Biochem Ch 2 Notes: Water, Weak Interactions, Ionization, Buffers, and pH

Principle 1: The solvent properties of water shape the evolution of living things

Water as solvent solubility
- Most small metabolic intermediates, nucleic acids, and proteins are soluble in water.
- Lipid bilayers (likely forerunners of membranes) form spontaneously in water and are stabilized by water interactions.
- The combined effects of multiple weak interactions (hydrogen bonds, ionic interactions, hydrophobic effect) powerfully influence the 3D shape and stability of biomolecules and structures.
Significance of water's interactions
- Hydrogen bonds, ionic interactions, and hydrophobic effect are individually weak but collectively strong in determining structure and stability.
- Water’s solvent properties shaped the evolution of living systems.

Principle 2: Ionization behavior of water and weak acids/bases is captured by equilibrium constants

Ionizable biomolecules
- Most biomolecules are ionizable; their structure and function depend on their ionization state, described by equilibrium constants.
Basic concept
- Acid dissociation constant: for HA ⇌ H+ + A−, Keq = [H+][A−]/[HA] = Ka.
- Stronger acids have larger Ka and typically smaller pKa; weaker acids have smaller Ka and larger pKa.
Ionization constants and strength
- Stronger acids dissociate more, giving larger Ka and smaller pKa values; weaker acids dissociate less, giving smaller Ka and larger pKa values.
Ion product of water
- The ionization of water yields the ion product Kw.
- In pure water at 25 °C: $K_w = [ ext{H}^+][ ext{OH}^-] = 1.0 imes 10^{-14} ext{ M}^2.$
- With water’s autoprotolysis: pure water [H2O] is treated as constant for Ka calculations; Kw can be used to relate pH and pOH via $ext{pH} + ext{pOH} = 14.$
Neutral pH and pH scale
- Neutral pH occurs when [H+] = [OH−] in pure water at a given temperature (25 °C gives pH = 7.0).
- The pH scale describes [H+] via $ext{pH} = - \log [ ext{H}^+].$
Examples of conjugate pairs and pKa values (selected)
- Acetic acid/acetate: Ka = $1.74 imes 10^{-5}$ , pKa = 4.76.
- Carbonic acid/bicarbonate: H2CO3 ⇌ H+ + HCO3−; Ka1 = $1.70 imes 10^{-3}$ , pKa1 ≈ 3.77; HCO3− ⇌ H+ + CO3^{2-}; Ka2 ≈ $6.31 imes 10^{-11}$ , pKa2 ≈ 10.2.
- Glycine (amino acid) conjugates: CO2H side (carboxyl) with pKa ≈ 2.34; amino group with pKa ≈ 9.60.
- Phosphoric acid (triprotic): H3PO4 with pKa1 ≈ 2.14; pKa2 ≈ 6.86; pKa3 ≈ 12.4.
Titration curves and pKa
- Titration curve: pH vs added NaOH; the pH at the midpoint of a buffering region equals the pKa of the weak acid (HA ⇌ H+ + A−).
- Henderson–Hasselbalch relation: $ext{pH} = ext{p}K_a + \log\frac{[A^-]}{[HA]}.$

