ELECTROLYTE

ELECTROLYTE

OVERVIEW

  • This lecture is part of a series of 3 lectures.

  • Focuses on understanding the behavior of electrolytes in aqueous solutions, relevant in pharmacy and pharmaceutical science.

BEFORE WE START

Importance of pH

  • pH is crucial in pharmacy and pharmaceutical sciences due to its impact on:

    • Medicines: such as antacids and proton pump inhibitors.

    • Disease states: e.g., respiratory acidosis/alkalosis.

Pharmaceutical Sciences Considerations

  • Most drugs are weak acids or bases, affecting:

    • Pharmacological Activity: pharmacokinetics and pharmacodynamics.

    • Drug Solubility: particularly in different pH levels.

    • Drug Stability: impacts effectiveness over time.

    • Drug Partitioning: influences absorption and distribution in the body.

    • Drug-Drug Interactions: can alter stomach pH.

    • Example: Ulipristal acetate interactions discussed in BNF.

Pharmacy Practice Implications

  • Important dispensing considerations include:

    • Advisory Labels: e.g., do not take indigestion remedies 2 hours before or after other specific medicines.

ACIDS AND BASES

  • Electrolytes: Substances that produce ions in water.

    • Cations: positively charged ions.

    • Anions: negatively charged ions.

  • Brönsted-Lowry Theory:

    • Acids: Proton donors (e.g., HA + H₂O ⇌ H₃O⁺ + A⁻).

    • Bases: Proton acceptors (e.g., B + H₂O ⇌ OH⁻ + BH⁺).

SELF IONISATION OF WATER

  • Hydrogen ions (H⁺) do not exist freely in solution; they are solvated in water, forming hydronium ions (H₃O⁺).

  • Reaction: H₂O + H⁺ ⇌ H₃O⁺.

  • Equilibrium:

    • The amount of ionization of water is very low.

    • Equilibrium constant: $K_w = [H^+][OH^-] = 1 imes 10^{-14}$ at 25°C.

pH CALCULATIONS

  • Definition: $pH = - ext{log}[H^+]$.

  • The pH scale indicates acidity/alkalinity:

    • Neutral solution: $[H^+] = 10^{-7}$ M → $pH = 7$.

    • Acidic solution: $[H^+] > 10^{-7}$ M → $pH < 7$.

    • Alkaline solution: $[H^+] < 10^{-7}$ M → $pH > 7$.

Relationship of pH and [H^+]

  • Each change in 1 pH unit corresponds to a 10-fold change in [H^+].

STRONG VS WEAK ACIDS AND BASES

Strong Acids

  • Complete Ionization: All molecules dissociate into ions (e.g., HCl → H⁺ + Cl⁻).

    • Calculation example for 0.1M HCl:

    • $pH = - ext{log}(0.1) = 1$.

Weak Acids

  • Partial Ionization: Only some molecules dissociate (e.g., acetic acid).

Strong Bases

  • Example of strong base calculation for 0.1M Ba(OH)₂:

    • Dissociation: Ba(OH)₂ → Ba²⁺ + 2OH⁻.

    • $pOH = - ext{log}[OH⁻] = - ext{log}(0.2) = 0.7$.

    • $pH = 14 - pOH = 13.3$.

CONJUGATE ACIDS AND BASES

  • Weak acid dissociates to produce $H_3O^+$ and its conjugate base:

    • Example: Acetic acid (HA) gives acetate (A⁻) as a conjugate base.

  • Acidity Constant ($K_a$) for weak acid dissociation:

    Ka = \frac{[H3O^+][A^-]}{[HA]}

DISSOCIATION CONSTANTS

  • Weak bases follow a similar model with a basicity constant ($K_b$).

  • Acid equilibrium constant equations can be simplified:

    pKa + pKb = pK_w = 14

% IONIZATION

  • Ionization extent depends on:

    • pKa/pKb of the compound

    • pH of the environment.

  • For weak acids:

    • Example with Naproxen ($pK_a = 4.2$):

    • % Ionization formula:

      ext{Percentage Ionization} = \frac{100}{1 + \text{antilog}(pK_a - pH)}

    • Calculated percentages at varying pH levels.

SUMMARY OF KEY TERMS

  • Know definitions and conversions for:

    • Electrolytes, pH, pKw, Ka, pKa, Kb, pKb

    • Acid-base strength relationship

    • Calculating pH, % ionization.

WHAT'S COMING UP

  • Future topics to explore:

    • Acid-base titrations

    • Buffers

    • Electrolytes

    • Log P