4. Reactions in Aqueous Solutions
solution: homogeneous mixture of two or more substances
solvent: bulk medium; usually the substance present with the
largest amount
solute(s): dissolved substance(s); usually present in smaller
amounts
examples:
gaseous – air
solid – alloy (brass: Cu/Zn)
liquid – seawater
aqueous solutions
solvent = water
water is inexpensive; is able to dissolve a large number of substances
many substances dissociate into ions in water =⇒ conduction
of electricity
pure water does not conduct electricity
water is an effective solvent for ionic compounds
GChem I
4.1
“−”
hydration
H
.
.....
.....
.....
.....
.
.
.
..
O
.....
.....
.....
.....
.....
.
“+”
H
aqueous solution:
water
nonelectrolyte: no ions present; no electric current
electrolyte:
weak electrolyte: substance is partially ionized: fair conductor
strong electrolyte: substance is completely ionized: good conductor
soluble ionic compounds are strong electrolytes
most molecular compounds are nonelectrolytes or weak electrolytes
examples:
MgCl2 (aq)
Mg2+ (aq) + 2 Cl− (aq)
−−−−−−→
ionic compound
HCl(aq)
strong electrolyte
H+ (aq) + Cl− (aq)
−−−−−−→
molecular compound
HC2 H3 O2 (aq)
strong electrolyte (acid)
−
−
*
)
−−
−−
−−
−−
−−
−−
−
−
reversible
molecular compound
– Typeset by FoilT X –
E
H+ (aq) + C2 H3 O2 − (aq)
weak electrolyte (acid)
dynamic equilibrium
here: HC2 H3 O2 predominant species
GChem I
4.2
C2 H5 OH(aq)
−−−
−−−→
molecular compound
nonelectrolyte (alcohol)
three important types of reactions in aqueous solutions:
(I) precipitation reactions
cation(aq) + anion(aq)
(beaker 1)
−−−−→
insoluble ionic solid
(beaker 2)
(precipitate)
net ionic equation
example
AgNO3 (aq) + NaI(aq) −→ AgI(s) + NaNO3 (aq)
questionable form: strong electrolytes present
better:
Ag+ (aq)+NO3 − (aq)+Na+ (aq)+I− (aq) −→ AgI(s)+Na+ (aq)+NO3 − (aq)
ionic form
+
−
“spectator ions”: Na (aq), NO3 (aq)
XXX
XXX
XX
XX
−
+X
−
+
−
XXX + NO
(aq)
Ag+ (aq) + NO
(aq)
Na
Na
(aq)
(aq)
X
XX + I (aq) −→ AgI(s) + XX + XX
3 X
3
X
X
X
X
Ag+ (aq) + I− (aq) −→ AgI(s) net ionic equation
GChem I
4.3
net ionic equations are doubly balanced
(i) atoms
(ii) electric charge
predict outcome of a reaction; for example:
AgNO3 (aq) + KBr(aq) −→ ?
Ag+ (aq) + NO3 − (aq) + K+ (aq) + Br− (aq) −→ ?
ionic compound is “insoluble”, if only a small amount can be
dissolved, say < 0.01 mol/L
Solubility of Ionic Compounds: Solubility Rules
Classes of soluble compounds
• compounds of the alkali metals (1A)
• ammonium (NH4 + ) compounds
• nitrates (NO3 − ), chlorates (ClO3 − ), perchlorates (ClO4 − ), ac−
etates (CH3 CO2 )
• chlorides (Cl− ), bromides (Br− ), iodides (I− ), except: Pb2+ ,
Ag+ , Hg2 2+
• sulfates (SO4 2− ), except: Sr2+ , Ba2+ , Pb2+ , Ca2+ , Hg2 2+ , Ag+
(the last three are slightly soluble)
GChem I
4.4
Classes of insoluble compounds
• carbonates (CO3 2− ), phosphates (PO4 3− ), oxalates (C2 O4 2− ),
2−
+
chromates (CrO4 ), except: alkali metals, NH4
• sulfides (S2− ), except: alkali metals, NH4 + , alkaline earth
metals (2A) (CaS, SrS, and BaS are slightly to moderately soluble),
• hydroxides (OH− ), oxides (O2− ), except: alkali metals, alkaline earth metals (Ca(OH)2 and Sr(OH)2 are only slightly
soluble; Mg(OH)2 is only very slightly soluble)
AgNO3 (aq) + KBr(aq) −→ ?
