Buffer system Bicarbonate Buffer Phosphate Buffer Protein Buffer.

2.  Buffer is a solution that resist changes in pH
 Buffer solutions contain:
- weak acid and salts
- weak base and salts
 There are three important systems in our
body:
1. Bicarbonate buffer system
2. Phosphate buffer system
3. Protein buffer system

3.  Is an acid-base homeostatic mechanism involving the
balance of carbonic acid, bicarbonate ion and carbon
dioxide in order to maintain pH in the blood and
duodenum, among other tissues, to support proper
metabolic function.
CO₂ + H₂O = H₂CO₃ + H₂O = HCO₃⁻ + H₃O⁺
 The pH balanced by the presence of both weak acid and it’s
conjugate base (HCO₃⁻).
 H₂CO₃ + H₂O ← HCO₃⁻ + H⁺ equilibrium shifts to the left,
however there is HCO₃⁻ present for the H⁺ to react with,
thus removing the excess H⁺ ions. If a process releases base
into the blood, which would react and lower the
concentration of the H⁺ , the equilibrium shifts to the right,
as long as there is H₂CO₃, thus replacing the lost H⁺ .

4.  Describes the derivation of pH as a measure of acidity,
using p𝐾 𝑎, the negative log of the aacid dissociation
constant, in biological and chemical systems. So that, the
equation is also used in estimating the pH of a Buffer
systems
pH = 𝑝𝐾 𝑎 + log₁₀
[A⁻ ]
[HA]
[HA] – molar conc. of undissociated acid
[A⁻ ] – molarity of this acid’s conjugate base
𝑝𝐾 𝑎 - is log₁₀ 𝐾 𝑎, and 𝐾 𝑎 is acid dissociation const.
 For bases: pOH = 𝑝𝐾 𝐵 + log₁₀
[BH⁺ ]
[B]
 For pH of basic solutions: pH = 𝑝𝐾 𝑎 + log₁₀
[B ]
[BH⁺]

5.  The phosphate buffer system [𝐻𝑃𝑂₄−2
]/ [H₂PO₄⁻]
plays a role in plasma and erythrocytes
H₃PO₄ = H₂PO₄⁻ + H⁺ = 𝐻𝑃𝑂₄−2
+ H⁺ = 𝐻𝑃𝑂₄−ᶟ + H⁺
 Protein contains –COO group, which is like an acetate ions
(CH₃COO⁻) can act as protein acceptors
 Protein also contain NH₃⁺ group which is like NH₄⁺
ammonium ions, can donate protons
 If acid cames into blood, hydronium ion can be neutralized
by the –COO groups
 -COO + H₃O → -COOH + H₂O
 If base is added, it can be neutralized by the NH₃⁺ groups
-NH₃⁺ + OH⁻ → -NH₂ + H₂O

6.  Acidic buffers – buffers with a pH of below 7 are formed from a
solution containing a weak acid and the salt of the same weak
acid (for ex. Ethanoic acid and sodium ethanoate). This gives a
buffer solution with a pH less then 7. CH₃COOH ↔ CH₃COO⁻ +
H⁺ equilibrium shifts to the left (large store of ethanoic acid
molecules)
CH₃COONa ↔ CH₃COO⁻ + Na⁺equilibrium shifts to the right
(large store of ethanoate ions) addition of small quantities of acid have no
effect on the pH, however H⁺ ions added react with the excess ethanoate ions in
eq.2 and are removed from the solution as ethanoic acid molecules. In case of addition
a base in small quantities, OH⁻ reacts with the hydrogen from eq.1 removing from the
right side. However there is a large reservoir of ethanoic acid on the left side, it’s able
to dissociate and make more H⁺ ions restoring the pH.
Basic buffers – contains a weak base and one of it’s salts. For ex.
Ammonia solution and ammonium chloride.
NH₃ + H₂O ↔ NH₄⁺ + OH⁻ equilibrium shifts to the left (large store of
ammonia molecules)
NH₄Cl ↔ NH₄⁺ + Cl⁻ equilibrium shifts to the right

7.  Acidosis refers to an excess of acid in the blood that
causes the pH to fall below 7.35
 Alkalosis refers to an excess of base in the blood
that causes the pH to rise above 7.45
 Buffering systems that resist changes in pH also
contribute to the regulation of acid and base
concentrations.