Introduction to pH, Buffers and Isotonic Solutions
pH and buffers are fundamental to pharmaceutical formulation. Every injectable, eye drop, nasal spray, and biological product requires precise pH control for stability, patient comfort, and drug activity. Buffer calculations using the Henderson-Hasselbalch equation are frequently tested in exams.
Syllabus & Topics
- 1Sorensen’s pH Scale: pH = -log[H⁺]. History and significance. pH scale from 0-14 (at 25°C). Neutral pH = 7 (pure water). Acidic pH < 7, Basic pH > 7.
- 2pH Determination – Electrometric Method: Glass electrode (H⁺ selective), Reference electrode (Calomel electrode), pH meter. Principle: Nernst equation. E = E° – (RT/nF) log [H⁺]. Advantages: accurate, applicable to colored/turbid solutions.
- 3pH Determination – Colorimetric Method: Using pH indicators (e.g., Litmus, Phenolphthalein, Universal indicator). pH indicator paper, Comparators. Limitations: colored solutions, interference.
- 4Applications of Buffers: To maintain pharmaceutical product pH; to maintain biological pH; to control drug ionization and permeability.
- 5Buffer Equation (Henderson-Hasselbalch): pH = pKa + log([A⁻]/[HA]) for weak acid buffer. pOH = pKb + log([BH⁺]/[B]) for weak base buffer. Derivation from Ka expression.
- 6Preparation of Buffers: Acetate buffer (acetic acid + sodium acetate, pH 3.5-5.7), Phosphate buffer (pH 5.8-8.0), Citrate buffer, Borate buffer, Tris buffer.
- 7Buffer Capacity (β): β = ΔB/ΔpH = the number of moles of strong acid/base added per litre to change pH by 1 unit. Maximum buffer capacity at pH = pKa (β_max = 0.576 × C). Van Slyke equation: β = 2.303 × Ka[H⁺]C/(Ka + [H⁺])².
- 8Buffers in Pharmaceutical Systems: Ophthalmic solutions (pH 7.3-7.4 for eye comfort), Injections (pH 7.4 for blood compatibility), Oral liquids.
- 9Buffers in Biological Systems: Blood buffer systems: (1) Bicarbonate buffer (HCO₃⁻/CO₂ – most important), (2) Haemoglobin buffer, (3) Phosphate buffer, (4) Plasma protein buffers. Normal blood pH = 7.35-7.45.
- 10Acidosis (pH < 7.35) and Alkalosis (pH > 7.45).
- 11Buffered Isotonic Solutions: Isotonic solutions have same osmotic pressure as blood (~285-295 mOsmol/L) or lacrimal fluid. Hypotonic solutions cause cell swelling/haemolysis. Hypertonic solutions cause cell shrinkage/crenation.
- 12Methods for adjusting isotonicity: Sodium chloride equivalent method, Freezing point depression method. NaCl equivalent for common drugs.
Learning Objectives
Frequently Asked Questions (FAQs)
Q1. What is the Henderson–Hasselbalch Equation?
The Henderson–Hasselbalch equation for a weak acid buffer is expressed as pH = pKa + log([A⁻]/[HA]). It explains the relationship between pH, pKa, and the ratio of conjugate base ([A⁻]) to weak acid ([HA]). When the concentrations of [A⁻] and [HA] are equal, pH equals pKa. Increasing the ratio [A⁻]/[HA] increases the pH. A buffer is most effective when the pH is within ±1 unit of its pKa, as it can resist changes in acidity or alkalinity within this range.
Q2. What is Buffer Capacity (β)?
Buffer capacity (β) is defined as β = ΔB/ΔpH and represents the resistance of a buffer solution to changes in pH upon the addition of acid or base. It indicates how much strong acid or base must be added to change the pH by one unit. Buffer capacity is maximum when pH equals pKa, and it increases with higher total buffer concentration (C). Practically, a buffer is most effective within the range of pKa ± 1 pH unit.
Q3. What Are the Most Important Buffer Systems in Blood?
The most important buffer system in blood is the bicarbonate/carbonic acid system (H₂CO₃/HCO₃⁻, pKa = 6.1), as its components are regulated by the lungs (which control CO₂) and the kidneys (which regulate bicarbonate levels). Other important systems include the haemoglobin buffer, which plays a major role inside red blood cells; the phosphate buffer system (H₂PO₄⁻/HPO₄²⁻), which is more significant in intracellular fluid and kidneys; and plasma protein buffers, which contribute through ionizable amino acid groups.
Q4. What is an Isotonic Solution?
An isotonic solution has the same osmotic pressure as blood plasma and lacrimal fluid, approximately 285–295 mOsmol/L. Such solutions do not cause any net movement of water across cell membranes. Examples include 0.9% sodium chloride (normal saline) and 5% dextrose solution. Hypotonic solutions cause hemolysis, where cells swell and burst due to water influx, while hypertonic solutions cause crenation, where cells shrink due to water loss.
Q5. Why Must Ophthalmic Solutions Be Buffered?
Ophthalmic solutions must be buffered to maintain a pH close to that of lacrimal fluid (approximately 7.4) and to ensure isotonicity around 290 mOsmol/L for patient comfort. Solutions with extreme pH values (below 3.5 or above 10.5) can cause irritation, pain, and reflex tearing, which may dilute the drug and reduce therapeutic effectiveness. Proper buffering also enhances drug stability, as many active pharmaceutical ingredients are sensitive to pH changes.