This document discusses buffer solutions, including their definition, examples, and mechanisms of action. It defines a buffer as a solution that resists changes to its pH when small amounts of acid or base are added. Common buffer systems include carbonic acid/sodium bicarbonate and monosodium phosphate/disodium phosphate. The Henderson-Hasselbach equation relates the pH of a buffer solution to the ratio of its conjugate acid and base forms and their pKa values. Buffers in the body include bicarbonate, phosphate, and protein buffers that help maintain blood pH within a narrow range.

1. Buffer Solutions
SOURCE: MUSHTAQ
MARYAM FIDA (O-1827)

2. Definition
 A Buffer solution is one that tends to maintain its pH when small amounts
of an acid or base are added to it
 Consist of a weakly dissociating acid (proton donor) and its conjugate base
or the salt i.e. the proton acceptor

3. Examples of Buffers
 H2CO3/NaHCO3 (Carbonic Acid/Sodium Bicarbonate)
 NaH2PO4 / Na2HPO4 (Monosodium Phosphate or Sodium Dihydrogen
Phosphate / Disodium Phosphate)
 NaH2PO4 acts as acid because it has one more H+ as compared to the
latter.

4. Suppose HCl is added to buffer system containing NaHCO3 and H2CO3
The following reaction will take place:
 NaHCO3 + HCl  H2CO3 + NaCl
 H2CO3 is a weak acid and a neutral salt is produced. Slight change in pH
 Therefore change in pH resisted

5.  If NaOH is added
 NaOH + H2CO3  NaHCO3 + H2O
 NaHCO3 is a much weaker base than NaOH, only slight rise in pH
 Therefore change in pH resisted

6.  The above two systems are the first line of defence against any strong acid
or base that gains entry into the body
 The other buffers include: Hemoglobin and plasma proteins

7. Indicators and pH-metry
 Indicators are used to find acidic or alkaline reactions of solutions
 Indicator: weakly ionizing acid or base, having distinct colour in the unionized state and a
different colour in the ionized state
 H.Indicator (blue)  H+ + Indicator – (red)
 Add acid. The above will remain undissociated. So colour will remain blue
 If we add base, the indicators will dissociate and give red colour

8. Determination of pH of a solution
Method 1: Indicators (liquids or pH papers)
Method 2: Electrometric Methods (pH meter)
 pH Meter: Has two electrodes. A glass electrode containing known concentration of 0.1M HCl.
 Calomel electrode which is the reference electrode containing mercury-mercurous chloride
paste surrounded by KCl solution
 The potential difference between the two on dipping in unknown solution is amplified and
measured by a potentiometer

9. Henderson-Hasselbach Equation
 HA = Weak Acid
 BA = Salt of HA
 Only HA can provide H+ ions
 This dissociation of a weak acid can be represented by:
 HA  (H+) + (A-)
 (H+)(A-) is a constant value and is called the dissociation or ionization
(HA) constant of an acid represented by Ka (a for acid)

10.  (H+)(A-) = Ka OR H+ = Ka (HA)
(HA) (A-)
 -log (H+) = -log Ka – log (HA) (We know –log(H+) = pH; similarly –log Ka=pKa)
(A-)
 pH = pKa – log (HA)
(A-)
 pH=pKa + log (A-)
(HA)
 pH=pKa + log Base (This is the Henderson – Hasselbach Equation)
Acid

11. Applications of Henderson
 To Determine the pH of the solution- The pH of a solution is a measure of the acidity and basicity of
the solution. The dependance of many reactions, as well as biological functions depends upon the pH
of the solution in which the chemistry occurs. We can calculate the pH of a solution of a chemical by
knowing its pKa and ratio of ionized and unionized chemical
 To determine the pKa of a chemical in solution- The pKa of a solution determines the acidity of the
compound. The pKa also determines the amount of ionization of a molecule in a solution of particular
pH.
 To determine the [A-] or salt concentration or the amount of ionized chemical and the [HA] or acid
concentration or the amount of unionized chemical- Knowing the ratio of the ionized and unionized
forms of the molecule is important for certain reactions as well as for pharmaceutics related problems
as a lot of biological functions such as absorption and excretion of drugs depends upon the
ionized/unionized state of the drug

12. Demonstration of buffering action of a
weak acid and its conjugate base
 Titrate weak acid with a base (usually NaOH)
 e.g. take 10ml of 0.1M solution of acetic acid and note its pH
 Add 10 ml NaOH 0.1M in small measured amounts at a time
 Note rise in pH after each addition till all 10ml of NaOH solution is consumed

13.  A pH value of 4.8 was seen when 5ml
of 0.1M NaOH was added
 This means that one half of the acetic
acid is dissociated i.e. at this pH 50%
of original acetic acid is present as
CH3COOH and 50% as CH3COO-
(acetate ion)
 This point is called the pKa of the acid
 At this point maximum buffering
action is seen
5 10

14.  Diprotic Acids (having two hydrogens e.g H2CO3) will have two pKa values;
one for each of the hydrogen ions
 Triprotic Acids (like phosphoric acid e.g. H3PO4) have three pKa values

15. BUFFERS
 A buffer solution is a solution that resists any change in its pH value on addition of acid or alkali
 The process by which the added H+ or OH– are removed is called as buffering action
 The extent of resistance of change in pH by buffer is called its buffering capacity

16. Buffers in the body
1. Principal buffers of extracellular fluids are
 Bicarbonate buffer
 Phosphate buffer
 Protein buffer
2. Principal intracellular buffers are
 Phosphate buffer
 Protein buffers (Hb)

17. Quiz
 An important buffer system of the blood plasma is NaHCO3/H2CO3. Their
ratio under normal circumstances is 20/1. Find out the pH of the plasma.
The pKa of H2CO3 is 6.1
 Note:
log of 20 = 1.3

18.  pH of plasma = pKa + log Base
Acid
= 6.1 + log 20/1 = 6.1 + 1.3 = 7.4