Reactions in Aqueous Solutions
7.1 Predicting Whether a Reaction Will
Occur?
 Driving Forces
 Formation of solid
 Formation of water
 Transfer of electrons
 Formation of a gas
7.2 Reactions in Which a Solid Forms?
 Driving forces
 Precipitation – formation of solid
E.g Ba(NO3)2(aq) + K2CrO4(aq)  yellow solid
yellow sol.
colorless
What happens When an Ionic Compound
Dissolves in Water?
 Ba(NO3)2(aq) => barium nitrate (white solid) has been
dissolved in water
 Containing Ba2+ (aq) and 2 NO3-(aq) ions
 Strong electrolytes: A unit of substance that dissolves in
water produces separate ions
 E.g K2CrO4(aq) contains K+, K+ and CrO42- ions
 => strong electrolyte
Strong Electrolytes
 Ba(NO3)2(aq) + K2CrO4(aq)  yellow solid
How to Decide What Products Form
 Write out all possible formula that can be formed
NO3K+
KNO3
Ba2+
Ba(NO3)2
CrO42-
Using
Solubility
Rules
 General Rules for Solubility of ionic compounds (salts) in water at
25oC
 Most nitrate (NO3-) salts are soluble
 Most salts of Na+, K+ and NH4+ are soluble
 Most chloride, bromide, iodide (Cl-, Br- and I-) salts are soluble.
Notable exceptions are AgCl, PbCl2 and Hg2Cl2.
 Most sulfate salts are soluble. Notable exceptions are BaSO4,
PbSO4 and CaSO4
 Most hydroxide (-OH) compounds are slightly soluble. The
important exception are NaOH and KOH. Ba(OH)2 and Ca(OH)2
are only moderately soluble
 Most sulfide (S2-), carbonate (CO32-) and phosphate (PO43-) salts are
slightly soluble
* The term insoluble and slightly soluble really mean the same thing
Identifying Precipitates in Reactions
Where a Solid Forms
 Step 1: Write the reactants as they actually exist before any
reaction occurs
 Step 2:
Consider the various solids that could form.
 Step 3: use solubility rules to decide whether a solid forms
E.g
AgNO3(aq) + KCl(aq)  white solid
Examples
 Using solubility rules to predict the product of Reactions
 KNO3(aq) + BaCl2(g) 
 Na2S(aq) + Cu(NO3)2(aq) 
 KOH(aq) + Fe(NO3)2(aq) 
Describing Reactions in Aqueous
Solutions
 Molecular equation: shows overall reaction but not
necessary the actual forms of the reactants and
products in solution
 Complete ionic equation: represents all reactants and
products that are strong electrolytes as ions. All
reactants and products are included.
 The net ionic equation includes only those
components that undergoes a change. Spectator ions are
not included
Examples
For each of the following reactions, write molecular
equation, the complete ionic equation and the net ionic
equation
 Aqueous sodium chloride is added to aqueous silver nitrate to
form solid silver chloride plus aqueous sodium nitrate
 Aqueous nickel (II) nitrate is added to aqueous potassium
carbonate to form solid nickel (II) carbonate and aqueous
potassium nitrate
7.4 Reactions That Form Water: Acids
and Bases
 Arrhenius acids: a substance that produces H+ ions
(protons) when it dissolved in water
 Strong acids: strong electrolytes
 Common strong acids: HCl, HBr, HI, H2SO4, HClO4, HNO3
dissolved in H2O
HCl(aq) -------------- H+(aq) + Cl-(aq)
7.4 Reactions That Form Water: Acids
and Bases
 Arrhenius Bases: substance that produces –OH ion
(hydroxide ions) in water
 Strong bases: Strong electrolytes
 Common strong bases: KOH, LiOH, NaOH, Ba(OH)2 and
Sr(OH)2
dissolved in H2O
NaOH(aq) ----------------- Na+(aq) + -OH(aq)
Writing Equations for Acid-Base
Reactions
 Nitric acid is a strong acid. Write the molecular, complete
ionic and net ion equations for the reaction of aqueous
nitric acid and aqueous potassium hydroxide.
 Hydrobromic acid is a strong acid. Write the molecular,
complete ionic and net ion equations for the reaction of
aqueous hydrobromic acid and aqueous barium hydroxide
7.5 Reactions of Metals with Nonmetals
(Oxidation Reduction)
 Oxidation – Reduction or Redox reaction: process of
transferring electrons from one to the other
 Step in determination Redox Reaction
 Determine the charges of each atom
 Oxidation: loss of electron(s); charge becomes more positive
 Reduction: gain of electron(s); charge becomes less positive
 E.g
2 Na(s) + Cl2(g)  2 NaCl(s)
Examples
 For each of the following reactions, show how electrons are
gained and lost
 2 Al(s) + 3 I2(g)  2 AlI3(s)
 2 Cs(s) + F2(g)  2 CsF(s)
 2SO2(g) + O2(g)  2SO3(g)
7.6 Ways to Classify Reactions
 Consider the driving forces
 Formation of solid


Precipitation reaction
Double replacement
 Formation of water

Acid – base reaction
 Transfer of electrons

Oxidation – Reduction reaction
7.7 Other Ways to Classify Reactions
 Combustion
 Organic Compound + O2(g) + heat  CO2(g) + H2O(g)
 Combination
 Element + element  compound
 Decomposition
 Compound  element + element
Examples
 Classify each of the following reactions in as many ways
as possible
 2 K(s) + Cl2(g)  2 KCl(s)
 BaCl2 (aq) + Na2SO4(aq)  BaSO4 (s) + 2 NaCl(aq)
 HNO3(aq) + NaOH(aq)  H2O(l) + NaNO3(aq)
 2C2H2 (g) + 5O2 (g)  4 CO2 (g) + 2H2O(l)
 PbO2 (s)  Pb(s) + O2 (g)
 Fe2O3(s) + 2 Al(s)  Al2O3(s) + 2 Fe(s)