Reactions in solution
A subset of chemical reactions
Learning objectives
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Define solution and its components
Distinguish among strong, weak and non-electrolyte
Identify strong acids and strong bases
Apply solubility rules to prediction of precipitate
formation
Classify types of chemical reaction
Predict course of reaction based on activity series
Define oxidation and reduction
Identify oxidizing and reducing agent in reactions
Determine oxidation numbers in ions and compounds
Solution
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A homogeneous mixture of two or more
substances
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Not just limited to liquid state
Solutions may or may not contain electrolytes
Electrolytes are substances that conduct
electricity when dissolved
Electrolytes and ionic compounds
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(a) All ionic compounds are electrolytes when dissolved in water
Not all ionic compounds are soluble
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How do we tell?
Rules to predict solubility
(b) Covalent molecular compounds* are non-electrolytes – no ions
produced
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*Except acids and bases
Dissociation and ionization:
same or different?
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Ionic compounds dissociate in water
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Ions already exist in the solid
Acids or bases* ionize in water
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Pure acid (HCl) or base (NH3) contains no ions
*Except strong inorganic bases (NaOH, Ca(OH)2
etc) are ionic (Na+OH-)
When the weak are made strong:
(Actually they aren’t)
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Strong electrolytes are characterized by near
complete dissociation
Weak electrolytes dissociate to much smaller extent.
Strong, weak or non electrolyte?
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All soluble (ionic) salts are strong electrolytes
Strong acids and bases are strong electrolytes
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Weak acids and bases are weak electrolytes
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Insoluble compounds are non-electrolytes
Molecular compounds are non-electrolytes (except
acids/bases)
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Examples of electrolyte types
Strong
electrolytes
Weak
Nonelectrolytes electrolytes
MINERAL ACIDS:
HCl, HBr, HI
HClO4, HNO3, H2SO4
ACIDS:
HF, H3PO4,
CH3CO2H
SOLUBLE SALTS:
KBr, Na3PO4
SOLUBLE
SALTS:
None
MINERAL BASES:
NaOH, Ba(OH)2
BASES:
NH3
Neutral
molecular
covalent
compounds:
H2O,
CH3OH,
C12H22O11
(sucrose)
Most organic
compounds
and
INSOLUBLE
salts
Know your acids
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The six (seven) strong acids
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HCl, HBr, HI (but not HF)
HNO3 (but not HNO2)
H2SO4 (but not H2SO3)
HClO4 (maybe HClO3)
All other acids are weak
Recognizing acids
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Mineral (inorganic) acids: HCl, HNO3 etc.
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Conventionally H appears first in formula
May be strong or weak
All strong acids are mineral
Organic acids (contain C, H and O): CH3COOH etc
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Harder to spot
Sometimes written with H in front – HCH3CO2
Always weak
Presence of –OH (-SH): necessary but not sufficient
NOTE: Not all –OH are acidic (CH3OH is not an acid)
Recognizing bases
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Mineral (inorganic) bases usually
distinguished by OH- ions– all strong
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NaOH, Ca(OH)2
Ammonia, NH3, is an exception – is weak
Organic bases (derivatives of NH3 (CH3NH2
etc) do not contain –OH) – all weak
Classifying chemical reactions
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Acid-base reactions
Oxidation-reduction reactions
Number of reactants and products
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Combination reactions
Decomposition reactions
Single displacement reactions
Double displacement (metathesis)/ (partner
exchange) reactions (in solution)
Neutralization
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Combine acid with base:
ACID + BASE = SALT + WATER
HCl(aq) + NaOH(aq) = H2O(l) + NaCl(aq)
Mg(OH)2(s) + 2HCl(aq) = MgCl2(aq) + 2H2O(l)
Salt contains anion of acid and cation of base:
HCl + NaOH = NaCl + H2O
HCl + KOH = KCl + H2O
HNO3 + KOH = KNO3 + H2O
2HCl + Ca(OH)2 = CaCl2 + 2H2O
HCN + NaOH = NaCN + H2O
Acid-base reaction with gas
formation
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Tums...
