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Reactions in Aqueous Solution
Chapter 4
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A solution is a homogenous mixture of 2 or more
substances
The solvent is the substance present in the larger
amount
The solute is(are) the substance(s) present in the
smaller amount(s)
Examples:
Solution
Solvent
Solute
Soft drink (l)
H2O
Sugar, CO2
Air (g)
N2
O2, Ar, CH4
Soft Solder (s)
Pb
Sn
4.1
An electrolyte is a substance that, when dissolved in
water, results in a solution that can conduct electricity.
A nonelectrolyte is a substance that, when dissolved
in water, results in a solution that does not conduct
electricity.
nonelectrolyte
weak electrolyte
strong electrolyte
4.1
Conduct electricity in solution?
Cations (+) and Anions (-)
Strong Electrolyte – 100% dissociation
NaCl (s)
H 2O
Na+ (aq) + Cl- (aq)
Weak Electrolyte – not completely dissociated
CH3COOH
CH3COO- (aq) + H+ (aq)
4.1
H2O
d-
A Polar molecule
+
d
Hydration is the process in which an ion is surrounded
by water molecules arranged in a specific manner.
d-
d+
H2O
It’s a battle.
Charge!
Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution
C6H12O6 (s)
H 2O
C6H12O6 (aq)
No dissociation
Examples:
Precipitation Reactions
Precipitate – insoluble solid that separates from solution
PbI2
4.2
Memorization time…
Solubility Rules for Common Ionic Compounds In water at 25oC
1. All alkali metals (Group 1A) compounds are soluble.
2. All ammonium (NH4+) compounds are soluble.
3. All nitrate (NO3-), chlorate (ClO3-), and perchlorate (ClO4-) compounds
are soluble.
4. Most hydroxides (OH-) are insoluble. The exceptions are barium
hydroxide [Ba(OH)2], which is very soluble, calcium hydroxide
[Ca(OH)2], which is slightly soluble, and previous examples.
5. Most compounds containing chlorides (Cl-), bromides (Br-), or iodides
(I-), are soluble. The exceptions are those containing Ag+, Hg22+, and
Pb2+.
6. Most sulfates (SO42-) are soluble, excepting previous rules. Calcium
sulfate [CaSO4] and silver sulfate [Ag2SO4] are slightly soluble.
Barium sulfate [BaSO4], mercury (II) sulfate [HgSO4], and lead sulfate
[PbSO4] are insoluble.
7. All carbonates (CO32-), phosphates (PO43-), and sulfides (S2-) are
insoluble, excepting previous rules.
Solubility Rules for Common Ionic Compounds
In water at 250C
Precipitation Reactions
Precipitate – insoluble solid that separates from solution
precipitate
Pb(NO3)2 (aq) + 2NaI (aq)
PbI2 (s) + 2NaNO3 (aq)
molecular equation
Pb2+ + 2NO3- + 2Na+ + 2I-
PbI2 (s) + 2Na+ + 2NO3-
ionic equation
Na+ and NO3- are spectator ions
PbI2
Pb2+ + 2I-
PbI2 (s)
net ionic equation
4.2
Writing Net Ionic Equations
1. Write the balanced molecular equation.
2. Write the ionic equation showing the strong electrolytes
3. Determine precipitate from solubility rules
4. Cancel the spectator ions on both sides of the ionic equation
Write the net ionic equation for the reaction of silver
nitrate with sodium chloride.
AgNO3 (aq) + NaCl (aq)
AgCl (s) + NaNO3 (aq)
Ag+ + NO3- + Na+ + Cl-
AgCl (s) + Na+ + NO3-
Ag+ + Cl-
AgCl (s)
4.2
Acids (according to Arrhenius)
• Have a sour taste. Vinegar owes its taste to acetic acid.
Citrus fruits contain citric acid.
• Cause certain plant dyes to change color. For example,
they turn litmus from blue to red.
• React with certain metals to produce hydrogen gas.
• React with carbonates (CO32-) and bicarbonates (HCO3-) to
produce carbon dioxide gas.
• Are electrolytes.
