1. 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

2. 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

3. 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.
No need for water although it can be present, it need not be.
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-

4. 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.

5. acid dissociation constant
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-
Ka is an equilibrium constant called the
acid dissociation constant

6. Water is AMPHOTERIC. It can act as an acid or a base
Acid and Base Strength
:NH3 + H2O NH4++OH-
(a molecular base)
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)
HA + H2O H3O++A-
Water is AMPHOTERIC. It can act as an acid or a base

7. Acid and Base Strength HClO4 ClO4- H2SO4 HSO4- HCl Cl- H3O+ H2OKa
~1010
1x10-2
6.8x10-4
1.75x10-5
9.5x10-8
5.7x10-10
4.7x10-11
1.8x10-16
HClO4 ClO4-
H2SO4 HSO4-
HCl Cl-
H3O+ H2O
HSO4- SO42-
HF F-
CH3COOH CH3COO-
H2S HS-
NH4+ NH3
HCO3- CO32-
H2O OH-
Stronger
Acid
Levelling Effect
Each acid will transfer a proton to a base below it in a mixed solution
Stronger
Base

8. Ionisation of water and pH
pH concept
pH = -log[H+]
pX = -logX
pH scale
[H+] > 10-7M, pH < 7
ACIDIC
[H+] < 10-7M, pH > 7
BASIC
[H+] = 10-7M, pH = 7
NEUTRAL
For any Bronsted conjugate Acid-Base pair
Ka . Kb = Kw

9. The pH Scale

10. 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.

11. Buffers: buffer the system against extreme changes in pH
Every life form is extremely sensitive to slight pH changes. Human blood for example needs to remain within the range
Buffers: buffer the system against extreme changes in pH
CH3COOH H++CH3COO-
Buffer solutions normally consist of two solutes: a weak Bronsted acid and its conjugate base

12. Buffers In general for: HAA- + H+ Henderson-Hasselbach Equation
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?

13. Acid-Base Reactions HCl + NaOH NaCl + H2O
 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.

14. Titration of a strong acid with a strong base
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 = nb mol NaOH
0.02mol/L x 50/1000 L = 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

15. Titration of a strong acid with a strong base
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
Initial conc.
pH = 1.88
Fraction of H+ remaining
Dilution factor
after equivalence pt.: 10.1mL NaOH added
pOH = 3.78
pH = 10.22
Initial conc. of base
Dilution factor

16. Titration Curves
Titration curve of a strong acid with a strong base

17. 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
K = 1/Kb = 1/(Kw / Ka) = 1.8 x 109
Ve = 10mL (as before)
Initial pH: a weak acid equilibrium problem
CH3COOH H++CH3COO-
0.02-x x x
x = 6 x 10-4, pH = 3.22

18. Titration of a weak acid with a strong base
Before eq. pt.: buffer system
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 3/10
Relative final: 7/10 3/10

19. Titration of a weak acid with a strong base
When volume of base added = 1/2Ve
at equivalence pt.: we have a solution of base in water
CH3COONa + H2O  CH3COOH + OH-
F-x x x
Kb = (Kw / Ka) = 5.56 x = x2/(F-x)
x = 3.05 x 10-6, pOH = 5.52, pH=8.48 (BASIC)

20. 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
pH = 10.22

21. 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 acid-strong base titrations.

22. Acid-Base Indicators HIn(aq.) H+ (aq.) +In- (aq.)
Usually dyes that are weak acids and display different colours in protonated/deprotonated forms.
HIn(aq.) H+ (aq.) +In- (aq.)
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

23. Acid-base indicators Indicator pH range pKa Acid Form Base Form
methyl violet yellow blue
thymol blue red yellow
methyl yellow red yellow
methyl orange red yellow
bromocresol green yellow blue
methyl red red yellow
bromothymol blue yellow blue
phenol red yellow red
thymol blue yellow blue
phenolphthalein colourless red
thymolphthalein colourless blue
alizarin yellow R yellow red
indigo carmine blue yellow

24. Solubility Product Ksp Consider for which
Ksp is the solubility-product constant. (BaSO4 is ignored because it is a pure solid).

25. 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:
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.