Ag+ (aq) + NO3 − (aq) + K+ (aq) + Br− (aq) −→ ?
Ag+ (aq) + Br− (aq) −→ AgBr(s)
(II) acid-base reactions
+
Arrhenius definition: acids provide H ions in aqueous solution;
−
bases provide OH ions in aqueous solution
acids:
• taste sour; acetic acid in vinegar, citric acid in lemon juice
GChem I
4.5
• cause color changes in plant dyes: litmus blue −→ red
• react with metals (e.g., Zn, Mg, Fe) to produce H2 gas
2 HCl(aq) + Mg(s) −→ MgCl2 (aq) + H2 (g)
• react with carbonates and hydrogen carbonates (bicarbonates) to produce CO2 gas
2 HCl(aq) + CaCO3 (s) −→ CaCl2 (aq) + H2 O(l) + CO2 (g)
HCl(aq) + NaHCO3 (s) −→ NaCl(aq) + H2 O(l) + CO2 (g)
• electrolytes
bases:
• taste bitter
• feel slippery (e.g., soaps)
• cause color changes in plant dyes; litmus red −→ blue
• electrolytes
Arrhenius definition limited to aqueous solutions
broader definition: Brønsted (Brønsted-Lowry)
+
acid = proton (H ) donor
base = proton acceptor
GChem I
4.6
examples of Brønsted acids:
HCl(g) + water −→ HCl(aq) −→ H+ (aq) + Cl− (aq)
more correctly
HCl(aq) + H2 O(l) −→ H3 O+ (aq) + Cl− (aq)
H3 O+ hydronium ion
HNO3 (aq) −→ H+ (aq) + NO3 − (aq)
+
examples of monoprotic acids: donate one H upon ionization
examples of strong acids: completely ionized in solution (strong
electrolytes); reaction goes to completion: −→
+
−
CH3 COOH(aq) −
)−
−*
− H (aq) + CH3 COO (aq)
monoprotic acid; example of a weak acid: equilibrium between
dissociated and undissociated form (weak electrolyte)
+
diprotic acids donate two H upon ionization
H2 SO4 −−→ H+ (aq) + HSO4 − (aq)
2−
+
HSO4 − (aq) −
)−
−*
− H (aq) + SO4 (aq)
GChem I
4.7
+
triprotic acids donate three H upon ionization
−
+
H3 PO4 −
)−
−*
− H (aq) + H2 PO4 (aq)
+
2−
H2 PO4 − (aq) −
)−
−*
− H (aq) + HPO4 (aq)
3−
+
HPO4 2− (aq) −
)−
−*
− H (aq) + PO4 (aq)
examples of Brønsted bases:
H2 O(l)
NaOH(s) −−−−−−→ Na+ (aq) + OH− (aq)
OH− (aq) + H+ (aq) −−→ H2 O(l)
NaOH is an example of a strong base: completely ionized in solution (strong electrolyte)
+
−
NH3 (aq) + H2 O(l) −
)−
−*
− NH4 (aq) + OH (aq)
NH3 is an example of a weak base: equilibrium between ionized
and nonionized form (weak electrolyte)
GChem I
4.8
Common Strong Acids and Strong Bases
Acids: HCl, HBr, HI, HClO4 , HNO3 , H2 SO4
(H2 SO4 ionizes in two
distinct steps. It is a strong acid only in its first ionization step.)
