13.1 Cells, Organelles, Water, and Biological Acid-Base Buffers

13.1 Cells, Organelles, Water, and Biological Acid-Base Buffers

Quick Answer: Biochemistry area A (10 items) covers cell types and organelles, water, and acid-base chemistry with buffers. Know which organelle hosts each pathway, why water is such a good biological solvent, and how to use the Henderson-Hasselbalch equation for phosphate, bicarbonate, and laboratory buffers. That includes the blood-gas calculation that distinguishes respiratory from metabolic acidosis.

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Prokaryotic and Eukaryotic Cells

Ribosome differences matter in medicine: antibiotics such as streptomycin and tetracycline target 70S bacterial ribosomes. Mitochondria and chloroplasts have their own circular DNA and 70S-like ribosomes, which supports the endosymbiotic theory of their origin.

Organelles and Their Functions

Where Pathways Run

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Water: The Biological Solvent

• Polarity and hydrogen bonding: the bent molecule (104.5 degrees) has a dipole moment of 1.85 D. In ice each molecule hydrogen-bonds to four neighbours.
• Thermal properties: a high specific heat (4.18 J/g K) and heat of vaporization (about 2.26 kJ/g at 100 C) buffer organisms against temperature change, and evaporation cools them efficiently.
• Dielectric constant of about 78.5 at 25 C weakens attractions between ions, so salts dissolve.
• The hydrophobic effect: nonpolar molecules force surrounding water into ordered cages. Clustering nonpolar groups together releases that water and increases entropy. This drives protein folding, membrane formation, and micelle assembly.
• Amphipathic molecules (fatty acid salts, phospholipids) form micelles and bilayers with their nonpolar tails hidden from water.

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Ionization of Water and Weak Acids

$K_w = [H^+][OH^-] = 1.0 \times 10^{-14}\ (25\text{ C}), \qquad pH = -\log[H^+]$

$K_w$ rises with temperature: at 37 C it is about $2.4 \times 10^{-14}$, so neutral pH is about 6.8 at body temperature.

Henderson-Hasselbalch equation:

$pH = pK_a + \log\frac{[A^-]}{[HA]}$

A buffer resists pH change best at $pH = pK_a$ and is useful within about $\pm 1$ unit of it.

Worked buffer preparation: To make 1.00 L of 0.100 M phosphate buffer at pH 7.40 ($pK_{a2} = 7.20$), use $\frac{[HPO_4^{2-}]}{[H_2PO_4^-]} = 10^{0.20} = 1.58$. That gives $[H_2PO_4^-] = \frac{0.100}{2.58} = 0.0387$ M and $[HPO_4^{2-}] = 0.0613$ M.

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Physiological Buffers

The Bicarbonate Buffer and Blood Gases

Dissolved $CO_2$ equals $0.0301 \times p_{CO_2}$ (mmHg) in mmol/L, so

$pH = 6.1 + \log\frac{[HCO_3^-]}{0.0301 \times p_{CO_2}}$

Normal arterial values (24 mM $HCO_3^-$, 40 mmHg $CO_2$) give $pH = 6.1 + \log\frac{24}{1.20} = 6.1 + 1.30 = 7.40$. The normal range is 7.35-7.45.

Laboratory (Good's) Buffers

Good's buffers were designed to be water-soluble, membrane-impermeable, and weakly metal-binding, with $pK_a$ values near 6-8.