Principle 3: Buffers resist pH changes in response to added acid/base

Buffers defined
- Buffers are aqueous systems consisting of a weak acid and its conjugate base that resist pH changes when small amounts of acid or base are added.
- A buffer is most effective when [HA] ≈ [A−], i.e., around the pKa of the weak acid.
Buffering region
- The flat region of a titration curve is the buffering region where pH ≈ pKa.
- Example: Acetate buffer (AcOH/AcO−): pH ≈ pKa when [HA] ≈ [A−].
Common buffer systems
- Acetate buffer: CH3COOH ⇌ CH3COO− + H+. pKa ≈ 4.76.
- Phosphate buffer system in cytoplasm: H2PO4− ⇌ HPO4^{2−} + H+. pKa ≈ 6.86; effective between pH ≈ 5.9–7.9.
- Bicarbonate buffer system in blood plasma: CO2(aq) + H2O ⇌ H2CO3 ⇌ H+ + HCO3−; buffering near pH ≈ 7.4 (pKa ≈ 6.1 for the combined system).
Bicarbonate system and CO2 hydration
- Dissolved CO2 equilibrates with aqueous CO2; hydration to H2CO3 is fast: $ext{CO}2( ext{aq}) + ext{H}2 ext{O} ightleftharpoons ext{H}2 ext{CO}3.$
- Kh is the hydration constant: $Kh = rac{[ ext{H}2 ext{CO}3]}{[ ext{CO}2( ext{aq})]}.$
- Combined equilibrium (hydration and dissociation) gives $K{ ext{combined}} = Kh Ka = rac{[ ext{H}^+][ ext{HCO}3^-]}{[ ext{CO}_2( ext{aq})]}.$
- Typical values at 37 °C: $Kh = 3.0 imes 10^{-3} ext{ M}, \, Ka = 2.7 imes 10^{-4} ext{ M}, \, K{ ext{combined}} = 8.1 imes 10^{-7} ext{ M}^2, \, pK{ ext{combined}} \,= \, 6.1.$
Blood pH calculation
- Using typical concentrations: [H2CO3] ≈ 1.2 mM; [HCO3−] ≈ 24 mM.
- Henderson–Hasselbalch yields: $ext{pH} = ext{p}Ka + \log\frac{[ ext{HCO}3^-]}{[ ext{H}2 ext{CO}3]} \approx 6.1 + \log\frac{24}{1.2} \approx 7.4.$
Physiological importance
- Buffering is essential to maintain macromolecule structure and function; blood pH must be kept in a narrow range (roughly 7.3–7.5) to avoid life-threatening effects.

Principle 4: Enzymes function best near physiological pH, with compartment-specific activity

Enzyme activity and pH
- Enzymes evolved to function near neutral pH but can have greater activity and stability at the pH of specific intracellular compartments (low or high pH).
Physiological/clinical relevance
- Extreme deviations from the normal pH (acidic or basic) can lead to loss of enzyme activity and life-threatening conditions.
Practical example of buffering in biology
- Proteins contain histidine residues that buffer effectively near neutral pH (pKa ≈ 6.0).
- The bicarbonate and phosphate buffers help stabilize pH in blood and cells, respectively.

