Chapter 10:Chapter 10:Acid Base Equilibria and Acid Base Equilibria and Buffer SolutionsBuffer Solutions

Chemistry For STPM

10.1 Theories of Acid and Bases

Arhenius Theory

Bronsted-Lowry

Lewis Theory

Acid + Water H+

Base + Water OH-Acid: Proton donor

Base: Proton acceptor

HCl + H2O Cl- + H3O+

RNH2 + H2O RNH3 + OH-

Water can acts as acid and base:Amphoteric Solvent

Acid: Electron pair acceptor[short of e, cation]

Base: Electron pair donor [extra e, anion]

NH3 + BF3 NH3BF3

HCl + NH3 Cl- + NH4+ Acid Base Conjugate base, conjugate acid

10.2 Electrolyte

Chemical Compound

Electrolyte Non-electrolyte

Strong: Fully dissociate into ions in aqueous solution, high degree of dissociation, good electrical conductor

Weak: Partially dissociate into ions in aqueous solution, low degree of dissociation, poor electrical conductor

10.3 The Strength of Acid and Base

Strong Acid and Weak Acid

Able to donate proton easily; completely dissociate

hardly donate proton ; partially dissociate

HCl + H2O Cl- + H3O+ Acid Base Conj. Base conj. acid

•The stronger the acid, the weaker its conjugate

• HCl is strong acid, easier to donate H+, equilibrium lies to the right.

• Conjugate base Cl- hardly receive H+.

Ostwald Dilution Law

CH3COOH CH3COO- + H+ t= 0 s (mol) c 0 0 t =equilibrium c (1-α) cα cα

Ka = Cα2 /(1-α) Dissociation Constant of Acid(KDissociation Constant of Acid(Kaa))

For weak acid, dissociation constant For weak acid, dissociation constant α is small, is small,

(1-(1-α) = 1, thus Ka = Cα2 .[H+] = Cα = c √ Ka/c = √ Ka x c

NH3 + H2O NH4+ + OH-

t= 0 s (mol) c 0 0t =equilibrium c (1-α) cα cα

Kb = Cα2 /(1-α) Dissociation Constant of Base(KDissociation Constant of Base(Kbb))

For weak base, dissociation constant For weak base, dissociation constant α I I small, small,

(1-(1-α) = 1, thus Kb = Cα2 .[OH-] = Cα = c √ K b/c = √ Kb xc

pKa = -log10 Ka

•The higher Ka, the stronger the acid.

•The higher pKa, the weaker the acid.

•The higher Kb, the stronger the base

• The higher pKb, the weaker the base

Various Value of Ka

The degree of second dissociation is less than first dissociation.

Second proton is hardly to be released due to strong attraction of proton by negatively charged ion.

Acid Equilibrium pKa

H3PO4 H3PO4 H2PO4- + H+ pk1 = 2.1

H2PO4- H2PO4

- HPO42- + H+ pk2 = 7.2

HPO42- HPO4

2- PO43- + H+ Pk3 = 12.4

10.4 Factor Affecting Acid and Base 10.4 Factor Affecting Acid and Base StrengthStrength

1. Which is behave as acid? Why? Na(OH) @ SO2(OH)2

2. Arrange the strength acid in decreasing order. Give reasons for the arrangement. CH3COOH, ClCH2COOH, Cl2CHCOOH

1. When an electronegative element is attached to OH, the O-H bond is weakened and proton tend to be released.Since Na is electropositive element, H unlikely to be lost.Thus, Na(OH) is a base whereas SO2(OH)2 is acid.

2. Carboxylic Acid, R(Akyl group) is attached electron withdrawing group, electron from O – H will be pulled away by Cl, thus H is easier to be given off.

Cl2CHCOOH , ClCH2COOH, CH3COOH

Consider

HH22O HO H++ + OH + OH--

Kc = [H+][OH-]/ [H2O]Given: Kw = [H+][OH-]

Then : Kw = Kc [H2O]Kw = Ka x Kb = 10-14 ………Explain!

pKw = pKa + pKb = 14 In pure water: [H+]=[OH-] = 10-7 mol dm-3 The Kw affected by the temperature

Dissociation of Water

Practice 1

1. When 0.1 mol of etanoic acid is dissolved in water and the volume made up to 1 dm3 , the concentration of H+ is 1.34 x 10-3 mol/dm3 .

2. Which is the stronger acid, HClO or HBrO? Explain your answer.

10.6 pH Scale pH = -log10 [H+] Type of solution:Neutral: [H+] = 10-7

Acidic : [H+] >10-7

Basic : [H+] <10-7

pKw = pKa + pKb = 14 From [10.5]Thus : pH = 14 + lg [OH] =14 – pOHNote: For Strong Acid, [H+] can be obtained directly

from the molarities. For Weak Acid, to find [H+], we need to know Ka.

1. Calculate the pH of 0.02 mol/dm3 sulphuric acid.2. Calculate the pH of a solution prepared by mixing 25.0

cm3 of 0.1 mol/dm3 hydrochloric acid with 15.0 cm3 of 0.1 mol/dm3 sodium hydroxide solution.

