Acid, Base, pH, Buffer, Henderson-Hasselbach equation following curriculum of Biochemistry for MBBS course

2. Acid Base pH Buffer
Dr. Farhana Atia
Associate Professor
Department of Biochemistry
Nilphamari Medical College, Nilphamari

3. Acid
• A substance that gives off proton [H⁺]
• Proton donor
• H2CO3 ↔ H⁺ + HCO₃⁻
Properties-
Sour in taste
Convert blue litmus into red
Produce salt and water with alkali
Gives H⁺ in aqueous solution

4. Base
• A substance that accepts proton
• Proton acceptor
• HCO₃⁻ : conjugate base of H2CO3
• Acid = Proton + Base
HA H⁺ + A⁻
HCl  H⁺ + Cl⁻
H₂PO₄⁻  H⁺ + HPO₄⁼
NH₄⁺  H⁺ + NH₃
CH₃COOH  H⁺ + CH₃COO⁻

5. • Usually strong acid has weak base: HCl / Cl⁻
Weak acid has strong base: H₂CO₃ / HCO₃⁻
Strong acid Weak acid
Dissociates quickly & completely Slowly & partially
Weak conjugate base Strong
Lower pK value Higher pK
Can not act as buffer Can act as buffer
Inorganic except H₂CO₃ Organic acid + H₂CO₃
In body- HCl, H₂SO₄ H₂CO₃, H₂PO₄

6. pH
• Introduced by Sorensen in 1909
• pH is defined as the negative log of the hydrogen ion
concentration.
The equation is:
pH = - log [H+]
• Pure water has an equal concentration of H+ & OH⁻ ions
[H+] = [HO⁻]= 10⁻⁷ mole/L
So, pH= 7 which is neutral.
• pH of blood plasma 7.35-7.45
• pH of gastric juice 1.5-3
• pH of urine 4.5-6.8

7. pH scale
• The pH scale measures how
acidic or basic a substance
is.
• Range : 0 to 14.
pH 7 = neutral.
pH < 7 = acidic.
pH > 7 = basic.
pH = 0, [H⁺] is maximum
• The pH scale is logarithmic
• pH 4
• 10 times more acidic than
pH 5
• 100 times (10 x 10) more
acidic than pH 6.
• pH 10
• 10 times more alkaline
than pH 9
• 100 times (10 x 10) more
alkaline than pH 8

8. [H⁺] of blood = 40 nmol/L

9. • pH measurement
Indicator dye
Indicator paper
pH meter
Indicator
Low pH
color
High pH
color
Thymol blue Red Yellow
Methyl red Red Yellow
Bromothymol
blue
Yellow Blue
Phenolphthalein Colorless Fuchsia
pH meter

10. pK
• Degree of dissociation of a substance is called dissociation
constant (K)
• pK is the negative logarithm of dissociation constant (K).
pK = - log K

11. pK
• Acid dissociates as,
HA  H⁺ + A⁻
• By the law of mass action: The product of the concentration of
products in a chemical reaction divided by the product of the
concentration of the reactants at equilibrium is a constant
K=
𝐻⁺ [𝐴⁻ ]
[𝐻𝐴]
• Relative strength of weak acid & base is expressed in terms of
their ‘K’
• For strong acid, K= high
• For weak acid, K= low

12. Henderson-Hasselbalch equation
• Important for understanding buffer action and acid-base
balance
• For the dissociation of a weak acid HA into H+ and A-, the
Henderson-Hasselbalch equation can be derived as follows:
• HA dissociates as follows
HA ↔ H⁺ + A⁻
• The equilibrium constant for this dissociation is-
K=
𝐻⁺ [𝐴⁻ ]
[𝐻𝐴]

14. Henderson-Hasselbalch equation
• Now invert -log [HA]/[A-], which involves changing its
sign, to obtain the Henderson-Hasselbalch equation:
• It shows the pK of a weak acid is equal to the pH of the
solution when [HA] = [A-] (i.e. the acid is exactly half
neutralized)

15. Henderson-Hasselbalch equation
• The equation can be used
to estimate pH of a buffer
solution
• For bicarbonate buffer,
pH = pK + log
HCO₃⁻
H₂CO₃

16. Regulation of [H⁺]
• Acids and bases continually enter and exit body.
• For normal function of body and normal enzyme activity a
normal [H⁺] (normal pH) is essential.

18. Buffer
• Buffers are solutions that resist changes in pH of a solution
when acid or alkali is added
Mixture of weak acid & conjugate (H₂CO₃-NaHCO₃)
Mixture of weak base & conjugate (NH₄OH-NH₄Cl)
• Buffer system has 2 components
• Buffers act within second
• Maximum buffering capacity: when pH=pK±1
• Buffer can not remove H⁺ from body. It only remove/ add H⁺ in
solution
• H⁺ is removed from body by renal system

19. Reduce free H⁺/OH temporarily
Thus resist the change in pH
Weak acid

20. Principal buffers in body fluid
Blood
• Bicarbonate
• Phosphate
• Plasma protein
• Hemoglobin
ICF
• Protein
• Phosphate
ECF
• Bicarbonate
• Protein
• Phosphate
Urine
• Ammonia
• Phosphate

21. Bicarbonate buffer
• HCO₃⁻/H₂CO₃
• Concentration: 24/1.2
• so ratio of two
components: 20/1)
• Principal buffer in ECF
• pK=6.1 , so weak buffer
• Can not buffer H₂CO₃,
cannot act in respiratory
acid base disorder
Mechanism of action:
• When strong acid is added-
Buffered by base part (HCO₃⁻):
HCl + NaHCO₃ NaCl + H₂CO₃
• When base/ alkali is added
Acid part acts (H⁺):
NaOH + H₂CO₃  NaHCO₃ + H₂O

22. Bicarbonate buffer
Bicarbonate buffer is the most effective buffer
• Available in ECF & ICF
• Salt: Acid=20:1 (acc. to HHE)
• Salt part more
• Our body is net acid producer
• HCO₃ excess: alkali reserve (20x)
• Buffer components are regulated by body easily
1. NaOH + H₂CO₃ NaHCO₃ + H₂O (excrete through kidney)
2. HCl + NaHCO₃ NaCl+ H₂CO₃
H₂CO3  H₂O + CO₂ (Exhaled)
So, it is an ‘open end’ buffer system.

23. Protein buffer
• Pr⁻ / HPr (basic protein/acid protein)
• Major buffer of I.C.F (intra cellular fluid)
• pK= 7.3 (6.4-7.8)
• Protein  acidic group - COOH
 basic group - NH₂
M/A
• Strong acid: H⁺ + Pr⁻HPr (weak acid)
• Base/alkali : OH⁻+HPr H₂O + Pr⁻ (weak base)

24. Hemoglobin buffer
• Hb ⁻/HHb or
HbO₂⁻/HHbO₂
• Major buffer in RBC
• pK=7.3 (6.6-7.8)
• M/A- as protein
Ammonia buffer
• NH₃/NH₄⁺
• Principal buffer of urine
• Facilitate renal acid
excretion

25. Phosphate buffer
• HPO₄⁼ / H2PO₄⁻
• pK= 6.8, but concentration in ECF is less
M/A
• Na₂HPO4 + HCl  NaCl + NaH₂PO4
• NaH₂PO4 + NaOH  Na₂HPO4 + H₂O
• Important Buffer of renal tubular fluid and I C F

26. Thank You