Chemical reactions
Classifications
Reactions in solution
Ionic equations
Learning objectives
Distinguish between chemical and physical change
Write and balance chemical equations
Describe concepts of oxidation and reduction
Classify reaction according to types of reactants and
products
Distinguish among strong, weak and non-electrolytes
Identify common acids and bases by from chemical
formula
Predict formation of precipitates by application of
solubility rules
Write total and net ionic equations from balanced
molecular equations
Chemical vs physical redux
Physical: No new substance!
Chemical: New substance formed!
Evidence for chemical reactions
The chemical equation
aA + bB = cC + dD
Reactant
side
coefficient
ELEMENT or
COMPOUND
Product
side
Law of Conservation of Matter states that matter is
neither created nor destroyed
Means: atoms on left equals atoms on right
Chemical book-keeping
Keys to balancing equations:
“Have I gained or lost any atoms?”
Put down the correct formula for each
reactant or product
Formulas cannot be changed in order to
balance the equation
 Reaction of hydrogen with oxygen to produce
water: reactants are H2 and O2, product is H2O
2 H2
+
O 2 → 2 H 2O
The big number
multiplies every
atom after it
 Count the atoms: 4 H and 2 O
The subscript
only multiplies
the atom before it
4 H and 2 O
Molecular representation of the
reaction
Balance the equations
A method of trial and error
CH4 + O2 = CO2 + H2O
Balance the equations
CH4 + O2 = CO2 + H2O
– CH4 + 2O2 = CO2 + 2H2O
C3H8 + O2 = CO2 + H2O
– C3H8 + 5O2 = 3CO2 + 4H2O
N2 + H2 = NH3
– N2 + 3H2 = 2NH3
Do balancing equation exercises
One approach to classification
Oxidation – reduction: focusing on
electrons
Oxidation is loss of electrons
Reduction is gain of electrons
Oxidation is always accompanied by
reduction
The total number of electrons is kept constant
Oxidizing agents oxidize and are
themselves reduced
Reducing agents reduce and are
themselves oxidized
Redox in chemistry
All reactions involve rearrangement of
atoms and molecules
Some reactions involve rearrangement of
atoms and molecules and electrons
– Photosynthesis, respiration, combustion...
These are called redox reactions
Any reaction involving elements must be
redox
Combination (synthesis)
reactions
Element + element 
compound (redox)
– S + O2 → SO2
– Metal + nonmetal  binary
ionic compound
– Nonmetal + nonmetal 
binary covalent compound
Compound + element 
compound (redox)
– CO + O2 → CO2
Compound + compound
 compound
– SO2 + H2O →H2SO3
Decomposition reactions
Compound  element +
element (redox)
– HgO → Hg + O2
Compound  element +
compound (redox)
– PCl5 → PCl3 + Cl2
Compound  compound
+ compound
– CaCO3 → CaO + CO2
Single replacement (displacement)
Element displaces another
element from compound
(redox)
Cu + 2 AgNO3 → Cu(NO3)2 + 2 Ag
Predicting single replacement
reactions: the activity series
Element higher in the
will displace one
lower in the series
The element higher is
a stronger reducing
agent
The element lower is
a stronger oxidizing
agent
Three types of double
displacement reaction
Compounds
exchanging partners
– Usually ionic
compounds in solution
Precipitation
Acid-base
neutralization
Gas formation
Precipitation
Identify ions and swap them
BaCl2 + Na2SO4 → BaSO4 + 2 NaCl
Acid – base neutralization:
special case of double replacement
BASE
ACID
SALT
WATER
KOH(aq) + HNO3(aq) = KNO3(aq) + H2O(l)
Product is liquid water not a solid
Gas formation
Product is either a gas or is unstable and
decomposes to a gas
CaCO3(s) + 2 HCl(aq) = CaCl2(aq) + H2O(l) + CO2(g)
Writing balanced molecular equations
for double replacement reactions
Use correct formulae
– Metal ion charge predicted
from group number
– Use table for correct
formula and charge for
polyatomic ions
Identify as solid (s), gas
(g), liquid (l) or dissolved
(aq)
Balance: atoms (groups)
on left = atoms (groups)
on right
Balancing double replacement equations
It’s very much a matter of states – show
them!
