Chapter 16 – Acid Base Equlibria #3
16.9 Acid-Base Properties of Salt Solutions
All soluble salts are strong electrolytes.
•
- In solution, they exist nearly entirely of ions.
- Acid-base properties of salts are due to the reactions of their ions in solution.
•
Many ions can react with water to form OH- or H+
•
This process is called hydrolysis.
Anions
–
–
–
Cations
–
–
•
Of strong acids are neutral.
• Example: Cl- (of HCl)
– Cl- + H2O  X
Of weak acids are basic.
• Example: F- of HF
– F- + H2O  HF + OHWith ionizable protons are amphoteric.
•
Example: HSO4Of strong bases are neutral.
• Example: Na+ of NaOH
– Na+ + H2O  X
All other cations are weak acids
• Example: Fe3+
Cation hydrolysis reaction:
Smaller and more highly charged ions = stronger (weak) acids
The pH of a solution may be qualitatively predicted:
Cation
Anion
Solution is:
Example
Neutral
(from strong base)
Neutral
(from strong acid)
NaCl
Neutral
(from strong base)
Basic
(from weak acid)
NaF
Acidic
(from a weak base)
Neutral
(from a strong acid)
FeCl3
Acidic
(from a weak base)
Basic
(from a weak acid)
NH CN
NH4Cl
4
Salts from a weak acid and weak base can be either acidic or basic.
Compare K of the cation and K of the anion
a
b
- If the K is larger the solution will be acidic
a
- If the K is larger the solution will be basic
b
For example, consider NH4CN.
The K of CN- is larger than the K of NH4+ so the solution will be basic.
b
a
Example 1: Predict whether aqueous solutions of the following compounds are acidic, basic or neutral.
(a) NH Br
(b) FeCl
(c) Na CO
4
3
d) KClO
2
3
(e) NaHC O
4
2
4
-1
Example 2: Using data from Appendix D, calculate [OH ] and pH for the following solution 0.10 M NaCN
NaCN is a soluble salt and will dissociate into Na +1 and CN -1 ions.
Na +1 ion does not interact with water, but CN -1 acts as a weak base:
Ka for HCN (given in Appendix D) = 4.9x10
-10
16.10 Acid-Base Behavior and Chemical Structure
Acidity is directly related to the strength of attraction for a pair of electrons to a central atom.
4 situations to consider:
1. Ions
Ionic Charge and Size
When comparing ions of similar structure:
More positive ions are stronger acids.
tie breaker: Smaller ion is stronger acid
+
Acid strength:
2+
2+
3+
Na < Ca < Cu < Al
3-
PO
4
2-
< HPO
4
< H PO - < H PO
2
4
3
4
Example 3: Predict which member of each pair produces the more acidic aqueous solution:
+1
(a) K
+2
or Cu
+2
(b) Fe
+3
or Fe
+3
(c) Al
+3
or Ga
2. Binary Acids:
•
•
Bond Polarity (Electronegativity) & Strength
The H–X bond strength is important in determining relative acid strength in any group in the periodic table.
– The weaker the bond the easier it will break
– The H–X bond strength tends to decrease down a group - acid strength increases down a group
H–X bond polarity is important in determining relative acid strength in any period of the periodic table.
– The H-X bond polarity tends to increase across a period - acid strength increases (from left to right)
across a period
3. Oxyacids (Acids with oxygen) with different central atoms
Generally, the larger the electronegativity of the central atom the stronger the acid.
– The stronger the pull on electrons the less tightly the H is held
Acid Strength: H BO < H CO < HNO
3
3
2
3
3
The higher EN of central atom means more electron density shift away from H - H is easier to remove – stronger acid
4. Oxyacids with the same central atom
Generally, the more oxygens attached to the central atom the stronger the acid.
– The more atoms pulling on electrons the less tightly the H is held
Acid Strength: HClO < HClO < HClO < HClO
2
3
4
More O atoms means more electron density shift away from H and H is easier to remove – stronger acid
Example 4:
Explain the following observations:
(a) HNO is a stronger acid than HNO
3
2
(b) H S is a stronger acid than H O
2
2
(c) H SO is a stronger acid than HSO42
4
(d) H SO is a stronger acid than H SeO
2
4
2
4
(e) CCl COOH is a stronger acid than CH COOH
3
3
16.11 Lewis Acids and Bases
A Brønsted-Lowry acid is a proton donor.
Focusing on electrons: A Brønsted-Lowry acid can be considered as an electron pair acceptor.
Lewis proposed a new definition of acids and bases that emphasizes the shared electron pair.
A Lewis acid is an electron pair acceptor.
A Lewis base is an electron pair donor.
Note: Lewis acids and bases do not need to contain protons.
Therefore, the Lewis definition is the most general definition of acids and bases.
What types of compounds can act as Lewis acids?
•
Lewis acids must have a vacant orbital (into which the electron pair can be donated).
•
Lewis acids sometimes have an incomplete octet (e.g , BF ).
•
Transition-metal ions can be Lewis acids (empty d orbitals)
•
Compounds with multiple bonds can act as Lewis acids.
Example 5:
.
3
Identify the Lewis acid and Lewis base in each of the following reactions:
(a) Fe(ClO ) + 6 H O  Fe(H O)
4 3
-1
(b) CN
2
2
+3
6
-1
+ 3 ClO
4
-1
+ H O  HCN + OH
2
(c) (CH ) N + BF  (CH )NBF
3 3
3
-1
(d) HIO + NH
2
3
 NH + IO
3
-1
3
Acid
H donor
e acceptor
empty orbitals/mult. bonds
incomplete octet/cation
Base
H acceptor
e donator
has lone pairs
often contains N