(s)+2Li+(aq)+2NO3–(aq)
NIE:
SI:
Pb+2(aq) + 2Cl–(aq)  PbCl2 (s)
Li+(aq) and NO3–(aq)