Ch. 2 - Water of Life

Water makes up ___% of most organisms

70% or more

Hydrogen bonds in liquid water have a bond dissociation energy of

~23 kJ/mol (vs 470 kJ/mol for covalent O-H bonds)

Lifetime of a single hydrogen bond in water

1–20 picoseconds

Average number of hydrogen bonds per water molecule in liquid vs ice

3.4 (liquid); 4 (ice)

H-O-H bond angle and why it's not 109.5°

104.5°; nonbonding orbitals on oxygen crowd the hydrogen atoms

Why ice is less dense than liquid water

Ice forms a full tetrahedral lattice with 4 H-bonds per molecule, creating a more open, ordered structure

"Flickering clusters" refers to

Short-lived groups of water molecules interlinked by hydrogen bonds in liquid water

Hydrogen bonds are strongest when

The donor atom, hydrogen, and acceptor atom are in a straight line

Hydrophilic

Charged or polar compounds that dissolve readily in water

Hydrophobic

Nonpolar molecules (e.g., lipids, waxes) poorly soluble in water

Amphipathic

Molecules with both polar/charged and nonpolar regions

Dielectric constant of water at 25°C

78.5 (high; effectively screens ionic charges)

Why nonpolar solutes are thermodynamically unfavorable in water

ΔH is slightly positive and ΔS is negative (ordered cage of water molecules forms), making ΔG positive

Hydrophobic effect

Clustering of nonpolar regions to minimize ordered water shell, driven by entropy increase

Micelles

Stable structures formed by amphipathic compounds in water; nonpolar cores sequestered away from water

van der Waals interactions arise from

Transient induced dipoles between nearby uncharged atoms

van der Waals contact

Distance at which net attraction between two nuclei is maximal

Why cumulative weak interactions confer high molecular stability

All interactions must be simultaneously disrupted to dissociate molecules; random simultaneous disruption is very unlikely

Ion product of water (Kw) at 25°C

1.0 × 10⁻¹⁴ M²

Kw equation

Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ M²

pH definition

pH = −log[H⁺]

At neutral pH (25°C), [H⁺] and [OH⁻] each equal

1 × 10⁻⁷ M

pH < 7

Acidic; pH > 7

pH scale is ___ not ___

Logarithmic; arithmetic (1 pH unit = 10-fold difference in [H⁺])

Acid dissociation constant (Ka) expression

Ka = [H⁺][A⁻] / [HA]

pKa definition

pKa = −log Ka; lower pKa = stronger acid

At the midpoint of titration

[HA] = [A⁻]; pH = pKa

Henderson-Hasselbalch equation

pH = pKa + log([A⁻]/[HA])

Buffering region of a weak acid

±1 pH unit around its pKa (approximately 10–90% titration)

Maximum buffering power occurs when

[proton donor] = [proton acceptor]; pH = pKa

Phosphate buffer system pKa and effective range

pKa = 6.86; effective ~pH 5.9–7.9

Bicarbonate buffer system components

H₂CO₃ (proton donor) and HCO₃⁻ (proton acceptor); pKcombined = 6.1

Why bicarbonate is effective at blood pH 7.4 despite pKa of 6.1

Large CO₂ reservoir in lungs continuously replenishes H₂CO₃; breathing rate adjusts equilibrium rapidly

Normal blood plasma pH range

7.35–7.45

Acidosis

Blood pH < 7.35; causes headache, nausea, stupor, coma, convulsions

Alkalosis

Blood pH > 7.45; causes dizziness, headache, weakness, fainting

Optimum pH

The pH at which an enzyme shows maximal catalytic activity

Untreated diabetes mellitus causes acidosis because

Fatty acid metabolism produces β-hydroxybutyric acid and acetoacetic acid, lowering blood pH

Hyperventilation causes ___ by ___

Alkalosis; excessive CO₂ exhaled raises blood pH above 7.45

Treatment for severe acidosis

Intravenous bicarbonate solution to raise [HCO₃⁻] and shift pH upward

Histidine side chain pKa and buffering significance

pKa = 6.0; buffers effectively near neutral pH in cytoplasm

Why macromolecules have less effect on osmolarity than equal mass of monomers

Osmolarity depends on particle number, not mass; one polymer = one particle regardless of size

Van't Hoff equation for osmotic pressure

Π = icRT (i = van't Hoff factor, c = molar concentration)

Isotonic / hypertonic / hypotonic solutions

no net water movement; cell shrinks; cell swells