Reactions in Aqueous
Solutions
Chapter 4
GENERAL PROPERTIES
Solution
A solution is a homogenous mixture of 2 or more
substances.
The solute is (are) the substance(s) present in the
smaller amount(s).
The solvent is the substance present in the larger
amount.
Solution
Solvent
Solute
Soft drink (l)
H2O
Sugar, CO2
Air (g)
N2
O2, Ar, CH4
Soft solder (s)
Pb
Sn
4
aqueous solutions of
KMnO4
Conduct electricity in solution?
Cations (+) and Anions (-)
ELECTROLYTE
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,
results in a solution that does not conduct electricity.
Hydration
process in which an ion is
surrounded by water
molecules arranged in a
specific manner.
d-
d+
H2O
Strong Electrolyte – 100% dissociation
NaCl (s)
H 2O
Na+ (aq) + Cl- (aq)
HCl(l)
H+(aq) + Cl−(aq)
CH3COOH(aq)
H+(aq) + CH3COO−(aq)
Weak Electrolyte – not completely dissociated
Review of Concepts
The diagrams here show three compounds AB2 (a), AC2 (b),
and AD2 (c) dissolved in water. Which is the strongest
electrolyte and which is the weakest? (For simplicity, water
molecules are not shown.)
PRECIPITATE REACTIONS
Precipitate – insoluble solid that separates from solution
CdS
PbS
Ni(OH)2
Al(OH)3
Solubility is the maximum amount of solute that will dissolve
in a given quantity of solvent at a specific temperature.
Example: 4.1
Classify the following ionic compounds as soluble or insoluble:
(a)
silver sulfate (Ag2SO4)
(b)
calcium carbonate (CaCO3)
(c)
sodium phosphate (Na3PO4).
precipitate
Pb(NO3)2 (aq) + 2NaI (aq)
PbI2 (s) + 2NaNO3 (aq)
molecular equation
Pb2+ + 2NO3- + 2Na+ + 2I-
PbI2 (s) + 2Na+ + 2NO3-
ionic equation
Pb2+ + 2IPbI2
PbI2 (s)
net ionic equation
Na+ and NO3- are spectator ions
15
Writing Net Ionic Equations
1. Write the balanced molecular equation.
2. Write the ionic equation showing the strong electrolytes
completely dissociated into cations and anions.
3. Cancel the spectator ions on both sides of the ionic
equation.
4. Check that charges and number of atoms are balanced in
the net ionic equation.
16
Example: 4.2
Predict what happens when a potassium phosphate (K3PO4)
solution is mixed with a calcium nitrate [Ca(NO3)2] solution.
Write a net ionic equation for the reaction.
Review of Concepts
Which of the diagrams here accurately describes the
reaction between Ca(NO3)2(aq) and Na2CO3(aq)? For
simplicity, only the Ca2+ (yellow) and CO32− (blue) ions are
shown.
ACID-BASE REACTIONS
General Properties
ACID
•
•
•
•
•
Sour taste
Color changes in plant
dyes
React with metals to
produce H2 gas
React with carbonates
and bicarbonates to
produce CO2 gas
Aqueous acid solutions
conduct electricity
BASE
•
•
•
•
Taste bitter
Feel slippery
Color changes in plant
dyes
Aqueous base solutions
conduct electricity
Arrhenius acid is a substance that produces H+ (H3O+) in water.
Arrhenius base is a substance that produces OH- in water.
21
Brønsted Acid and Bases
ACID
•
BASE
Proton donor
•
Proton acceptor
A Brønsted acid must contain at least one ionizable
proton!
base
acid
acid
base
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
23
STRONG ACIDS
STRONG BASES
HI
HBr
HClO4
HCl
H2SO4
HNO3
NaOH
KOH
LiOH
RbOH
CsOH
Ca(OH)2
Ba(OH)2
Strong acids/bases are strong electrolytes and
will completely
Sr(OH)
2
dissociate in water.