4.3
Bases (according to Arrhenius)
• Have a bitter taste.
• Cause certain plant dyes to change color. For example,
they turn litmus from red to blue.
• Feel slippery (they make water insoluble organic molecules
into water soluble molecules). Many soaps contain bases.
• Are electrolytes.
4.3
Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
4.3
A Brønsted acid is a proton donor
A Brønsted base is a proton acceptor
base
acid
acid
base
A Brønsted acid must contain at least one
ionizable proton!
4.3
Monoprotic acids
HCl
H+ + Cl-
HNO3
H+ + NO3H+ + CH3COO-
CH3COOH
Strong electrolyte, strong acid
Strong electrolyte, strong acid
Weak electrolyte, weak acid
Diprotic acids
H2SO4
H+ + HSO4-
Strong electrolyte, strong acid
HSO4-
H+ + SO42-
Weak electrolyte, weak acid
Triprotic acids
H3PO4
H2PO4HPO42-
H+ + H2PO4H+ + HPO42H+ + PO43-
Weak electrolyte, weak acid
Weak electrolyte, weak acid
Weak electrolyte, weak acid
4.3
Neutralization Reaction
acid + base
HCl (aq) + NaOH (aq)
H+ + Cl- + Na+ + OH-
H+ + OH-
salt + water
NaCl (aq) + H2O
Na+ + Cl- + H2O
H2O
net ionic equation for strong acid / strong base reaction
4.3
Neutralization Reaction
acid + base
salt + water
HCl (aq) + NH3 (aq)
NH4Cl(aq)
H+ + Cl- + NH3(aq)
NH4+ + Cl-
H+ + NH3(aq)
NH4+
net ionic equation for strong acid / weak base reaction
4.3
Oxidation-Reduction Reactions
(electron transfer reactions)
2Mg (s) + O2 (g)
2Mg
O2 + 4e-
2MgO (s)
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
2O2-
Reduction half-reaction (gain e-)
2Mg + O2 + 4e2Mg + O2
2Mg2+ + 2O2- + 4e2MgO
4.4
Fig. 4.9a-c
Iron (Fe) reacting with hydrochloric acid (HCl)
Fe (s) + 2HCl (aq)
Fe
Fe2+ + 2e- Fe is oxidized
2H+ + 2e-
H2 H+ is reduced
FeCl2 (aq) + H2 (g)
Fe is the reducing agent
H+ is the oxidizing agent
4.4
The Activity Series for Metals
Displacement Reaction
M + BC
AC + B
M is metal
BC is acid or H2O
B is H2
Ca + 2H2O
Ca(OH)2 + H2
Pb + 2H2O
Pb(OH)2 + H2
4.4
4.4
Zn (s) + CuSO4 (aq)
Zn2+ + 2e- Zn is oxidized
Zn
Cu2+ + 2e-
ZnSO4 (aq) + Cu (s)
Zn is the reducing agent
Cu Cu2+ is reduced Cu2+ is the oxidizing agent
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Cu (s) + 2AgNO3 (aq)
Cu
Ag+ + 1e-
Cu2+ + 2eAg
Cu(NO3)2 (aq) + 2Ag (s)
Cu is oxidized Cu is the reducing agent
Ag+ is reduced Ag+ is the oxidizing agent
4.4
Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation
number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2. In monatomic ions, the oxidation number is equal to
the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is usually –2. In H2O2
and O22- it is –1.
4.4
4. The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds. In these
cases, its oxidation number is –1.
5. Group IA metals are +1, IIA metals are +2 and fluorine is
always –1.
6. The sum of the oxidation numbers of all the atoms in a
molecule or ion is equal to the charge on the
molecule or ion.
HCO3Oxidation numbers of all
the elements in HCO3- ?
O = -2
H = +1
3x(-2) + 1 + ? = -1
C = +4
4.4
Fig. 4.10
IF3
Oxidation numbers of all
the elements in the
following ?