Bases:
LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2 , Sr(OH)2 ,
Ba(OH)2
neutralization reaction
acid + base −−→ salt + water
example:
HCl(aq) + NaOH(aq) −−→ NaCl(aq) + H2 O(l)
better: ionic form
H+ (aq)+Cl− (aq)+Na+ (aq)+OH− (aq) −−→ Na+ (aq)+Cl− (aq)+H2 O(l)
net ionic equation
H+ (aq) + OH− (aq) −−→ H2 O(l)
acid-base reactions are proton transfer reactions
GChem I
4.9
Some Common Gas-Forming Reactions
HSO3 − (aq) + H+ (aq) −→ SO2 (g) + H2 O(l)
SO3 2− (aq) + 2 H+ (aq) −→ SO2 (g) + H2 O(l)
HCO3 − (aq) + H+ (aq) −→ CO2 (g) + H2 O(l)
CO3 2− (aq) + 2 H+ (aq) −→ CO2 (g) + H2 O(l)
S2− (aq) + 2 H+ (aq) −→ H2 S(g)
NH4 + (aq) + OH− (aq) −→ NH3 (g) + H2 O(l)
(III) oxidation-reduction reactions
2 Mg(s) + O2 (g) −→ 2 MgO(s)
CH4 (g) + 2 O2 (g) −→ CO2 (g) + 2 H2 O(l)
∆
Fe2 O3 (s) + 3 CO(g) −→ 2 Fe(l) + 3 CO2 (g)
historical definition: oxidation = gain of O; reduction = loss of O
Mg(s) + Cl2 (g) −→ MgCl2 (s)
compare with 2 Mg(s) + O2 (g) −→ 2 MgO(s)
GChem I
4.10
decompose the reactions into half-reactions:
2 Mg −→ 2 Mg2+ + 4 e− oxidation half-reaction; electrons on the right
O2 + 4 e− −→ 2 O2− reduction half-reaction; electrons on the left
Mg −→ Mg2+ + 2 e− oxidation half-reaction
Cl2 + 2 e− −→ 2 Cl− reduction half-reaction
modern view: oxidation-reduction reactions or redox reactions
are electron transfer reactions
oxidation = loss of electrons; reduction = gain of electrons
oxidation and reduction always occur together; there is no net
loss or gain of electrons overall
works well for ionic compounds
generalize to all compounds, including molecular compounds,
by assigning oxidation numbers to the elements in the reactants
and products
an atom’s oxidation number, also called oxidation state, represents the number of charges the atom would have in a molecule or ionic compound if electrons were transferred completely
0
0
+2 −2
2 Mg(s) + O2 (g) −→ 2 Mg O(s)
GChem I
4.11
0
0
+2 −1
Mg(s) + Cl2 (g) −→ MgCl2 (s)
0
0
+1 −1
H2 (g) + Cl2 (g) −→ 2 H Cl(g)
0
0
+4−2
S(s) + O2 (g) −→ SO 2 (g)
Mg : 0 −→ +2
H : 0 −→ +1
S : 0 −→ +4
O : 0 −→ −2
Cl : 0 −→ −1
oxidation: oxidation state of an atom increases
reduction: oxidation state of an atom decreases
redox reaction = oxidation half-reaction + reduction halfreaction
redox reactions in aqueous solution
Zn(s) + 2 HCl(aq) −→ ZnCl2 (aq) + H2 (g)
GChem I
4.12
ionic equation
Zn(s) + 2 H+ (aq) + 2 Cl− (aq) −→ Zn2+ (aq) + 2 Cl− (aq) + H2 (g)
net ionic equation
Zn(s) + 2 H+ (aq) −→ Zn2+ (aq) + H2 (g)
determine oxidation states
0
+1
+2
+
2+
0
Zn(s) + 2 H (aq) −→ Zn (aq) + H2 (g)
