J1i: Principles Underlying Acid-Base Chemistry

Importance to Regulate pH

  • pH is the negative logarithm to the base 10 of the hydrogen ion [H+]

pH = – log10 [H+]

  • Normal arterial pH = 7.4 = [H+] 40 mol/L
  • The negative logarithmic scale of pH is non-linear & the relationship between pH & [H+] is inverse
Logarithmic scale
  • ∴↓ pH 1 unit = x10 fold ↑ H+ activity
  • ∴↓ pH 7.4 → 7.3 = ↑H+ from 40 → 50mol/L
  • ∴tight regulation [H+] for 2 reasons:
    1. SMALL MOLECULES: completely ionise at neutral pH ∴become trapped in cells of organelles
    2. PROTEINS: perform optimally at a specific pH ∴altering pH effects
      • Enzyme activity
      • Membrane excitability
      • Electron production
      • Hormone release
      • CNS reflexes

Bronsted-Lowry Definition

  • Acid = proton donor
  • Base = proton acceptor
  • H+ = an atom without an electron i.e. a proton
  • Strong acid = an acid with a strong tendency to completely dissociate & discharge its H+ into solution
    • HA (strong acid) ↔ H+ (proton) + A–
  • Strong base = base that reacts powerfully with H+, mopping it up from solution

→ Strong acids & bases always fully dissociate/associate ∴only exist in a charged form

  • Weak acid = less readily release H+
  • Weak base = less readily accepts H+

→ Most acids & bases are weak i.e. only partially dissociate

  • H2CO3 → most important weak acid
  • HCO3– → most important weak base
  • Distinguish strength of acid/base → based on pKa
  • pKA : the negative logarithm of dissociation constant of a substance & the pH where the substance is 50% dissociated

pKA

  • pKa < 4 = strong acid
  • pKa > 12 = strong base
  • pKa 4 – 12 = weak acid / base

VOLATILE ACID = an acid that can only be excreted by lungs

FIXED ACID = an acid that requires excretion by kidneys i.e. cannot be excreted by lungs

Henderson-Hasslebach Approach

  • LAW OF MASS ACTION → the velocity of a chemical reaction is proportional to the active concentrations of the reactions
  • The equilibrium constant (K) indicates which side of the reaction the equilibrium is:
Equilibrium
  • HENDERSON applied the LAW OF MASS ACTION to the equilibrium reaction for carbonic acid

CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3–

k1
    • If [H2O] is large enough to be considered as constant → ∴[H2CO3] = K3 x [CO2]
    • Then substitute a function of the derivable [CO2] in place of [H2CO3] which can’t be measured:
H
    • [CO2] can be calculated from PaCO2 using Henry’s Law
    • K1, K2, K3, K4 are all numerically different constants
  • Then SORENSEN introduced the pH scale
  • It was known that ∆ [HCO3–] reflected accumulation of non-volatile acids
  • pH measurement showed that ∆pCO2 was also affecting pH
  • ∴concept of metabolic/resp acid-base derangement was formed
  • HASSELBACH re-arranged Henderson’s equation:
H2