Pb(NO3)2(aq) + 2KI(aq) = 2KNO3(aq) + PbI2(s)
Balance polyatomic ions as units:
– Pb2+, K+, I-, NO3-
Left hand side
Right hand side
1 Pb2+
1 Pb2+
2 NO3-
2 NO3-
2 K+
2 K+
2 I-
2 I-
Molecular equation for reaction of
Na2SO4 + Ba(NO3)2
Combustion
Element or compound
reacting with oxygen
(redox)
– CH4 + O2 → CO2 + H2O
Associated with
production of heat and
light
Involves hydrocarbons
(CxHy), nonmetals (S) and
metals (Mg)
Sorting solution reactions:
dissolved species
Electrolytes:
– Ionic compounds produce ions in solution
(NaCl, NH4NO3 etc.)
Non-electrolytes:
– Covalent compounds do not produce ions in
solution (CH3OH, C6H12O6 etc.)
Electrolytes: distinguishing by
strength
All soluble substances that produce ions are
electrolytes
Strong electrolytes are characterized by
complete dissociation in water
Weak electrolytes dissociate to a much smaller
extent.
Strong, weak or non electrolyte?
All soluble salts are strong electrolytes
Strong acids and bases are strong
electrolytes
Weak acids and bases are weak
electrolytes
Insoluble compounds are non-electrolytes
Molecular compounds are non-electrolytes
Classification of electrolytes
Strong
electrolytes
Weak
Nonelectrolytes electrolytes
ACIDS:
HCl, HBr, HI
HClO4, HNO3, H2SO4
ACIDS:
HF, H3PO4,
CH3CO2H
SALTS:
KBr, Na3PO4
SALTS:
None
BASES:
NaOH, Ba(OH)2
BASES:
NH3
Molecular
covalent
compounds:
H2O,
CH3OH,
C12H22O11
(sucrose)
Most organic
compounds
and
INSOLUBLE
salts
Flow chart for determining type of
electrolyte
Yes
No
1. Soluble in H2O?
Yes
No
2. Acid or base?
Weak
Strong
3. Weak or strong?
Weak
electrolyte
Cov
Ionic
3. Ionic or covalent?
Strong
electrolyte
Nonelectrolyte
Recognizing acids and bases
Acids usually have H at the beginning of the
formula – HCl
Bases usually have OH in the formula – NaOH
– Not in organic compounds though - CH3OH
Acid formula
Name
Base formula
Name
HCl
Hydrochloric acid
NaOH
Sodium hydroxide
H2SO4
Sulfuric acid
KOH
Potassium
hydroxide
H3PO4
Phosphoric acid
Ba(OH)2
Barium hydroxide
HNO3
Nitric acid
NH3
Ammonia
HClO4
Perchloric acid
(CH3)3N
Trimethylamine
CH3CO2H
Acetic acid
HCO2H
Formic acid
Citric acid
The strong acids and bases
Strong acids
(Only six)
Strong bases
(g1A and g2A)
HCl
Hydrochloric acid
LiOH
Lithium hydroxide
HBr
Hydrobromic acid NaOH
Sodium hydroxide
HI
Hydroiodic acid
KOH
Potassium
hydroxide
HNO3
Nitric acid
Ca(OH)2
Calcium
hydroxide
H2SO4
Sulfuric acid
Sr(OH)2
Strontium
hydroxide
HClO4
Perchloric acid
Ba(OH)2
Barium hydroxide
Solubility rools
Group I and ammonium (NH4+) compounds soluble
Nitrates (NO3-), acetates (CH3CO2-) soluble
Chlorides, bromides and iodides generally soluble
{except Pb(II), Ag(I) and Hg(I)}
Sulphates (SO42-) generally soluble (except g2A and
Pb2+)
Carbonates (CO32-), phosphates (PO43-) generally
insoluble (except gIA)
Hydroxides (OH-), sulphides (S2-) generally insoluble
(except gIA and gIIA)
Total ionic equations
Pb(NO3)2(aq) + K2CrO4(aq) = 2KNO3(aq) + PbCrO4(s)
Total ionic equation
Dissolved substances:
– Strong electrolytes show as ions
– Weak or non- electrolytes show as molecular
formula
Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + CrO42-(aq) =
2K+(aq) + 2NO3-(aq) + PbCrO4(s)
Net ionic equations
Spectator ions are those ions that do not
undergo a change; they do not participate
in the chemical change and are the same
on both sides of the equation
Remove all spectator ions from the
equation
Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + CrO42-(aq) =
2K+(aq) + 2NO3-(aq) + PbCrO4(s)
Net ionic equations
Pb2+(aq) + CrO42-(aq) = PbCrO4(s)
Mass and charge must still balance, although
overall charge may not be neutral in a net ionic
equation
Net ionic equation for reaction of
Na2SO4 + Pb(NO3)2