2.1 Weak Interactions in Aqueous Systems: Hydrogen Bonding and Water’s Properties

Hydrogen bonding gives water its unusual properties
- Water has higher melting point, boiling point, and heat of vaporization than many solvents.
- Hydrogen bond: electrostatic attraction between the oxygen atom of one water molecule and the hydrogen of another.
Strength and dynamics of hydrogen bonds
- Bond dissociation energy: ~ $23 ext{ kJ/mol}$ in liquid H2O (≈ 10% covalent, 90% electrostatic).
- Lifetime of a hydrogen bond: ~1–20 picoseconds.
- Each liquid-water molecule forms ~3.4 hydrogen bonds on average; in ice, each water molecule forms 4 hydrogen bonds.
Water’s high melting/boiling behavior
- During melting or evaporation, heat is absorbed; entropy of the system increases.
- For melting: $ext{H}2 ext{O(s)} ightarrow ext{H}2 ext{O(l)}; \, riangle H = +5.9\ ext{kJ/mol}.$ For evaporation: $ext{H}2 ext{O(l)} ightarrow ext{H}2 ext{O(g)}; \, riangle H = +44.0\ ext{kJ/mol}.$
- Since \Delta H > 0, the process is driven by an increase in entropy (\Delta S), making the overall process favorable (negative \\Delta G).
Water interacts with polar solutes
- Water forms hydrogen bonds with electronegative atoms acting as hydrogen acceptors (O, N, etc.).
Hydrophilic vs. hydrophobic vs. amphipathic concepts
- Hydrophilic: dissolves easily in water; typically polar or charged.
- Hydrophobic: nonpolar molecules (e.g., lipids, waxes) poorly soluble.
- Amphipathic: possess both polar/charged and nonpolar regions (e.g., phosphatidylcholine).
Polar, nonpolar, and amphipathic biomolecules (examples from Table 2-1)
- Polar: Glucose; Glycine; Lactate; Glycerol.
- Nonpolar: Typical waxes; long-chain hydrocarbons.
- Amphipathic: Phosphatidylcholine.
Water as solvent and dielectric effects
- Water dissolves salts and charged biomolecules by screening electrostatic interactions; entropy increase upon dissolution is a major driver.
- Water’s dielectric constant at 25 °C: oldsymbol{\epsilon \,\approx \,78.5}. (Compared with nonpolar benzene, oldsymbol{\varepsilon \,= \,4.6}.)
Solubility and gas solubility in water
- Nonpolar gases (CO2, O2, N2) are poorly soluble because dissolution constrains their motion and decreases entropy.
- Nonpolar solutes order water around them, increasing enthalpy and decreasing entropy; overall dissolution is unfriendly (positive \Delta G).
The hydrophobic effect
- Nonpolar regions cluster together in water; polar regions maximize interactions with each other and the solvent.
- Micelle formation is a thermodynamically favorable outcome for amphipathic molecules in water.
Water–macromolecule interactions
- Weak interactions (van der Waals, hydrogen bonds, ionic interactions) stabilize macromolecular structure.
- The cumulative effect of many weak interactions drives proper folding; water molecules can be tightly bound as part of crystal structures.
Proteins and water
- Water can be essential to protein function (e.g., cytochrome f has a chain of bound water molecules that may facilitate proton transport through membranes).
Enzyme-substrate binding and ordered water
- Binding displaces ordered water around substrates; release of this ordered water contributes to stabilization of the enzyme–substrate complex via hydrogen bonding, ionic interactions, and the hydrophobic effect.
van der Waals interactions
- Weak, distance-dependent attractions/repulsions between transient dipoles.
- van der Waals radii provide a measure of how close atoms can approach; the distance of a van der Waals interaction is roughly the sum of the van der Waals radii for the two atoms.
- Example radii (nanometers): H 0.11; O 0.15; N 0.15; C 0.17; S 0.18; P 0.19; I 0.21. Covalent radii (for single bonds) are smaller (H 0.03; O 0.066; N 0.070; C 0.077; S 0.104; P 0.110). The C–C single-bond length is about 0.154 nm (sum of covalent radii).
Importance for macromolecular structure
- Noncovalent interactions are weaker than covalent bonds but continually form/break, enabling dynamic folding and function.
Osmotic effects and water movement
- Solutes alter colligative properties: vapor pressure, boiling point, freezing point, and osmotic pressure.
- Osmotic pressure is the force required to resist water movement across a semipermeable membrane.
- Semipermeable membranes permit water movement but restrict solute diffusion, driving osmosis.

2.2 Ionization of Water, Weak Acids, and Weak Bases

Pure water ionization and proton mobility

Pure water ionization (autoprotolysis)
- Water has a slight tendency to ionize to yield a proton (H+) and a hydroxide (OH−): $ext{H}2 ext{O} ⇌ ext{H}^+ + ext{OH}^-.$ Water protons are rapidly hydrated to hydronium: $ext{H}^+ + ext{H}2 ext{O} ⇌ ext{H}_3 ext{O}^+.$
Proton hopping and mobility
- Protons move via a relay mechanism (proton hopping) through successive hydronium and water molecules, enabling high ionic mobility in aqueous solutions.

Equilibrium constants and acid–base strength

General equilibrium constant (Ka)
- For HA ⇌ H+ + A−, Ka = $rac{[ ext{H}^+][ ext{A}^-]}{[ ext{HA}]}.$
Ionization of water expressed by Kw
- For water ionization: $K_w = [ ext{H}^+][ ext{OH}^-].$
Neutral pH and pH scale
- pH is defined as $ext{pH} = -\, ext{log}_{10}[ ext{H}^+].$
- At 25 °C, neutral water has pH = 7.0 (equal [H+] and [OH−]).
pH and pOH relation
- In all cases: $ext{pH} + ext{pOH} = 14.$
Weak acids and bases and conjugate pairs
- Each weak acid has a conjugate base; the stronger the acid, the larger its Ka and the smaller its pKa; conversely, weaker acids have smaller Ka and larger pKa.
- Examples of conjugate pairs include acetate (pKa 4.76), bicarbonate (pKa1 ≈ 3.77) and carbonate (pKa2 ≈ 10.2), ammonium (pKa ≈ 9.25), glycine zwitterionic forms, and phosphoric acid’s multiple pKa values.

