Acid base ph

1. Acid, base N buffers

2. • Arrhenious concept
• An acid is a hydrogen containing compd which
gives hydrogen ion when dissolved in water.
HCl + H2O-------H+(aq) + Cl- (aq)
H2SO4 + H2O-----H+ + SO4—
CH3COOH
• Base---gives OH- ions
NaOH
NH4OH

3. limitations
• Acid base behaviour in aqueous conditions
only..not acid in gaseous state or dissolved in
other solvents
Acids and bases have to contain H and OH- ions
respectively…
NaNH2 is a strong base, it gives Na+ and NH2-
but no OH-

4. Bronsted-lowry concept
• Not limited to water as a solvent
• acid = species donating protons
• Base = species accepting protons
NH3 + H2O ------- NH4+ + OH-
• Water donates a proton and NH3 accepts
HCl + H2O----- H3O+ + Cl-
HCl donates and water accepts

5. limitation
• It is applicable to aqueous and non-aqueous
protic solvents equally well.
• However, acids have to contain hydrogen
• CO2, SO2 acts as acid

6. Conjugate acids and bases
• When an acid loses a proton in solution to act
as an acid, the remaining part of its molecule
which is an anion will have a tendency to gain
the proton back and thus acts as a base
• CH3COOH + H2O--------CH3COO- + H3O+
• A+B----CB+CA

7. Lewis concept
• Acid= accepts a pair of electron
• Base = electron donor
:NH3 + BF3----H3N:->BF3
Base acid ammonia boron trifluoride

8. neutralization
• The reaction of an acid with a base in which
characteristic properties of both are destroyed
or neutralized, is called neutralization.
• When an acid reacts with a base it produces a
compd called salt and water
• HCl+ NaOH-----NaCl + H2O

9. salt
• Compound formed by the partial or complete
neutralization of an acid by a base.
• H2SO4 + NaOH----NaHSO4 + H2O
• H2SO4 + Ca(OH)2----CaSO4 + 2H2O

10. Classification of salt
• Normal salt- complete neutralization
• Acidic salt- partial neutralization of acid by base
H2SO4 + NaOH----NaHSO4 + H2O
• Basic salt- partial neutralization of base by acid
Zn(OH)2 + HCl-----Zn(OH)Cl + H2O
• Mixed salt- consists of different cations or anions
NaKSO4

11. pH & buffer
• pH=-logH+
• General form of the buffering reaction
Buffer +H+ =============== H Buffer

12. Henderson hasselbalch eqn
• pH= pKa + log base/acid
Or
• pH= pKa + log salt/acid
• When base = acid, pH=pKa
• Therefore, when the concentration of base and
acid are the same, then pH is equal to pKa.
– The pH of a buffer on addition of known quantity of
acid and alkali can therefore be predicted by the eqn
– Moreover, the concentration of salt or acid can be
found out by measuring th pH

13. buffer
• Aqueous system that tend to resist any
changes in pH when small amount of acid or
base is added.
• Are formed by weak acid and weak base
• Never a strong acid or base
• Types
– Acidic buffer
– Basic buffer

14. • Acidic buffer
• Mixture of wk acid and conjugate base ( salt of
same acid with a strong base)
• CH3COOH/CH3COONa
• Basic buffer
• Mixture of wk base and conjugate base (salt of
same base with a strong acid)
• NH4OH/NH4Cl

15. Mechanism of buffer action
• Acetate buffer
• CH3COOH/CH3COONa-
– When an acid is added
• HCl + CH3COONa----- CH3COOH + NaCl
• St.acid wk acid
– When a base is added
• NaOH + CH3COOH------CH3COONa + H2O
• St.base wk base

16. Biological buffer
• Bicarbonate buffer
• Phosphate buffer
• Protein buffer
• Hb buffer

17. • H+ concn and pH
• Ka dissociation constant and pKa
• The pH at which an acid is half ionized is called
pKa of an acid which is constant at a particular
temperature and pressure.
• Strong acids will have a low pKa and weak
acids have a higher pKa

18. Henderson hasselbalch eqn
• pH= pKa + log base/acid
Or
• pH= pKa + log salt/acid
• When base = acid, pH=pKa
• Therefore, when the concentration of base and
acid are the same, then pH is equal to pKa.
– The pH of a buffer on addition of known quantity of
acid and alkali can therefore be predicted by the eqn
– Moreover, the concentration of salt or acid can be
found out by measuring th pH

19. • Normal pH= 7.4
• Range 7.38-7.42
• Acidosis
– Below 7.38
– Below 7.25- life threatening—CNS depression and
coma
– Below 7.0-death
• Alkalosis
– Above 7.42
– Above 7.55 dangerous-neuromuscular
hyperexcitability and tetany
– Above 7.6-death

20. Mechanism of regulation of pH
1. buffers of body fluids
2. Respiratory system
3. Renal excretion
These mechanisms are inter-related
Buffers are the first line of defense against acid load. The
buffers are effective as long as the acid load is not
excessive, and the alkali reserve is not exhausted.
Once the base is utilised in this reaction, it is to be
replenished to meet further challenge.

21. Bicarbonate buffer
• NaHCO3 -sodium bicarbonate-kidney
• H2CO3- carbonic acid-respiratory component
65% of buffering capacity in plasma
40% in whole body
pH= 7.4 (blood)
pKa of H2CO3 = 6.1
Ratio of NaHCO3/H2CO3=20/1
Also called as alkali reserve
Strong and non volatile acids are converted to weak
and volatile form.e.g, HCl, H2SO4, lactic acid, ketoacids

22. • pH= pKa + log salt/acid
• pH= 6.1 + log 24/ 1.2
• pH = 6.1 + log 20 = 6.1 + 1.3
• Mechanism
• Acidic condition
• HCl + NaHCO3------ NaCl + H2CO3-----H2O +
CO2----exhaled
• Basic condition
• NaOH + H2CO3-----NaHCO3 + H2O----excreted
through kidney

23. Phosphate buffer
• Also called as intracellular buffer
• Na2HPO4
• NaH2PO4
• pH=7.4
• pKa of NaH2PO4=6.8
• Ratio = 4/1
• Acidic condition
• HCl + Na2HPO4----NaCl + NaH2PO4-----kidney acidic urine
• Basic
• NaOH + NaH2PO4----H2O + Na2HPO4---kidney alkaline
urine

24. Protein buffer
• Buffering capacity depends on pKa value of
ionizable side chain
• Accepts H+ from medium and acts as base in
acidic condition and vice-versa

25. Hb buffer
• Buffering capacity of Hb is due to imidazole ring
of histidine
• One Hb has 38 histidine
• Imidazole ring contains 2 groups
1. Iron containing grp
2. Immidazole nitrogen- donates or accept proton
depending on pH of the medium
Oxygenation of Hb acts as acid
Deoxygenation of Hb acts as base

27. N
H
N--- Fe++
N
N--- Fe++
-- H+ +HCO3------ H2CO3-----H2O+ CO2------exhaled
H+ + HCO3- -----H2CO3 --------H2O + CO2
Oxygenation state
In lungs
Deoxygenation state
In tissues