Acids and Bases-Review
The properties of acids include the following:
•
Taste sour (but don't taste them!!)
•
Their water solutions conduct electrical current (electrolytes)
•
They react with bases to form salts and water
•
Turns Blue Litmus Paper to Red
The properties of bases include the following:
•
Have a slippery feel between the fingers
•
Have a bitter taste (but don't taste them!!)
•
React with acids to form salts and water
•
Turns Red Litmus Blue
•
Their water solutions conduct electrical current (electrolytes)
Acids and Bases
Arrhenius in 1884 discovered that acids give off H+ ions and
allow for a good flow of electricity through a solution.
Arrhenius also discovered that bases give off OH- ions and
OH- ions also allow for a good flow of electricity through the
solution.
Traditionally Professor Arrhenius defined:
Acid released Hydrogen ion (as Hydronium ions,
H3O+) in water solution.
Base produced Hydroxide ion in water solution.
The limitations on these definitions were:
1. The need for water
2. The need for a protic acid
3. The need for Hydroxide bases
Bronsted/Lowry acids and bases
Bronsted and Lowry defined these two terms the following:
Acid-Proton donor Base-Proton acceptor
These definitions are not as restrictive as Arrhenius’ definitions.
1. No need for water although it can be present, it need not be.
2. Bases do not have to be Hydroxide compounds.
However, one restriction still remaining is the need for a protic acid.
Each Bronsted acid is coupled to a
conjugate base to constitute a
CONJUGATE ACID-BASE PAIR
CH3COOH + H2O H3O++CH3COO-
Lewis Acids and Bases
G.N. Lewis defined these in an even less restrictive manner:
Acid- Electron pair acceptor
Base- Electron pair donor
In this set of definitions there is no longer a need for a protic
acid. In other words only electron exchange must occur.
These definition sets are NOT contradictory. A Proton donor is the same as
an electron acceptor. A Proton acceptor is the same as an electron donor. Also
the first set of definitions are less inclusive so that all of the Arrenhius acids
are found under the Bronsted definition but not all Bronsted acids will be
Arrenhius acids. All Arrenhius and Bronsted acids will be under the Lewis
definition but not all Lewis acids will be Bronsted or Arrenhius acids.
Acid and Base Strength
Strong acids (memorise) dissociate completely in water
HClO4, HClO3, HCl, HBr, HI, HNO3 and H2SO4
Strong bases are the metal hydroxides of Group 1 and
heavy Group 2
E.g. LiOH, NaOH, KOH, Ba(OH)2 etc
Weak acids and bases are not completely ionised in solution
CH3COOH + H2O H3O++CH3COO-

H O CH COO 


Ka
3

3
CH3COOH
Ka is an equilibrium constant
called the
acid dissociation
constant
Acid and Base Strength
(a molecular base)

NH OH 


:NH3 + H2O NH4++OH-
Kb

4
: NH3 
The magnitude of the Ka or Kb, using water as a
common proton donor/acceptor, determines the
strength of the acid or base
In general (for acids)

H O A 


HA + H2O H3O++A-
Ka

3
HA 
Water is AMPHOTERIC. It can act as an acid or a base
Acid and Base Strength
Ka
Stronger
Acid
~1010
HClO4
H2SO4
HCl
H3O+
1x10-2
HSO46.8x10-4
HF
CH3COOH 1.75x10-5
9.5x10-8
H2S
5.7x10-10
NH4+
HCO34.7x10-11
H2O
1.8x10-16
ClO4Levelling
HSO4Effect
ClEach acid
H2O
will transfer
SO42a proton to a
Fbase below it
CH3COOin a mixed
HSsolution
NH3
CO32Stronger
OHBase
Ionisation of water and pH
H 2 O  aq.  OH   H 
Kc  Ka
autoionisa tion
pH concept
H OH   10 10   1.80 10