Acids, buffers, and titration

Titration curves reveal pKa
- The pH at the midpoint of a weak acid buffer equals pKa.
Buffers and their concentrations
- Buffers typically consist of a weak acid and its conjugate base; Henderson–Hasselbalch relates pH to the ratio of conjugate base to weak acid.
Selected acid–base systems and pKa values
- Acetic acid: pKa = 4.76 (Ka = $1.74 imes 10^{-5}$ ).
- Carbonic acid (first dissociation): pKa1 ≈ 3.77 (Ka1 ≈ $1.70 imes 10^{-3}$ ).
- Bicarbonate (second dissociation): pKa2 ≈ 10.2 (Ka2 ≈ $6.31 imes 10^{-11}$ ).
- Glycine (carboxyl group): pKa ≈ 2.34; amino group: pKa ≈ 9.60.
- Phosphoric acid (triprotic): pKa1 ≈ 2.14; pKa2 ≈ 6.86; pKa3 ≈ 12.4.
pKa and buffer design
- A buffer’s effective range roughly spans around its pKa (±1 pH unit). The Henderson–Hasselbalch equation governs how [A−] and [HA] determine the pH.

The buffering systems in biological contexts

The acetate buffer system
- CH3COOH ⇌ CH3COO− + H+. Buffering described by Ka and Henderson–Hasselbalch.
The phosphate buffer system (cytoplasm)
- H2PO4− ⇌ HPO4^{2−} + H+. Effective near pH 6.86, across approx. pH 5.9–7.9.
The bicarbonate buffer system (blood plasma)
- CO2(aq) + H2O ⇌ H2CO3 ⇌ H+ + HCO3−; buffer effectiveness near pH 7.4.
- The pH depends on the concentrations of HCO3− and CO2 (pCO2).
Additional equilibria related to CO2 hydration
- CO2(aq) ⇌ CO2(g) (gas–solution equilibrium) with Ka describing the dissolution.
- The overall bicarbonate system is characterized by: $K{ ext{combined}} = Kh Ka = \frac{[ ext{H}^+][ ext{HCO}3^-]}{[ ext{CO}_2( ext{aq})]}.$
Practical calculation of blood pH
- Using typical values: [H2CO3] ≈ 1.2 mM; [HCO3−] ≈ 24 mM.
- pH ≈ 7.4 using the Henderson–Hasselbalch form for the bicarbonate buffer.

