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5A · The chemistry of life
Water and its solutions
Water's polarity and hydrogen bonding make it the universal biological solvent — and set up the chapter's high-yield core: acid–base chemistry (pH, strong vs. weak, buffers, titrations) and solutions (solubility, Ksp, colligative properties).
Properties of water
Water is a bent, polar molecule that forms hydrogen bonds, giving it cohesion, a high boiling point and heat capacity, and its role as the solvent of life.
Oxygen's electronegativity pulls electron density from the two hydrogens, making each O–H bond polar and the bent molecule a net dipole. Each water can hydrogen-bond to up to four neighbors, which explains its anomalously high melting/boiling points, high specific heat (temperature buffering in organisms), high heat of vaporization (sweating), surface tension, and the fact that ice floats (the H-bond lattice is less dense than liquid water). As a solvent it dissolves ions and polar molecules ("like dissolves like") while excluding nonpolar ones — the hydrophobic effect that folds proteins and forms membranes.
Polarity and hydrogen bonding
A hydrogen bond is a strong dipole–dipole attraction between an H bonded to N, O, or F and a lone pair on another N/O/F. It's much stronger than ordinary dipole forces but far weaker than a covalent bond.
Hydrogen bonding requires both a donor (H on N/O/F) and an acceptor (a lone pair on N/O/F). It is the strongest intermolecular force on the MCAT and underlies water's properties, the pairing of DNA bases, and the secondary structure of proteins. Crucially it is an intermolecular force, not an actual chemical bond — a distinction the exam tests directly (see 5B).
Cohesion, heat capacity, and the solvent role
Hydrogen bonding gives water cohesion (surface tension, capillary action), a high heat capacity (thermal buffering), a high heat of vaporization (evaporative cooling), and broad solvent power for ions and polar molecules.
These are all downstream of H-bonding. High specific heat lets organisms resist temperature swings; high heat of vaporization makes sweating an efficient coolant; cohesion and adhesion pull water up xylem and through capillaries. The same polarity that dissolves salts and sugars excludes nonpolar molecules, driving the hydrophobic effect that organizes membranes and protein cores — the single most important consequence for biochemistry.
Acids and bases
Water autoionizes (Kw = [H⁺][OH⁻] = 10⁻¹⁴), defining the pH scale. Acids/bases are strong (fully dissociate) or weak (partial, governed by Ka/Kb); buffers resist pH change near a weak acid's pKa, and titrations track pH as acid meets base.
The pH scale comes straight from water's autoionization: pH = −log[H⁺], pOH = −log[OH⁻], and pH + pOH = 14 at 25 °C. The strong/weak distinction is the conceptual fork — strong acids/bases dissociate completely (pH from concentration directly), weak ones reach an equilibrium set by Ka/Kb. The two highest-yield applications are buffers and titration curves, both of which live on the Henderson–Hasselbalch relationship.
Three definitions of acids and bases
Arrhenius: donates H⁺ / OH⁻ in water. Brønsted–Lowry: proton (H⁺) donor / acceptor — the MCAT default. Lewis: electron-pair acceptor / donor (the most general).
The three definitions widen in scope. Brønsted–Lowry is the workhorse and introduces conjugate acid–base pairs: every acid has a conjugate base (the acid minus its proton). Lewis broadens "base" to any electron-pair donor — so a nucleophile is a Lewis base and an electrophile a Lewis acid, tying this directly to organic reactivity in 5D. A metal cation accepting electron pairs (e.g., in a complex ion) is a Lewis acid even with no proton involved.
pH, pOH, and Kw
pH = −log[H⁺]; Kw = [H⁺][OH⁻] = 10⁻¹⁴ at 25 °C; pH + pOH = 14. Each whole pH unit is a ten-fold change in [H⁺].
Because pH is logarithmic, a solution at pH 3 has 100× the [H⁺] of pH 5. For estimation without a calculator: if [H⁺] = a×10⁻ᵇ, then pH ≈ b − log a (a number a bit below b). Note Kw rises with temperature (autoionization is endothermic), so neutral pH is below 7 in a warm body — a subtle but testable point.
Strong vs. weak; Ka, Kb, and conjugates
Strong acids/bases dissociate completely; weak ones partially, with Ka (or Kb) setting the equilibrium. A lower pKa = stronger acid. The stronger the acid, the weaker its conjugate base (Ka·Kb = Kw).
Memorize the handful of strong acids (HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄) and strong bases (group I/II hydroxides); everything else is weak. For a weak acid, Ka is small and pKa = −log Ka; the smaller the pKa, the more it dissociates. The inverse relationship Ka·Kb = Kw means a very strong acid has a negligibly weak conjugate base (Cl⁻ doesn't affect pH), while a weak acid's conjugate base is a meaningful base.
Don't confuse
Acid strength (degree of dissociation, set by Ka) is not the same as concentration (how much is dissolved). A concentrated weak acid can have a higher pH than a dilute strong acid. The exam reliably separates these.
Buffers and Henderson–Hasselbalch
A buffer is a weak acid plus its conjugate base; it resists pH change. pH = pKa + log([A⁻]/[HA]) (Henderson–Hasselbalch). A buffer works best, and pH = pKa, when [A⁻] = [HA].