Ion Product


7
H 2O
7
16
55.4
H OH   K (55.4)  10

pH = -log[H+]

a
14
 Kw
pX = -logX
pH scale
[H+] > 10-7M, pH < 7
pH  pOH  14
ACIDIC
[H+] < 10-7M, pH > 7
For any Bronsted conjugate Acid-Base pair
Ka . Kb = Kw
BASIC
[H+] = 10-7M, pH = 7
NEUTRAL
The pH Scale
The Common Ion Effect
• The solubility of a partially soluble salt is decreased when a
common ion is added.
• Consider the equilibrium established when ethanoic acid is added
to water.
• At equilibrium H+ and C2H3O2- are constantly moving into and out
of solution, but the concentrations of ions is constant and equal.
• If a common ion is added, e.g. C2H3O2- from NaC2H3O2 (which is
a strong electrolyte) then [C2H3O2-] increases and the system is no
longer at equilibrium.
• So, [H+] must decrease, according to Le Chatelier’s Principle.
Buffers
Every life form is extremely sensitive to slight pH changes. Human
blood for example needs to remain within the range 7.38-7.42.
Buffers: buffer the system against extreme changes in pH
CH3COOH H++CH3COOBuffer solutions
normally consist
of two solutes: a
weak Bronsted
acid and its
conjugate base
pH  pK a  log

H CH COO 


Ka
-
3
CH 3COOH
H    K  CH3COOH
a
CH COO 
-
3
CH COO 
-
3
CH3COOH
Buffers
In general for:
HAA- + H+
Henderson-Hasselbach Equation
pH  pK a

A
 log
-
HA 
Buffer capacity
Q. If we generate 0.15mol H+ in a reaction vessel of 1L (with no accompanying
volume change) containing 1mol each of CH3COOH and CH3COO-, what will the
solution pH change be?
For the same reaction in water what is the pH change?
Acid-Base Reactions
Acid/Base reactions are reactions that involve the neutralisation of an acid through
the use of a base.
HCl + NaOH NaCl + H2O
In this reaction, the Na+ and the Cl- are called spectator ions because they play no
role in the overall outcome of the reaction. The only thing that reacts is the H+
(from the HCl) and the OH- (from the NaOH). So the reaction that actually takes
place is:
H+ + OH- H2O
If in the end, the OH- was the limiting reagent and there are H+'s still left in the
solution then the solution is acidic, but if the H+ was the limiting reagent and OH-'s
were left in the solution then the solution is basic.
Titration
Titration is the process of mixing acids and bases to analyse one of the solutions.
For example, if you were given an unknown acidic solution and a 1 molar NaOH
solution, titration could be used to determine what the concentration of the other
solution was.
Acid-Base Titrations
The goal of titration is to determine the equivalence point. The equivalence point
is the point in which all the H+ and the OH- ions have been used to produce water.
Titration also usually involves an indicator. An indicator is a liquid that turns a
specific colour at a specific pH. (Different indicators change colours at different
pH's). Indicators are chosen to allow a colour change at the equivalence point.
Titration of a strong acid with a strong base
50.00mL of 0.020M HCl with 0.100M NaOH
H+ + OH- H2O
Kc=1/Kw=1014
at equivalence pt.:
nb mol HCl
0.02mol/L x 50/1000 L
=
nb mol NaOH
=
0.1mol/L x Ve(mL)/1000 L
Ve = 0.001mol HCl  (0.1mol/L x 1/1000 L) = 10 mL
pH determined by dissociation of H20: Kw = [H+][OH-] = 10-14
[H+] = 10-14 = 10-7 mol/L => pH = 7.00
Acid-Base Titrations
Titration of a strong acid with a strong base
Initial pH:
0.02mol/L strong acid.
pH = 1.70
before equivalence pt.:
when 3.00mL of NaOH has been added
 10  3 
 50 


[H  ]  
0.02mol/L


  0.0132mol/ L
 10 
 50  3 
Fraction of
H+ remaining
Initial
conc.
after equivalence pt.:
Dilution factor
pH = 1.88
10.1mL NaOH added
 0.1 
[OH ]  0.1mol/L 
  0.000166mo l/L
 50  10.1 

Initial conc.
of base
Dilution factor
pOH = 3.78
pH = 10.22
Titration Curves
Titration curve of a strong acid with a strong base
12
10
Equivalence pt.
pH
8
6
4
2
0
-2
0
2
4
6
8
10
12
14
Volume NaOH added (mL)
16
Titration of a weak acid with a strong base
Take the example of a titration of 50.0mL 0.020M CH3COOH
(Ka = 1.8 x 10-5) with 0.10M NaOH
CH3COOH + NaOH  CH3COONa + H2O
Reaction is the reverse of Kb for
CH3COO- base
Ve = 10mL (as before)
K = 1/Kb = 1/(Kw / Ka) = 1.8 x 109
Initial pH:
a weak acid equilibrium problem
H CH COO 