Additional notes: Key concepts, constants, and implications

Ionization and pH in biological systems
- Ionization states regulate structure and function of biomolecules; many critical processes depend on local pH.
Dielectric properties and solubility
- Water’s high dielectric constant (ε ≈ 78.5 at 25 °C) stabilizes ionic solutes and reduces electrostatic attraction between charged species.
Osmosis and osmotic pressure
- Osmotic pressure Π for solutions is approximated by the van ’t Hoff equation: $\Pi = i c R T,$ where i is the van’t Hoff factor (extent of dissociation), c is molar concentration, R is gas constant, and T is absolute temperature.
- Osmolarity is defined as i c; osmosis is water movement driven by differences in osmotic pressure across a semipermeable membrane.
Hydrophobic effect and water structure around nonpolar solutes
- Nonpolar solutes disrupt water's hydrogen-bond network; water forms an ordered “cage” around nonpolar solutes, increasing enthalpy and decreasing entropy, making dissolution unfavorable without compensating hydrophobic interactions.
Hydration and proton transport in membranes
- Water molecules can be integral to proton relay pathways in proteins (e.g., cytochrome f) by forming hydrogen-bond networks.
Thermodynamic framing of binding and catalysis
- Enzyme–substrate binding often involves release of ordered water, contributing favorably to free energy changes through the hydrophobic effect and hydrogen bonding/ionic interactions.
van der Waals forces and radii
- van der Waals interactions are weak and distance-dependent; radii (van der Waals and covalent) define how close atoms can approach.
Practical implications for health and disease
- Acidosis and alkalosis refer to pH deviations from normal physiological ranges; serious conditions can arise if blood pH leaves the narrow viable window (roughly 7.3–7.5).
- Untreated diabetes can lead to metabolic acidosis due to accumulation of organic acids, lowering blood pH and impairing enzyme function.
Quick reference values
- Neutral pH at 25 °C: $ext{pH} = 7.0,$ $[ ext{H}^+] = [ ext{OH}^-] = 1.0 imes 10^{-7} ext{ M}.$
- Water ion product: $K_w = [ ext{H}^+][ ext{OH}^-] = 1.0 imes 10^{-14} ext{ M}^2.$
- Dielectric constant of water: oldsymbol{\varepsilon \approx 78.5} ext{ at } 25^ ext{°C}.
- Hydrogen-bond energy: ~ $23 ext{ kJ/mol}$ ; hydrogen bonds are short-lived (1–20 ps).
- Hydrogen bond coordination in liquid water: ~3.4 H-bonds per molecule; in ice: 4 H-bonds per molecule.
- Heat of fusion: $riangle H_ ext{fusion} \approx 5.9\ ext{kJ/mol};$ Heat of vaporization: $riangle H_ ext{vap} \approx 44.0\ ext{kJ/mol}.$
- Water’s role in buffering near pH 7.4
- Blood bicarbonate system buffers around physiologic pH; pK combined ≈ 6.1.

Quick reference formulas

Acid–base equilibria and constants
- Ka = $\frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}.$
- pH = $-\log [\text{H}^+].$
- Henderson–Hasselbalch: $\text{pH} = \text{p}K_a + \log \frac{[A^-]}{[HA]}.$
Water ionization and Kw
- $K_w = [\text{H}^+][\text{OH}^-] = 1.0\times 10^{-14}\text{ M}^2.$
- $\text{pH} + \text{pOH} = 14.$
Osmotic pressure
- $\Pi = i c R T.$ (van ’t Hoff form)
Bicarbonate system (three equilibria and combined constant)
- CO2(g) ⇌ CO2(aq) with Ka (dissolution equilibrium)
- CO2(aq) + H2O ⇌ H2CO3 with Kh
- H2CO3 ⇌ H+ + HCO3− with Ka1; HCO3− ⇌ H+ + CO3^{2-} with Ka2
- Combined: $K{ ext{combined}} = Kh Ka = \frac{[\text{H}^+][\text{HCO}3^-]}{[\text{CO}_2( ext{aq})]}.$
- Blood pH example (typical): pH ≈ 7.4 when [H2CO3] ≈ 1.2 mM and [HCO3−] ≈ 24 mM.
Buffers and pH relationships
- Buffer capacity is greatest near pKa; pH ∼ pKa when [HA] ≈ [A−].
- In buffering regions, small additions of acid/base cause minimal pH changes due to conjugate base/acid pair formation.
Nonpolar solubility and hydrophobic effect (conceptual)
- Nonpolar solutes force water into more ordered solvent shells, increasing enthalpy and decreasing entropy; hydrophobic effect drives nonpolar solutes to aggregate, reducing structured water exposure (e.g., micelles).
Hydrogen bonds as directional interactions
- Directionality is strongest when the acceptor is in line with the covalent donor–H bond, maximizing electrostatic interactions.
Miscellaneous biomolecule examples (polar/nonpolar/amphipathic)
- Polar: Glucose; Glycine; Lactate; Glycerol.
- Nonpolar: Typical waxes; long-chain hydrocarbons.
- Amphipathic: Phosphatidylcholine (phospholipid head is polar; hydrophobic tails).
Practical clinical relevance
- Acidosis: blood pH below ~7.35; Alkalosis: pH above ~7.45; extreme deviations are life-threatening.
- Untreated diabetes can cause metabolic acidosis due to accumulation of organic acids (e.g., β-hydroxybutyric acid and acetoacetate), lowering blood pH and impairing enzyme function.