Free MCAT study guide — Chemical and Physical Foundations of Biological Systems
Acid-base chemistry is one of the most frequently tested topics on the MCAT, appearing in both general chemistry and biochemistry contexts. You need to understand the Bronsted-Lowry definition (acids donate protons, bases accept protons), the Lewis definition (acids accept electron pairs, bases donate electron pairs), and how these apply to organic and biological reactions. Key quantitative concepts include Ka, Kb, pKa, pH, pOH, and the relationship Kw = Ka x Kb = 1.0 x 10^-14 at 25 degrees Celsius. You must be comfortable with the Henderson-Hasselbalch equation (pH = pKa + log[A-]/[HA]) and its applications to buffer systems, titration curves, and amino acid chemistry. Buffer capacity, the selection of appropriate buffers, and the physiological bicarbonate buffer system are high-yield topics. Titration curves for monoprotic and polyprotic acids (including amino acids) require you to identify equivalence points, half-equivalence points (where pH = pKa), and buffer regions. The MCAT tests both conceptual understanding and quantitative problem-solving, so practice both.
Strong acids (HCl, HBr, HI, H2SO4, HNO3, HClO4) completely dissociate in water; their conjugate bases are negligible bases. Strong bases (LiOH, NaOH, KOH, Ca(OH)2, Ba(OH)2) also completely dissociate. Weak acids partially dissociate and have Ka values less than 1 (and pKa values greater than 0). The smaller the Ka (higher the pKa), the weaker the acid. For weak acids, the equilibrium HA <-> H+ + A- is described by Ka = [H+][A-]/[HA]. For weak bases, Kb = [BH+][OH-]/[B]. The relationship Ka x Kb = Kw (1.0 x 10^-14) links conjugate acid-base pairs, and pKa + pKb = 14. To calculate pH of a weak acid solution, use the approximation [H+] = sqrt(Ka x C) when Ka << C. Percent ionization = ([H+]/C) x 100% and increases with dilution. Polyprotic acids (H2SO4, H3PO4, amino acids) have multiple Ka values, with Ka1 >> Ka2 >> Ka3, meaning each successive proton is harder to remove. The MCAT expects you to estimate pH values using logarithm approximations rather than a calculator.
The Henderson-Hasselbalch equation, pH = pKa + log([A-]/[HA]), is essential for buffer calculations and titration analysis. When [A-] = [HA], the log term is zero and pH = pKa -- this is the half-equivalence point in a titration and the point of maximum buffer capacity. When [A-] > [HA], pH > pKa (more basic). When [HA] > [A-], pH < pKa (more acidic). For amino acids, the Henderson-Hasselbalch equation applies to each ionizable group independently. The isoelectric point (pI) is the pH at which the amino acid has no net charge: for amino acids without charged side chains, pI = (pKa1 + pKa2)/2. For acidic amino acids (Asp, Glu), pI = (pKa1 + pKaR)/2, averaging the two lowest pKa values. For basic amino acids (Lys, Arg, His), pI = (pKa2 + pKaR)/2, averaging the two highest pKa values. At pH < pI, the amino acid has a net positive charge (migrates toward cathode in electrophoresis); at pH > pI, it has a net negative charge (migrates toward anode).
A buffer is a solution that resists pH change upon addition of small amounts of acid or base. Buffers consist of a weak acid and its conjugate base (or a weak base and its conjugate acid) in roughly equal concentrations. The effective buffer range is pH = pKa plus or minus 1. Buffer capacity depends on the absolute concentrations of the buffer components -- more concentrated buffers resist pH change better. When strong acid is added to a buffer, it reacts with the conjugate base (A- + H+ -> HA). When strong base is added, it reacts with the weak acid (HA + OH- -> A- + H2O). The bicarbonate buffer system is the most important physiological buffer: CO2 + H2O <-> H2CO3 <-> H+ + HCO3-. The lungs regulate CO2 (volatile acid), and the kidneys regulate HCO3-. Despite its pKa of 6.1 (seemingly far from blood pH 7.4), the bicarbonate system is effective because the lungs continuously adjust CO2 levels. The phosphate buffer system (H2PO4-/HPO4^2-, pKa = 6.8) is important intracellularly, and proteins (especially hemoglobin) also serve as buffers through their ionizable amino acid side chains.
A titration curve plots pH versus volume of titrant added. Titrating a weak acid with a strong base produces a curve with a buffer region (where pH changes slowly, centered at pH = pKa), a half-equivalence point (pH = pKa, halfway to equivalence), and an equivalence point (where moles of acid = moles of base, pH > 7 due to hydrolysis of the conjugate base). The initial pH is determined by the weak acid alone. The equivalence point pH is determined by the conjugate base concentration. Indicators should have pKa values near the equivalence point pH. For diprotic acids, there are two buffer regions, two half-equivalence points, and two equivalence points. Amino acid titration curves are particularly important: a simple amino acid like glycine shows two buffering regions (pKa1 around 2.3 for the carboxyl group, pKa2 around 9.6 for the amino group) and the isoelectric point is the pH midway between. Amino acids with ionizable side chains (Asp, Glu, Cys, Tyr, Lys, Arg, His) have three pKa values and three buffering regions.
