1. This document discusses acid-base theories including Arrhenius, Bronsted-Lowry, and acid-base equilibria. 2. It explains the ion product constant of water (Kw) and how pH and pOH scales are used to measure hydrogen and hydroxide ion concentrations. 3. Weak acids and bases only partially dissociate in water and their equilibria are expressed using acid (Ka) and base (Kb) dissociation constants.

1. ACID BASE EQUILIBRIUM
GROUP MEMBERS :
NOOR AZURAH ABDUL RAZAK
(D20101037502)
SITI NORHAFIZA HAFINAS MOHD ZANUDIN
(D20101037503)

2. Ion Product constant
Kw=[H†][OH⁻]
=(1.0x10⁻⁷)(1.0x10⁻⁷) Theory
= 1.0x10⁻ˡ⁴ Arheniuss
Bronsted-Lowry
pH = -log[H†]
pOH = -log[OH⁻]
ACID BASE
hpH + pOH = 14.00
EQUIL RIUM
IB
Weak Acid Weak Base
Strong Acid Strong Base
HCN,H2S NH3
HCl, HBr NaOH
 H+   A−   BH +   OH − 
Ka =     Kb =   
Ionize completely(100%) in [ HA ] [ B]
water
Dissociate partially in water

3. ACID-BASE THEORY

4. Arrhenius Theory
• The concept began in 1887
• Acid - a substance that, when dissolved in
water, increases the concentration of H+ ions.
• Base - a substance that, when dissolved in
water, increases the concentration of OH-
ions.
• Disadvantage- Apply only to aqueous solution

5. Bronsted-Lowry Theory
 In 1923, Johannes Bronsted and Thomas
Lowry proposed a more general definition
of acids bases
 Acid - a substance (molecule or ion) that
donates a proton to another substance
Base - a substance that accepts a proton
Not necessarily to be in aqueous solution

6. ION PRODUCT OF WATER
 Ion-product of water refers to autoionization of water
H₂O(l) ↔ H†(aq) + OH‾ (aq)
 The equilibrium constant for the autonization of water
is :
Kw =[H†][OH‾ ]
Kw is called ion-product constant, which is the product of
the molar conentration of H† and OH‾ at a particular
temperature

7. In pure water at 25⁰c, [H†]=[OH⁻]=1.0x10⁻⁷
Kw =[H†][OH⁻]
= (1.0x10⁻⁷)(1.0x10⁻⁷)
= 1.0x10⁻ˡ⁴

8. pH SCALE
 Use to measure the concentration of H† in a solution
 Measuring the acidity of the solution
 pH : negative base-10 log of the concentration of hydrogen oin,
[H†]
pH = -log[H†]
 pOH : negative base-10 log of the concentration of hydroxide oin,
[OH⁻]
pOH = -log[OH⁻]
pH + pOH = 14.00

10. Example :
Given that the pH of hydrogen sulfide,H₂S
is 4.82. Find the concentration of
hydrogen ion, [H†]
Solution :
pH = -log[H†]
4.82 = -log[H†]
[H†] = antilog (-4.82)
= 1.51 x 10⁻⁵

11. Strong
acid
Weak
Weak
acid
Acid &Base base
Strong
base

12. Strong acid
+ −
HCl( aq ) + H 2O( l)     H 3O
→
100%ionization
( aq ) + Cl ( aq )

13. Strong bases
• Dissociate completely in an aqueous solution
to produce high concentration of OH-
• 100% ionization or 100% dissociation
• Example NaOH and KOH
NaOH( aq )     Na + ( aq ) + OH − ( aq )
100%ionization
→

14. Weak acids
CH 3 COOH ( aq ) + H 2 O ( l ) ⇔ CH 3 COO − ( aq ) + H 3 O + ( aq )

15. Weak bases
+ −
NH 3( aq ) + H 2 O ⇔ NH 4 ( aq ) + OH ( aq )

16. ACIDS AND EQUILIBRIUM
• Equilibrium constant can be used for the ionization reaction to
express the extent to which a weak acids ionizes.
HA( aq ) + H 2O( l ) ⇔ H 3O + ( aq ) + A− ( aq )
HA( aq ) ⇔ H + ( aq ) + A− ( aq )
Ka =
[ H O ][ A ]
3
+ −
or Ka =
[ H ][ A ]
+ −
[ HA] [ HA]
• A smaller Ka value meaning the degree of the dissociation of
the acid is low

17. BASES AND EQUILIBRIUM
• the base is a weak base (ionizes <100%), the
equilibrium is dependent on the base
equilibrium constant, Kb.
• A base equilibrium expression can only be
written for weak bases.
B( aq ) + H 2 O( l ) ⇔ BH + ( aq ) + OH − ( aq )
Kb =
[ BH ][OH ]
+ −
[ B]