HCl(aq) + NaHCO3(aq) = NaCl(aq) + H2CO3(aq)
H2CO3 is unstable:
H2CO3(aq) = H2O(l) + CO2(g)
Rewrite equation:
HCl(aq) + NaHCO3(aq) = NaCl(aq) + H2O(l) + CO2(g)
 Bad egg gas:
2HCl(aq) + Na2S(s) = H2S(g) + 2NaCl(aq)
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Oxidation - reduction
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Oxidation is loss of
electrons
Reduction is gain of
electrons
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Oxidation is always
accompanied by reduction
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The total number of electrons
is kept constant
Oxidizing agents oxidize
and are themselves reduced
Reducing agents reduce and
are themselves oxidized
Identifying oxidation and reduction
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Follow the electrons
Only reactants (things on left) need to be considered
Metal elements:
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Generally form positive ions (lose electrons)
Are reducing agents
Are oxidized
Na → Na+ + e-
Nonmetal elements:
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Form negative ions (gain electrons)
Are oxidizing agents
Are reduced
Cl2 + 2 e-→ 2 Cl-
Identifying oxidation and reduction
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With elements forming ionic compounds identifying
oxidation and reduction is usually straightforward
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Follow path of electrons from reactant to product
Metals are oxidized, nonmetals are reduced
What about covalent molecules and reactions
involving only compounds?
System of oxidation numbers is used
Oxidation numbers keep track of electrons
Oxidation numbers
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Oxidation number is the number of electrons gained
or lost by the element in making a compound
Metals are typically
considered more 'cation-like'
and would possess positive
oxidation numbers, while
nonmetals are considered
more 'anion-like' and would
possess negative oxidation
numbers.
Predicting oxidation numbers
1. Oxidation number of atoms in element is zero
2. Oxidation number of element in monatomic ion
equals charge
3. Sum of oxidation numbers in compound is zero
4. Sum of oxidation numbers in polyatomic ion equals
charge
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F has ON –1
H has ON +1; except in metal hydrides where it is –1
Oxygen is usually –2. Exceptions:
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O is –1 in hydrogen peroxide, and other peroxides
O is –1/2 in superoxides KO2
In OF2 O is +2
Position of element in periodic table
determines oxidation number
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G1A is +1
G2A is +2
G3A is +3 (some rare exceptions)
G5A are –3 in compounds with metals, H or with NH4+
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G6A below O (S, Se etc.) are –2 in binary compounds with
metals, H or NH4+
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Exceptions are compounds with elements to right (e.g. NO2, PF5);
in which case use rules 3 and 4.
When combined with O or lighter halogen (e.g. SeO2, SF6) use
rules 3 and 4.
G7A elements are –1 in binary compounds with metals, H
or NH4+ or with a heavier halogen (e.g. Cl in BrCl3)
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When combined with O or a lighter halogen, use rules 3 and 4 (e.g.
Br in BrCl3 or Cl in ClO4-).
Identifying reagents
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Those elements that tend to give up electrons (metals)
are typically categorized as reducing agents and those
that tend to accept electrons (nonmetals) are referred
to as oxidizing agents.
Identify redox by change in
oxidation numbers
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Reducing agent increases its oxidation number (Na)
Oxidizing agent decreases its oxidation number (H in
H2O)
Nuggets of redox processes
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Where there is oxidation there is always reduction
Oxidizing agent
Reducing agent
Is itself reduced
Is itself oxidized
Gains electrons
Loses electrons
Causes oxidation
Causes reduction
More active metals are strongly
reducing
Iron reduces Cu2+ to Cu
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Iron reduces Cu2+ ions to Cu
Cu does not reduce Fe2+
Applying activity series to metals in
acids
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Mg is higher than H in
activity series – forms
H2
Cu is lower than H in
activity series – no H2
produced
Element can be oxidizer and reducer
depending on relative positions in
activity series
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Fe reduces Cu2+
Cu reduces Ag+ (lower activity)
Fe2+ is reduced by Zn (higher activity)
Combination reactions
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Element + element  compound (redox)
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Metal + nonmetal  binary ionic compound
Nonmetal + nonmetal  binary covalent compound
Compound + element  compound (redox)
Compound + compound  compound
Decomposition reactions
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Compound  element + element (redox)
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Compound  element + compound (redox)
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Compound  compound + compound
Single replacement (displacement)
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Element displaces
another element from
compound (redox)
Metathesis (double displacement)
reactions involve changing partners
AX + BY = AY + BX
Driven by removal of ions from solution
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Formation of an insoluble solid (precipitate)
Formation of nonionized molecules (eg H2O)
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Acid-base neutralization
Formation of a gas (eg CO2)
Precipitation reactions
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Does one of the possible cation-anion
combinations produce an insoluble salt?
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Initial compounds are all soluble
Use solubility rules to investigate
If yes, a precipitate is produced
Solubility rools OK
Applied not remembered
Production of a gas
If product is a gas that
has low solubility in
water, reaction produces
gas
 Any carbonate with an
acid for example:
Na2CO3 + H2SO4 =
Na2SO4 + H2O + CO2
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