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Stoichiometry
College Chemistry
Chapter 3
Law of Conservation of Mass
 All chemical and physical reactions must follow the LCM.
 Lavoisier was the first to state this law and he was also given
credit as the founder of quantitative chemistry.
 All equations must be balanced in order to adhere to the
LCM.
Law of Conservation of Mass
 Regular reactions: atoms and masses balance
 Redox reactions: atoms, masses, and charges balance
 Here there will be change in the oxidation state of ions during the
reaction.
 One element will be oxidized; that means that it will lose electrons
and become more positive.
 One element will be reduced; that means that it will gain electrons
and become more negative.
 A balanced equation may require that half reactions be considered in
order to balance the charges.
Redox examples
 Mg0(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
Half reactions:
Oxidation: Mg0(s)  Mg+2 + 2eReduction: 2H+(aq) + 2e-  H20
Spectator ions: Cl-(aq)
Types of Reactions
 Synthesis (direct combination) – A+BC
 May use elements
 May use compounds
 Using metal oxides: get hydroxides
 Using nonmetal oxides: get acids
 Decomposition (analysis) – C  A+B
 Ordinary binary compounds
 Chlorates – chloride salt and oxygen
 Carbonates – oxides and carbon dioxide
 Hydroxides – metal oxide and water
 Acids – nonmetallic oxide and water
Examples
 Ordinary binary compounds
 2NaCl(s)  2Na0(s) + Cl20(g)
 Chlorates – chloride salt and oxygen
 2KClO3(s)  2KCl(s) + 3O2(g)
 Carbonates – oxides and carbon dioxide
 BaCO3(s) BaO(s) + CO2(g)
 Hydroxides – metal oxide and water
 NaOH(l)  Na2O(s) + H2O(l)
 Acids – nonmetallic oxide and water
 H2SO4(s)  SO3(g) + H2O(l)
Types of Reactions
 Single displacement – A + BC  AC + B
 Must use activity series to determine if reaction happens
 May be used with halogens with
F2 > Cl2 > Br2 > I2
 Double replacement – AB + CD  AD + CB
 Must consider the driving forces to determine if reaction happens
 Formation of a precipitate
 Formation of a gas
 Formation of water
 Formation of a small molecular compound
Solubility Rules
 Formation of a precipitate occurs if the product is insoluble in
aqueous solution.
 Soluble compounds are those containing
 NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1
 NO3 C2H3O2- except with Ag+
 Cl-, Br-, I- except with Ag+, Hg2+2, and Pb+2
 SO4-2 except with Sr+2, Ba+2, Hg2+2, Pb+2
Solubility Rules
 Insoluble compounds are those containing
 S-2 except with NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1, Sr+2, Ba+2, Ca+2
 CO3-2 except with NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1
 PO4-3 except with NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1
 OH- except with NH4+, Li+1, Na+1, K+1, Rb+1, Cs+1, Sr+2, Ba+2, Ca+2
REMEMBER
 Diatomics such as H2, O2, N2, F2, Cl2, Br2, or I2
 The charge on a compound must net zero.
 The charges on the representative elements are predictable;
those on the transition elements are not except for zinc,
silver, and cadmium.
Types of Reactions
 Combustion – one reactant will be oxygen and the other one
may be an organic compound, a metallic compound, or a
nonmetallic compound
 Complete – occurs when there is excess oxygen  CO2(g) +
H2O(l)
 Incomplete – occurs when there is a limited amount of oxygen
 CO(g) + H2O(l)
Types of Solution Reactions
 Precipitation reactions – a precipitate forms as a result of
two solutions reacting in solution
 Acid-base reactions (neutralizations) – water and a salt form
as a result of an acid and a base reacting in solution (may be a
titration)
 Oxidation-reduction reactions (redox) – products form such
that the charges on reactant ions change as the reaction
proceeds.
 Two half reactions are sometimes required to determine the
balanced equation.
 May be in acidic or basic solution.
Atomic and Molecular Weights
 Atomic masses are based on the standard of carbon-12.
 The masses on the periodic table are not the mass of any
isotopes; they are the weighted average of all the isotopes.
Empirical Formulas
 Directly
 Masses
 Percent composition
 Combustion analysis