Zn is oxidized; H+ is reduced
ox:
Zn(s) −→ Zn2+ (aq) + 2 e−
+
−
red: 2 H (aq) + 2 e −→ H2 (g)
Zn(s) + 2 H+ (aq) −→ Zn2+ (aq) + H2 (g)
oxidizing agent or oxidant: the substance that is reduced; in the
+
previous example: H
reducing agent or reductant: the substance that is oxidized; in
the previous example: Zn
types of redox reactions
GChem I
4.13
(1) combination reactions
A + B −→ C
0
0
+4−2
S(s) + O2 (g) −→ SO 2 (g)
0
0
+2
−3
3 Mg(s) + N2 (g) −→ Mg3 N 2 (s)
(2) decomposition reactions
C −→ A + B
0
+2 −2
0
2 Hg O(s) −→ 2 Hg(l) + O2 (g)
+1 +5 −2
+1 −1
0
2 K Cl O 3 (s) −→ 2 K Cl(s) + 3 O2 (g)
(3) displacement reactions
A + BC −→ AC + B
(i) hydrogen displacement
active metals displace hydrogen from cold water
0
+1 −2
+1 −2+1
0
2 Na(s) + 2 H 2 O(l) −→ 2 Na O H(aq) + H2 (g)
GChem I
4.14
0
+1 −2
0
+2 −2+1
Ca(s) + 2 H 2 O(l) −→ Ca( O H)2 (s) + H2 (g)
less reactive metals displace hydrogen from steam
0
+1 −2
+3 −2
0
2 Al(s) + 3 H 2 O(g) −→ Al2 O 3 (s) + 3 H2 (g)
many metals can displace hydrogen from acids
0
+1 −1
+2 −1
0
Zn(s) + 2 H Cl(aq) −→ ZnCl2 (aq) + H2 (g)
0
+1
0
+2
Zn(s) + 2 H + (aq) −→ Zn2+ (aq) + H2 (g)
activity series
(ii) metal displacement
0
+1
0
+2
Cu(s) + 2 AgNO3 (aq) −→ Cu(NO3 )2 + 2 Ag(s)
0
+2
+2
0
Zn(s) + CuSO4 (aq) −→ ZnSO4 (aq) + Cu(s)
activity series
any species will react with any species (in a compound) below it
in the activity series
(iii) halogen displacement
F2 > Cl2 > Br2 > I2
GChem I
4.15
0
+1 −1
+1 −1
0
Cl2 (g) + 2 K Br(aq) −→ 2 K Cl(aq) + Br2 (l)
0
+1−1
0
+1 −1
Br2 (g) + 2 K I (aq) −→ 2 K Br(aq) + I2 (l)
(4) disproportionation reactions
simultaneous oxidation and reduction of an element with multiple oxidation states
+1 −1
+1 −2
0
2 H 2 O 2 (aq) −→ 2 H 2 O(l) + O2 (g)
0
−2+1
−
+1 −2
−
−1
−
+1 −2
Cl2 (g) + 2 O H (aq) −→ Cl O (aq) + Cl (aq) + H 2 O(l)
concentration of solutions
concentration = amount of solute in a given quantity of solvent
or solution
molarity =
GChem I
amount of solute (in moles)
volume of solution (in L)
4.16
unit: M = mol/L
c=
molar
n
V
molarity refers to the amount of solute originally dissolved
1M glucose solution: 1 mole C6 H12 O6 molecules in 1L of solution
1M KCl solution: 1 mole KCl dissolved in enough water to provide
1 L of solution
KCl is a strong electrolyte:
KCl(aq) −→ K+ (aq) + Cl− (aq)
+
+
therefore 1 L of solution contains one mole K ions, [K ] = 1 M,
−
−
and one mole Cl ions, [Cl ] = 1 M, i.e., a total of two moles of
ions
dilution of solutions
stock solution (concentrated solution) −→ dilute solution
c i −→ c f
c i Vi = c f V f
GChem I
4.17