3. Calculate the pH for 0.01 mol/dm3 of Ba(OH)2.4. Calculate the pH of a solution prepared by mixing 25.0

cm3 of 0.1 mol/dm3 hydrochloric acid with 25.0 cm3 of 0.2 mol/dm3 sodium hydroxide solution.

5. The pH of 0.15 mol/dm3 butyl amine solution is 12.0 at 198 K. Calculate the dissociation constant of the base and the degree of dissociation.

PRACTICE 2

10.7 Acid- Base Indicators

1. The indicators can change color because their ions have colors that are different from the un-dissociated molecule.

2. Litmus is a weak acid: Hln + H2O H3O

+ + ln-

Red Blue Why litmus paper show red in acidic solution?

In acidic solution, the [H3O+] is high. Thus, the

equilibrium shift to the left.

Indicator Dissociation Constant

KHln = [H3O+] [ln-]

[Hln] Rearrange

[H3O+] = KHln [Hln]

[ln-]

Or pH = p KHln - lg [Hln] [ln-] When [Hln] = [ln-]; pH = p KHln

Indicators pKHln pH range Acid color Alkali color

Methyl Orange 3.7 3.2-4.2 Red Yellow

Methyl Red 5.1 4.2-6.3 Red Yellow

Bromothymol Blue

7.1 6.0-7.6 Yellow Blue

Phenolphthalein

9.3 8.2-10.0 Colorless Red

The colour changes for indicators

Choosing Indicators For an Acid Base titrationChoosing Indicators For an Acid Base titration

1. An acid-base titration is a method used to determine the concentration of either an acid or base.

2. The end-point is the point at which the indicator change colour

3. Type of acid base titration: a) strong acid and strong base b) weak acid and strong base c) strong acid and weak base d) weak acid and weak base

Strong acid and strong base

2

87

12

0 10 20 25 30 cm3 NaOH

pH

phenolphthelein

Bromothymo blue

Methyl orange

HCl + NaOH NaCl + H2O

The pH changes sharply from 3.5 to 10.5

Weak acid and strong base

2

98

12

0 10 20 25 30 cm3 NaOH

pH CH3COOH + NaOH NaCl + H2O

Phenolphthalein

Ka = 1.8 x 10-5 mol/dm3 c = 0.1 mol/dm3

The pH changes sharply from 6.5 to 10.5

Strong acid and Weak base

2

65

12

0 10 20 25 30 cm3 NaOH

pH HCl + NH3 NH4Cl

Methyl orange

Weak acid and Weak base

7

0 10 20 25 30 cm3 NaOH

pH CH3COOH + NH3 CH3COONH4

There is no sharp increase of pH at the endpoint

Titration of Weak Polyprotic Acid

12

10

8

6

4

2

pH H3PO4 + NaOH

0 25 50 cm3 NaOH

Formation of NaH2PO4

Formation of Na2HPO4 Phenolphthalein

Methyl orange

1. A weak polyprotic acid shows a titration curve with separate equivalence point for each acidic hydrogen displaced.

H3PO4 + NaOH NaH2PO4 + H2O …First equivalent point [3.6]

NaH2PO4 + NaOH Na2HPO4 + H2O . Second equivalent point [9.0]

Na2HPO4 + NaOH Na3PO4 + H2O …Third equivalent point [13.0]

Titration of Weak Polyprotic Acid

Sketch the graph to show the changes in pH during the titration of 25 cm3 of 0.1 mol/dm3 sodium hydroxide solution with 0.1 mol/dm3 sulphuric acid.

PRACTICE 3

Ask yourself: Next what u need to know?

or

Correct

?

?

?

Salt Hydrolysis

For Strong Acid and Base: Na+ + Cl- + H2O No reaction

For Strong acid and weak base: NH4Cl NH+ + Cl-

Cl- + H2O No reaction NH4+ + H2O NH3 + H3O

+

For Weak acid + Strong base: CH3COONa CH3COO- + Na+

Na+ + H2O No reaction

CH3COO- + H2O CH3COOH + OH-

Type of Salt Example Nature of salt

Strong acid + Strong base Neutral

Strong acid + Weak base Acidic

Weak acid + Strong base Alkaline

Weak acid + Weak base Depends Ka

10.8 Buffer solution

Acidic Buffer solution Alkaline Buffer solution

Mixture of weak acid and its salt whose pH value changes very slightly when a small amount of an acid or alkali is added

Mixture of weak base and its salt whose pH value changes very slightly when a small amount of an acid or alkali is added

The action of buffer solution

CH3COOH CH3COO- + H+

 CH3COONa CH3COO- + Na+

NH3 + H2O NH4+ + OH-

NH4Cl NH4+ + Cl-

Calculating the pH of Acidic and Alkaline Buffer

Let’s: HA H+ + A – @ NH4 + H+ + NH3

Ka = [H+][A-]/[HA]

How To Get

pH = pKa + lg [Salt]/[Acid]

Ka = [H+][NH3]/[NH4]

How To Get

pOH = pKb + lg [Salt]/[Base]