Fundamentals of Chemistry: Acids, Bases, and pH

Laboratory Session: The first laboratory for the Biochemistry module occurs this Thursday.
Requirements: Students must complete Health and Safety videos/quizzes, review Learning Science virtual simulations, and bring a paper copy of the proforma.
Attendance: Late arrivals for the health and safety briefing will be barred from the session.
Assessment: The Bradford Assay practical is formative, but these skills will be assessed in a summative in-class practical test in January 2026.

pH Definition: Stands for ‘power of hydrogen’; a measure of the activity of dissolved hydrogen ions ( $H^+$ ) or hydronium ions ( $H_3O^+$ ).
Dissociation of Water: $H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-$ or $H_2O \rightleftharpoons H^+ + OH^-$ .
Solution States at $25\,^\circ C$ :
- Neutral: $[H^+] = [OH^-]$ ; pH is $7.0$ .
- Acidic: [H^+] > [OH^-]; pH is < 7.0.
- Basic: [OH^-] > [H^+]; pH is > 7.0.
Environmental Impact: Pure water exposed to the atmosphere absorbs $CO_2$ to form carbonic acid, lowering pH to approximately $5.7$ .

The Scal: Ranges from $0$ (very acidic) to $14$ (very basic).
Logarithmic Nature: The scale is logarithmic; a difference of 1 pH unit represents a 10-fold change in $[H^+]$ .
Formula:
- $pH = -\log[H_3O^+]$
- $pH = -\log[H^+]$
Biological pH Examples:
- Gastric Acid: $0.7$
- Lysosome: $5.5$
- Arterial Blood Plasma: $7.4$
- Pancreas Secretions: $8.1$

Strong Acids: Ionize completely in solution (e.g., Hydrochloric acid ( $HCl$ ), Nitric acid ( $HNO_3$ ), Sulphuric acid ( $H_2SO_4$ )).
Weak Acids: Dissociate only partially; an equilibrium is reached between the acid and its conjugate base (e.g., Ethanoic acid ( $CH_3COOH$ ), Carbonic acid ( $H_2CO_3$ )).

Definition: $K_a$ (or equilibrium constant $K_{eq}$ ) measures the strength of an acid.
Formula: $K_a = \frac{[H^+] \cdot [X^-]}{[HX]}$
Interpretation: A higher $K_a$ indicates a stronger acid.
Weak Acid pH Calculation: $[H^+] = \sqrt{K_a \cdot [HX]}$ , then apply $pH = -\log[H^+]$ .

Strong Bases: Hydrolyze completely (e.g., Sodium hydroxide ( $NaOH$ ), Potassium hydroxide ( $KOH$ )).
Weak Bases: Only partially ionize (e.g., Ammonia ( $NH_3$ )).
Water Ionization Constant ( $K_w$ ): $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$ .
p-Scale Relationships: $pK_w = pH + pOH = 14$ .
Neutralization: Reaction between acid and base produces a salt and water: $HCl(aq) + NaOH(aq) \rightarrow H_2O(l) + NaCl(aq)$ .

Function: Solutions that resist large changes in pH when small amounts of acid or base are added.
Composition: A mixture of a weak acid and its conjugate base (or salt) or a weak base and its conjugate acid.
Mechanism: Added $H^+$ or $OH^-$ ions are neutralized as the equilibrium shifts to accommodate the change.
Henderson–Hasselbalch Equation: Used to estimate buffer pH. Optimal buffering occurs when $pH = pK_a$ .
Biological Importance: Essential for enzyme function. Blood plasma is buffered to pH $7.35 - 7.45$ by the bicarbonate-carbonic acid system.

Q1: pH is calculated via which of the following? Answer: c) $-Log_{10} H^+$ ion concentration.
Q2: If a solution is described as a weak base, its properties will be? Answer: a) pH > 7 and partially ionised.
Q3: You are using a hydrochloric acid solution of pH $2.5$ . The closest estimation of the original concentration is? Answer: c) $+0.003\text{ moles}$ .
Q4: Which statement about the acidity constant, $K_a$ , is true? Answer: d) The lower the $K_a$ value the weaker the acid.
Q5: Which statement is incorrect? Answer: d) All the equations given in the lecture need to be memorised (Note: Slide implies understanding is key, but usually formulas are provided or specific ones targeted for exams).