Acids & Bases Overview Quiz

Acids & Bases, pH and Buffers

• Blood pH levels range from 6 to 9, with 7.35-7.45 being the normal range; values outside this range indicate acidosis (low pH) or alkalosis (high pH).
• Death can result from severe acidosis or alkalosis, as biochemical reactions are sensitive to pH.
• Buffer solutions maintain a constant pH in biological systems, crucial for proper bodily function. Examples include saliva, sweat, blood, stomach, intestine, and urine.
• The pH scale is logarithmic, meaning pH values change by powers of 10 (e.g., pH 1 is 10 times more acidic than pH 2).
• The pH scale ranges from 0 to 14, with neutral pH = 7.
• pH = -log[H+]
• pOH = -log[OH-]
• K_w= [H⁺][OH^-] = 1 x 10^-14

Learning Outcomes

• Students should be able to explain acids, bases, pH, and pOH.
• Students should be able to define strong and weak acids and their importance in biology.
• Students should understand water auto-ionization.
• Students should know about conjugate acid/base pairs (Ka, Kb, pKa, pKb)
• Students should be able to calculate pH.
• Students should explain buffer systems and the factors affecting choice of buffer, using the Henderson-Hasselbalch equation. and perform buffer calculations.

Suggested Reading

• Crowe & Bradshaw, Chemistry for the Biosciences, 3rd Edition (e-book in online library)
• "Catch up Chemistry" (online resource)
• Chapter 17: "Acids, bases and buffer solutions"

Biochemical Reactions in the Body

• Biochemical reactions in the body are highly sensitive to pH.
• Buffer solutions are crucial for maintaining this pH, and this includes saliva, sweat, blood, stomach, intestines, and urine.

Acids and Bases

• Electrolytes, upon addition to water, dissociate into cations and anions.
• Brønsted-Lowry Acid is a proton donor (HA → H⁺ + A^–).
• Brønsted-Lowry Base is a proton acceptor (B + H⁺ → BH⁺).
• Autoionization is a reaction between identical neutral molecules to form a cation and an anion (e.g., 2H₂O ⇌ H₃O⁺ + OH^-, Kw = [H⁺][OH^-]).

Equilibrium Constants

• For any reversible reaction with reactants A & B and products C & D, the equilibrium constant is Keq = [C][D]/[A][B].
• The water equilibrium is a reverse reaction (2H₂O ⇌ H₃O⁺ + OH^–), which lies predominantly to the left, implying a small degree of dissociation.
• The equilibrium constant for water dissociation, K_w, is [H⁺][OH^-]/[H₂O]. However, the water concentration is assumed to be constant, therefore K_w = [H⁺][OH^-) = 1 x 10^-14 M².

pH

pH scale is a logarithmic scale expressing the concentration of H+ ions in solution.

Calculating pH

• pH = -log[H+].

Strong/Weak Acids/Bases

• Strong acids completely dissociate in solution, while weak acids only partially dissociate.
• Strong bases completely dissociate in solution, while weak bases only partially dissociate.
• Note: the strength of an acid or base is not dependent on concentration, but rather the equilibrium constant and pK values.

Conjugate Acids and Bases

• A conjugate base is what remains of an acid after it has donated a proton.
  • Weak acid dissociates to give H⁺ and its conjugate base.
  • E.g. CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO^- conjugate base.
• A conjugate acid is formed when a base accepts a proton.
  • E.g. NH₃ + H₂O ⇌ NH₄⁺ + OH^-.

Acid Equilibrium Constant

• K_a = [H₃O⁺][A^-]/[HA]

Base Equilibrium Constant

• K_b = [BH⁺][OH^-]/[B]

pKa, pKb

• pK_a = -logK_a
• pK_b = -logK_b

Polyprotic Acids

• Polyprotic acids can donate more than one proton (e.g., H₂SO₄, H₃PO₄).
• Each dissociation step has its own equilibrium constant.

Buffers

• Buffer solutions resist changes in pH when an acid or base is added.
• Often contain substances with acidic and basic forms (Conjugate pairs of weak acids).
• The Henderson-Hasselbalch equation gives the relationship between pH, pKa, and the concentrations of the acid and its conjugate base. pH = pKa + log[base]/[acid].
• Buffers work best around a pH range of ±1 of their pKa.

Choice of Buffer Solution

• The choice of a buffer solution depends on the desired pH range. Biological fluids commonly use a phosphate buffer as it works well at slightly above neutral pH (e.g pH 7).

Acids & Bases in Biological Systems

• Many biological processes require specific pH levels.
• For example, the small intestine has a pH of around 6, to ensure proper functioning of digestive enzymes and absorption of nutrients.
• The pH of the stomach is extremely low (1 to 2) to help break down food.
• Certain drugs must be in a non-ionized form to be absorbed through cell membranes.
• pKa's and ionisation play key roles in medicine absorption; for example, aspirin is better absorbed in the stomach.