Example: 4.3
Classify each of the following species in aqueous solution as a
Brønsted acid or base:
(a)
(b)
(c)
HBr
Review of Concepts
Which of the following diagrams best represents a weak
acid? Very weak acid? Strong acid? The proton exists in
water as the hydronium ion. All acids are monoprotic. (For
simplicity, water molecules are not shown.)
Acid-Base Neutralization
acid + base
HCl (aq) + NaOH (aq)
salt + water
NaCl (aq) + H2O(l)
H+ + Cl- + Na+ + OH-
Na+ + Cl- + H2O (l)
H+ + OH-
H2O
Neutralization Reaction Involving a
Weak Electrolyte
weak acid + base
HCN (aq) + NaOH (aq)
HCN + Na+ + OH-
HCN + OH-
28
salt + water
NaCN (aq) + H2O
Na+ + CN- + H2O
CN- + H2O
Example: 4.4
Write molecular, ionic, and net ionic equations for each of the
following acid-base reactions:
(a) hydrobromic acid(aq) + barium hydroxide(aq)
(b) sulfuric acid(aq) + potassium hydroxide(aq)
Gas formation
• Certain salts react with acids to produce gaseous
products
• 2 HNO2 breaks down into H2O(l) + NO2(g) + NO(g)
• H2CO3 breaks down into H2O(l) + CO2(g)
• H2SO3 breaks down into H2O(l) + SO2(g)
• NH4OH breaks down into H2O(l) + NH3(g)
• H2S(g)
• CO2(g)
• H2(g)
• If you get one of these as a product in your molecular
equation, they immediately breakdown as above
• Gasses do not ionize
2HCl (aq) + Na2CO3 (aq)
2H+ + 2Cl- + 2Na+ + CO322H+ + CO32-
2NaCl (aq) + H2O +CO2
2Na+ + 2Cl- + H2O + CO2
H2O + CO2
Double Replacement Rxns
Review
Driving Force
How do you recognize it?
Precipitate
You must memorize the solubility rules. Any compound
formed from two ions can be recognized as soluble
(written as separate ions) or as a precipitate (written as
a molecule).
Gas formed
You must memorize the combinations that decompose
into gases (there are 4). You must also memorize the
gases that form. For example, when you H2SO3 as a
product, you must know it decomposes into H2O and
SO2 gas.
Weak electrolyte
You must memorize the short list of strong acids and
strong bases so you will recognize all the weak acids
and bases that dissolve, but do not dissociate into ions.
The weak base ammonia, NH3, is in this category. It
exits in water as NH3(aq) and only slightly forms the
ions NH4+ + OH−
2Mg + O2
2MgO
OXIDATION-REDUCTION
REACTIONS
OIL RIG
Half-reaction
OXIDATION REACTION
•
•
•
•
Reaction that involves the
loss of electrons
Contains reducing agentdonates electrons
Oxidation number becomes
more positive
Electrons are on the product
side of the half reaction
2Mg
2Mg2+ + 4e-
2Mg + O2 + 4e-
REDUCTION REACTION
•
•
•
•
Involves the gain of
electrons
Contains oxidizing agentaccepts electrons
Oxidation number becomes
more negative
Electrons are on the
reactant side of the half
reaction
O2 + 4e-
2Mg2+ + 2O2- + 4e-
2O2-
Zn(s) + CuSO4(aq)
Zn
Zn2+ + 2e-
ZnSO4(aq) + Cu(s)
Zn is oxidized
Zn is the reducing agent
Cu2+ + 2e-
Cu
Cu2+ is reduced
Cu2+ is the oxidizing agent
35
Oxidation Number
•
•
Charge the atom would have in a molecule if
electrons were transferred completely
Rules
•
•
•
•
Uncombined elements = 0
Neutral compounds sum = 0
Ion = ion charge (polyatomic ions sum to charge)
Exceptions
•
Hydrogen +1 w/ nonmetals, −1 w/ metals
•
Oxygen −2 except w/ fluorine (+2), in peroxides (−1)
•
Fluorine ALWAYS −1
Example: 4.5
Assign oxidation numbers to all the elements in the following
compounds and ion:
(a)
Li2O
(b)
HNO3
(c)
More common
oxidation numbers are
in red.