CH3COOH
0.02-x
H++CH3COOx
Ka
x
x = 6 x 10-4, pH = 3.22
-
3
CH3COOH
x2
Ka 
0.02 - x
Titration of a weak acid with a strong base
Before eq. pt.:
buffer system
pH  pK a

A
 log
-
HA 
One of the simplest ways to treat these problems is to evaluate the quotient in the log
using relative concentration before and after the reaction.
Imagine we have added 3.00mLs of base
CH3COOH + NaOH  CH3COONa + H2O
Relative Initial: 1
Relative final: 7/10
3/10
3/10
3/10
pH  4.74  log
 4.37
7/10
Titration of a weak acid with a strong base
When volume of base added = 1/2Ve
5/10
pH  4.74  log
 4.74  pK a
5/10
at equivalence pt.: we have a solution of base in water
CH3COONa + H2O  CH3COOH + OHF-x
x
x
 50 
F  0.02mol/L 
  0.0167mol/ L
 50  10 
Kb = (Kw / Ka) = 5.56 x 10-10 = x2/(F-x)
x = 3.05 x 10-6, pOH = 5.52, pH=8.48 (BASIC)
Titration of a weak acid with a strong base
after equivalence pt.: pH is determined by excess base added
For 10.1mL base added in total
pOH = 3.78
 0.1 
[OH  ]  0.10mol/L 
  0.000166mo l/L
 50  10.1 
pH = 10.22
12
10
equivalence pt.
pH
8
6
4
2
-2
0
2
4
6
8
10
12
Volume of NaOH added
14
16
Acid-Base Titrations
Weak Acid-Strong Base Titrations
• The weaker the acid, the
smaller the equivalence point
inflection.
• For very weak acids, it is
impossible to detect the
equivalence point.
• Choose an indicator with a Ka
range suited to the weak acid.
•Titration of weak bases with strong acids have similar features to weak acidstrong base titrations.
Acid-Base Indicators
Usually dyes that are weak acids and display different
colours in protonated/deprotonated forms.
HIn(aq.) H+ (aq.) +In- (aq.)

H In 


Ka
-
HIn 
In general we seek an indicator whose transition range (±1pH
unit from the indicator pKa) overlaps the steepest part of the
titration curve as closely as possible
Acid-base indicators
Indicator
pH range pKa
Acid Form
Base Form
methyl violet
0.0- 1.6
0.8
yellow
blue
thymol blue
1.2- 2.8
1.6
red
yellow
methyl yellow
2.9- 4.0
3.3
red
yellow
methyl orange
3.1- 4.4
4.2
red
yellow
bromocresol green
3.8- 5.4
4.7
yellow
blue
methyl red
4.2- 6.2
5.0
red
yellow
bromothymol blue
6.0- 7.6
7.1
yellow
blue
phenol red
6.4- 8.0
7.4
yellow
red
thymol blue
8.0- 9.6
8.9
yellow
blue
phenolphthalein
8.0- 9.8
9.7
colourless
red
thymolphthalein
9.3-10.5
9.9
colourless
blue
alizarin yellow R
10.1-12.0 11.0
yellow
red
indigo carmine
11.4-13.0 12.2
blue
yellow
Solubility Product
Ksp
• Consider
• for which
BaSO4(s)
Ba2+(aq) + SO42-(aq)
2
2
K sp  [Ba ][SO4 ]
• Ksp is the solubility-product constant. (BaSO4 is ignored because it
is a pure solid).
Factors That Affect Solubility
Common-Ion Effect
• Solubility is decreased when a
common ion is added.
• This is an application of
Le Châtelier’s principle:
CaF2(s)
Ca2+(aq) + 2F-(aq)
• as F- (from NaF, say) is added, the
equilibrium shifts away from the
increase.
• Therefore, CaF2(s) is formed
and precipitation occurs.
• As NaF is added to the system, the solubility of CaF2 decreases.