The Lewis definition expands acid-base chemistry beyond proton transfer. A Lewis acid is an electron pair acceptor (electrophile), and a Lewis base is an electron pair donor (nucleophile). All Bronsted acids are Lewis acids, but not all Lewis acids are Bronsted acids. Metal cations (Fe3+, Zn2+, Mg2+) are Lewis acids because they accept electron pairs from ligands in coordination complexes. BF3 and AlCl3 are classic Lewis acids with empty orbitals. In organic chemistry, Lewis acid-base interactions are fundamental: carbonyl carbons are electrophilic (Lewis acid), and nucleophiles like hydroxide, cyanide, and Grignard reagents are Lewis bases. In biochemistry, metal ion cofactors in enzymes act as Lewis acids to stabilize negative charges in transition states (e.g., Zn2+ in carbonic anhydrase and carboxypeptidase). The MCAT may ask you to identify Lewis acids/bases in reaction mechanisms or enzyme active sites.
The solubility product constant (Ksp) describes the equilibrium between a solid ionic compound and its dissolved ions. For a salt AxBy, Ksp = [A]^x[B]^y. If the ion product (Q) exceeds Ksp, precipitation occurs; if Q < Ksp, the solution is unsaturated. The common ion effect reduces the solubility of a salt when one of its ions is already present in solution -- for example, adding NaCl to a saturated solution of AgCl reduces Ag+ solubility because the increased Cl- concentration shifts the equilibrium toward precipitation. pH affects the solubility of salts containing basic anions: metal hydroxides and carbonates are more soluble in acidic solutions because H+ reacts with OH- or CO3^2-, removing them from the equilibrium and shifting it toward dissolution. The MCAT often tests Ksp calculations in the context of kidney stones (calcium oxalate), bone mineral (hydroxyapatite), and biological precipitation phenomena.
Strong acids: HCl, HBr, HI, H2SO4, HNO3, HClO4. Everything else is weak.
pH = -log[H+]; pOH = -log[OH-]; pH + pOH = 14 at 25 degrees Celsius.
Henderson-Hasselbalch: pH = pKa + log([A-]/[HA]). At the half-equivalence point, pH = pKa.
Buffer range is pKa +/- 1. Maximum buffer capacity when [A-] = [HA].
The bicarbonate buffer: CO2 + H2O <-> H2CO3 <-> H+ + HCO3-. pKa = 6.1.
At the equivalence point of a weak acid-strong base titration, pH > 7.
Amino acid pI: average of the two pKa values flanking the zwitterionic form.
Kw = Ka x Kb = 1.0 x 10^-14; pKa + pKb = 14.
Lewis acids accept electron pairs; Lewis bases donate electron pairs.
The common ion effect decreases solubility by adding an ion already present.
Percent ionization of a weak acid increases with dilution.
For polyprotic acids, Ka1 >> Ka2 >> Ka3; the first proton is easiest to remove.
Blood pH is maintained at 7.35-7.45 by the bicarbonate buffer, respiratory, and renal systems.
Calculating pH of a strong acid as if it were weak -- strong acids completely dissociate, so [H+] = concentration directly.
Forgetting that the equivalence point pH is NOT 7 for weak acid-strong base titrations (it is above 7).
Using Henderson-Hasselbalch when the acid is strong -- H-H only applies to weak acid/conjugate base pairs.
Confusing pKa with pH -- pKa is a property of the acid, pH is a property of the solution.
Averaging the wrong two pKa values when calculating pI for amino acids with ionizable side chains.
Forgetting that adding water to a buffer dilutes both components equally, so the ratio stays the same and pH barely changes.
Confusing Bronsted-Lowry and Lewis definitions -- H+ is both a Bronsted acid and a Lewis acid.
Practice estimating pH without a calculator: know that log(2) is approximately 0.3, log(3) is approximately 0.5, and log(5) is approximately 0.7. For a weak acid with Ka = 1.8 x 10^-5 at 0.1 M concentration, estimate [H+] = sqrt(1.8 x 10^-6) which is approximately 1.3 x 10^-3, so pH is approximately 2.9. The MCAT expects this level of mental math. Practice with various Ka values until you are fast and accurate.
Draw titration curves from scratch for monoprotic acids, diprotic acids, and amino acids. Label the initial pH, buffer region, half-equivalence point (pH = pKa), equivalence point, and final pH. For each region, identify the dominant species in solution. Practice predicting the charge on an amino acid at any given pH relative to its pI, and determine the direction of migration in electrophoresis. These skills are tested repeatedly on the MCAT.
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