Types of Redox Reactions
Combination Reaction
A+B
0
C
+3 -1
0
2Al + 3Br2
2AlBr3
Decomposition Reaction
C
+1 +5 -2
2KClO3
39
A+B
+1 -1
0
2KCl + 3O2
Combustion Reaction
A + O2
B
0
0
S + O2
0
0
2Mg + O2
40
+4 -2
SO2
+2 -2
2MgO
Displacement Reaction
A + BC
0
+1
+2
Sr + 2H2O
+4
0
TiCl4 + 2Mg
0
-1
Cl2 + 2KBr
41
AC + B
0
Sr(OH)2 + H2
0
Hydrogen Displacement
+2
Ti + 2MgCl2
-1
Metal Displacement
0
2KCl + Br2
Halogen Displacement
Activity Series
For Halogens:
F2 > Cl2 > Br2 > I2
Disproportionation Reaction
The same element is simultaneously oxidized
and reduced.
Example:
reduced
+1
0
Cl2 + 2OHoxidized
43
-1
ClO- + Cl- + H2O
Elements most likely to
undergo disproportionation
Example: 4.6
Classify the following redox reactions and indicate changes in
the oxidation numbers of the elements:
(a)
(b)
(c)
(d)
Concentration
Molarity = moles of solute
liters of solution
Ex: 1M KCl solution
KCl(s)
H2O
K+(aq) + Cl−(aq)
Ex: 1M Ba(NO3)2 solution
Ba(NO3)2(s)
H2O
Ba2+(aq) + 2 NO3−(aq)
n
M=
V
Preparing a Solution of Known Concentration
47
Example: 4.7
How many grams of potassium
dichromate (K2Cr2O7) are
required to prepare a 250-mL
solution whose concentration
is 2.16 M?
A K2Cr2O7 solution.
Example: 4.8
In a biochemical assay, a chemist needs to add 3.81 g of
glucose to a reaction mixture. Calculate the volume in milliliters
of a 2.53 M glucose solution she should use for the addition.
Dilution
procedure for preparing a less concentrated solution from a
more concentrated solution.
Dilution
Add Solvent
Moles of solute
before dilution (i)
MiVi
=
=
Moles of solute
after dilution (f)
MfVf
Example: 4.9
Describe how you would prepare 5.00 × 102 mL of a
1.75 M H2SO4 solution, starting with an 8.61 M stock
solution of H2SO4.
Review of Concepts
What is the final concentration of a 0.6M NaCl
solution if its volume is doubled and the
number of moles of solute is tripled?
Quantitative analysis
•
•
Gravimetric analysis
Titrations
•
•
Acid-base
redox
Gravimetric Analysis
1. Dissolve unknown substance in water
2. React unknown with known substance to form a precipitate
3. Filter and dry precipitate
4. Weigh precipitate
5. Use chemical formula and mass of precipitate to determine
amount of unknown ion
Example: 4.10
A 0.5662-g sample of an ionic compound containing chloride
ions and an unknown metal is dissolved in water and treated
with an excess of AgNO3. If 1.0882 g of AgCl precipitate forms,
what is the percent by mass of Cl in the original compound?
Titrations
In a titration, a solution of accurately known concentration is
added 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
56
Titrations can be used in the analysis of
Acid-base reactions
H2SO4 + 2NaOH
2H2O + Na2SO4
Redox reactions
5Fe2+ + MnO4- + 8H+
57
Mn2+ + 5Fe3+ + 4H2O
Example: 4.11
In a titration experiment, a student finds that 23.48 mL of a
NaOH solution are needed to neutralize 0.5468 g of KHP. What
is the concentration (in molarity) of the NaOH solution?
Example: 4.12
How many milliliters (mL) of a 0.610 M NaOH solution are
needed to neutralize 20.0 mL of a 0.245 M H2SO4 solution?
Redox titrations
Example: 4.13
A 16.42-mL volume of 0.1327 M KMnO4
solution is needed to oxidize 25.00 mL of a
FeSO4 solution in an acidic medium. What is
the concentration of the FeSO4 solution in
molarity? The net ionic equation is