F = -1
3x(-1) + ? = 0
I = +3
NaIO3
Na = +1 O = -2
3x(-2) + 1 + ? = 0
I = +5
K2Cr2O7
O = -2
K = +1
7x(-2) + 2x(+1) + 2x(?) = 0
Cr = +6
4.4
Types of Oxidation-Reduction Reactions
Combination Reaction
A+B
C
0
+4 -2
0
S + O2
SO2
Decomposition Reaction
C
+1 +5 -2
2KClO3
A+B
+1 -1
0
2KCl + 3O2
4.4
Types of Oxidation-Reduction Reactions
Displacement Reaction
A + BC
0
+1
+2
Sr + 2H2O
+4
0
TiCl4 + 2Mg
0
AC + B
-1
Cl2 + 2KBr
0
Sr(OH)2 + H2 Hydrogen Displacement
0
+2
Ti + 2MgCl2
-1
Metal Displacement
0
2KCl + Br2
Halogen Displacement
4.4
The Activity Series for Metals
Displacement Reaction
M + BC
AC + B
M is metal
BC is acid or H2O
B is H2
Ca + 2H2O
Ca(OH)2 + H2
Pb + 2H2O
Pb(OH)2 + H2
4.4
Classify the following reactions.
Ca2+ + CO32NH3 + H+
Zn + 2HCl
Ca + F2
CaCO3
NH4+
ZnCl2 + H2
CaF2
Precipitation
Acid-Base
Redox (H2 Displacement)
Redox (Combination)
4.4
Solution Stoichiometry
The concentration of a solution is the amount of solute
present in a given quantity of solvent or solution.
M = molarity =
moles of solute
liters of solution
What mass of KI is required to make 500. mL of
a 2.80 M KI solution?
M KI
volume KI
500. mL x
moles KI
1L
1000 mL
x
2.80 mol KI
1 L soln
M KI
x
grams KI
166 g KI
1 mol KI
= 232 g KI
4.5
4.5
Dilution is the procedure for preparing a less concentrated
solution from a more concentrated solution.
Dilution
Add Solvent
Moles of solute
before dilution (i)
=
Moles of solute
after dilution (f)
MiVi
=
MfVf
4.5
How would you prepare 60.0 mL of 0.200 M
HNO3 from a stock solution of 4.00 M HNO3?
MiVi = MfVf
Mi = 4.00
Vi =
MfVf
Mi
Mf = 0.200
Vf = 0.0600 L
Vi = ? L
0.200 x 0.0600
=
= 0.00300 L = 3.00 mL
4.00
3.00 mL of acid + 57.0 mL of water = 60.0 mL of solution
4.5
Titrations
In a titration a solution of accurately known concentration is
gradually added to another solution of unknown concentration
until the chemical reaction between the two solutions is
complete.
Equivalence point – the point at which the reaction is complete
Indicator – substance that changes color at (or near) the
equivalence point
Slowly add base
to unknown acid
UNTIL
the indicator
changes color
4.7
What volume of a 1.420 M NaOH solution is
Required to titrate 25.00 mL of a 4.50 M H2SO4
solution?
WRITE THE CHEMICAL EQUATION!
H2SO4 + 2NaOH
M
volume acid
25.00 mL x
acid
2H2O + Na2SO4
rx
moles acid
4.50 mol H2SO4
1000 mL soln
x
coef.
M
moles base
2 mol NaOH
1 mol H2SO4
x
base
volume base
1000 ml soln
1.420 mol NaOH
= 158 mL
4.7
What volume of a 0.800 M HCl solution is
required to titrate 40.00 mL of a 1.600 M NH3
solution?
HCl + NH3
40.00 mL x
1.600 mol NH3
1000 mL soln
x
1 mol HCl
1 mol NH3
NH4Cl
x
1000 ml soln
= 80.0 mL
0.800 mol HCl
4.7
What volume of a 1.200 M Ba(OH)2 solution is
required to titrate 90.00 mL of a 3.600 M H3PO4
solution?
3Ba(OH)2 + 2H3PO4
6H2O + Ba3(PO4)2
3.600 mol H3PO4 3 mol Ba2(OH)
1000 ml soln
= 405.0 mL
x
x
90.00 mL x
1000 mL soln
2 mol H3PO4 1.200 mol Ba2(OH)
4.7