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-Review

Acids and BasesArrhenius 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 basesBronsted 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.

Strong acids (memorise) dissociate completely in water

HClO4, HClO3, HCl, HBr, HI, HNO3 and H2SO4

Acid and Base Strength

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-

COOHCH

COOCH OHK

3

33a

Ka is an equilibrium constant

called the

acid dissociation constant

Acid and Base Strength:NH3 + H2O NH4

++OH-

(a molecular base)

3

4b NH:

OH NHK

The magnitude of the Ka or Kb, using water as a common proton donor/acceptor, determines the

strength of the acid or base

Water is AMPHOTERIC. It can act as an acid or a base

HA + H2O H3O++A-

HA

A OHK 3

a

In general (for acids)

Acid and Base Strength

HClO4 ClO4-

H2SO4 HSO4-HCl Cl-

H3O+ H2O

HSO4- SO4

2-

HF F-

CH3COOH CH3COO-

H2S HS-

NH4+ NH3

HCO3- CO3

2-

H2O OH-

Stronger

Acid

Stronger

Base

Ka

~1010

1x10-2

6.8x10-4

1.75x10-5

9.5x10-8

5.7x10-10

4.7x10-11

1.8x10-16

Levelling Effect

Each acid will transfer a proton to a base below it

in a mixed solution

Ionisation of water and pH

For any Bronsted conjugate Acid-Base pair

14pOHpH

K10(55.4)K OH H Product Ion

101.8055.4

1010

OH

OH HKK

tionautoionisa HOHaq.OH

w14

a

1677

2ac

2

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

NEUTRALKa . Kb = Kw

The pH ScaleThe pH Scale

The Common Ion EffectThe 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.

BuffersEvery 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

Buffer solutions normally consist of two solutes: a weak Bronsted acid and its conjugate base

CH3COOH H++CH3COO-

-

3

3a

3

-3

a

COOCH

COOHCHKH

COOHCH

COOCH HK

COOHCH

COOCHlogpKpH

3

-3

a

BuffersIn general for: HAA- + H+

HA

AlogpKpH

-

a

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?

Acid-Base Reactions Acid/Base reactions are reactions that involve the neutralisation of an acid through the use of a base.

HCl + NaOH NaCl + H2OIn 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.  

TitrationTitration 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 TitrationsThe 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 NaOHH+ + 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

Acid-Base TitrationsTitration 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

L0.0132mol/350

500.02mol/L

10

310][H

Fraction of H+ remaining

Dilution factorInitial conc.

pH = 1.88

after equivalence pt.: 10.1mL NaOH added

l/L0.000166mo10.150

0.10.1mol/L][OH

Initial conc. of base Dilution factor

pOH = 3.78

pH = 10.22

Titration Curves

-2 0 2 4 6 8 10 12 14 160

2

4

6

8

10

12

Equivalence pt.

pH

Volume NaOH added (mL)

Titration curve of a strong acid with a strong base

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

Initial pH: a weak acid equilibrium problem

x-0.02K

COOHCH

COOCH HK

2

a

3

-3

a

x

x = 6 x 10-4, pH = 3.22

CH3COOH H++CH3COO-

0.02-x x x

Ve = 10mL (as before)Reaction is the reverse of Kb for CH3COO- base

K = 1/Kb = 1/(Kw / Ka) = 1.8 x 109

Titration of a weak acid with a strong base

Before eq. pt.: buffer system HA

AlogpKpH

-

a

Imagine we have added 3.00mLs of base

CH3COOH + NaOH CH3COONa + H2ORelative Initial: 1 3/10Relative final: 7/10 3/10

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.

37.47/10

3/10log74.4pH

Titration of a weak acid with a strong base

Kb = (Kw / Ka) = 5.56 x 10-10 = x2/(F-x)

x = 3.05 x 10-6, pOH = 5.52, pH=8.48 (BASIC)

When volume of base added = 1/2Ve

apK4.745/10

5/10log4.74pH

at equivalence pt.: we have a solution of base in water

CH3COONa + H2O CH3COOH + OH-

F-x x x

L0.0167mol/1050

500.02mol/LF

Titration of a weak acid with a strong baseafter equivalence pt.: pH is determined by excess base added

For 10.1mL base added in total

l/L0.000166mo1.0150

0.10.10mol/L][OH

pOH = 3.78

pH = 10.22

-2 0 2 4 6 8 10 12 14 162

4

6

8

10

12

equivalence pt.

pH

Volume of NaOH added

Acid-Base TitrationsAcid-Base Titrations

Weak Acid-Strong Base TitrationsWeak 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.

Acid-Base IndicatorsUsually 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

HIn

In HK

-

a

Acid-base indicatorsIndicator 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

BaSO4(s) Ba2+(aq) + SO42-(aq)

KKspsp• Consider

• for which

• Ksp is the solubility-product constant. (BaSO4 is ignored because it is a pure solid).

]SO][Ba[ -24

2spK

Factors That Affect SolubilityFactors That Affect Solubility

Common-Ion EffectCommon-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.

CaF2(s) Ca2+(aq) + 2F-(aq)