Buffers absorb added acid or base by converting between the weak acid and its conjugate base, holding pH nearly constant. Henderson–Hasselbalch shows pH is set by the ratio of base to acid, not the absolute amounts — so dilution barely shifts pH. Buffering capacity is greatest within about ±1 pH unit of the pKa; the body's bicarbonate buffer (pKa ~6.1, plus respiratory CO₂ control) keeps blood near 7.4. The carbonic-acid/bicarbonate system links this to gas exchange in 4B.
Titration curves
A titration plots pH vs. added titrant. At the half-equivalence point pH = pKa (max buffering); at the equivalence point moles of acid = moles of base. A weak-acid/strong-base equivalence point is basic (pH > 7), not neutral.
For a weak acid titrated with strong base, the curve has a buffering plateau (centered at pH = pKa, the half-equivalence point) and a steep jump at equivalence. The equivalence point pH is set by the conjugate: weak acid + strong base → basic salt (pH > 7); strong acid + weak base → acidic salt (pH < 7); strong + strong → neutral. A polyprotic acid shows one plateau and one equivalence step per proton.
Don't confuse
Equivalence ≠ neutral. Equivalence means stoichiometrically equal moles, but the pH there depends on the conjugate's strength — usually not 7 unless both partners are strong. And the half-equivalence point (not equivalence) is where pH = pKa.
Solutions and solubility
Solubility is governed by "like dissolves like." Sparingly soluble salts have a solubility product Ksp, lowered by a common ion. Dissolved particles change a solvent's colligative properties — boiling, freezing, and osmotic pressure.
Whether something dissolves is set by the match between solute and solvent polarity and by entropy. For ionic solids, the equilibrium constant for dissolving is Ksp; adding an ion already in the salt shifts that equilibrium back (the common-ion effect), reducing solubility. Gas solubility follows Henry's law (dissolved gas ∝ partial pressure) and falls with temperature.
Concentration units and dilution
Molarity M = mol solute / L solution is the default. Molality m = mol solute / kg solvent (used for colligative properties; temperature-independent). Mole fraction = mol component / mol total. Dilution: M₁V₁ = M₂V₂.
Pick the unit the problem needs: molarity for solution stoichiometry (but it drifts with temperature as volume changes); molality for boiling-point/freezing-point work, since it's based on solvent mass and is temperature-independent; mole fraction for gas and vapor-pressure problems; normality (equivalents per liter) for acid–base/redox titrations. The dilution relation M₁V₁ = M₂V₂ (moles are conserved when solvent is added) handles every "dilute to…" question.
Solubility, Ksp, and the common-ion effect
For a sparingly soluble salt, Ksp is the equilibrium product of dissolved ion concentrations. Adding a common ion shifts the equilibrium toward the solid (Le Châtelier), lowering solubility.
Compare the reaction quotient Q to Ksp to predict precipitation: Q > Ksp precipitates, Q < Ksp dissolves further. The common-ion effect explains why a salt is far less soluble in a solution already containing one of its ions, and pH can shift solubility when the anion is basic (e.g., hydroxides, carbonates). This is the same equilibrium logic as 5E.
Salts, electrolytes, and hydrolysis
A dissolved salt can make a solution acidic, basic, or neutral depending on its ions: the conjugate of a weak acid or base hydrolyzes water. Strong-acid + strong-base salts are neutral. Strong electrolytes dissociate fully; weak ones partially; nonelectrolytes not at all.
To predict a salt's pH, trace where each ion came from: the cation of a weak base (e.g., NH₄⁺) is acidic; the anion of a weak acid (e.g., acetate, F⁻) is basic; ions from strong acids/bases (Na⁺, Cl⁻) are spectators and leave the solution neutral. So NH₄Cl is acidic, sodium acetate is basic, NaCl is neutral. Electrolyte strength (how fully a solute splits into ions) sets conductivity and feeds the van 't Hoff factor in colligative properties.
Don't confuse
A salt is not automatically pH-neutral. Only strong-acid + strong-base salts are; whenever a salt carries the conjugate of a weak partner, that ion hydrolyzes and shifts the pH.
Colligative properties
Colligative properties depend on the number of dissolved particles, not their identity: boiling-point elevation, freezing-point depression, osmotic pressure (Π = iMRT), and vapor-pressure lowering.
What matters is particle count, captured by the van 't Hoff factor i (NaCl → i ≈ 2, glucose → i = 1). So 1 M NaCl depresses freezing point about twice as much as 1 M glucose. Osmotic pressure (Π = iMRT) drives water across membranes from low to high solute concentration and underlies tonicity, IV-fluid design, and turgor. Vapor pressure of the solvent drops as nonvolatile solute is added (Raoult's law).
Don't confuse
Colligative effects scale with particle count, so you must apply the van 't Hoff factor — 1 M CaCl₂ gives ~3 particles, more effect than 1 M glucose. Comparing molarities without dissociation is the trap.
Worked question
A buffer is prepared from a weak acid (pKa = 4.7) and its conjugate base such that [A⁻] is ten times [HA]. What is the approximate pH?