Introduction to Biochemistry

Biomolecules

• All contain carbon.
• Include proteins (amino acids), sugars (monosaccharides), nucleic acids (nucleotides), and lipids (fatty acids & glycerol).
• Assembling these requires energy.

Energy and Thermodynamics

• Thermodynamics describes energy.
• First Law: Energy cannot be created or destroyed, only transformed.
• Second Law: Entropy (disorder) of the universe increases.

Gibbs Free Energy (G)

• Components: Enthalpy (H) and Entropy (S)
• Enthalpy (H): Heat content of a system.
• Entropy (S): How energy is dispersed in a system.
• Formula: \Delta G = \Delta H - T\Delta S
• Exothermic reactions release heat ($\$\Delta H < 0\$).
• Endothermic reactions absorb heat ($\$\Delta H > 0\$).

Spontaneous Reactions

• Spontaneous (exergonic): $\$\Delta G < 0\$
• Nonspontaneous (endergonic): $\$\Delta G > 0\$
• At equilibrium: $\$\Delta G = 0\$
• Reactions with decreased enthalpy and increased entropy are spontaneous at all temperatures.
• Nonspontaneous reactions can proceed in vivo by coupling with spontaneous reactions.

Cell Types

• Prokaryotes: Small, unicellular, lack a discrete nucleus and internal membranes.
• Eukaryotes: Larger, uni/multicellular, have a discrete nucleus and internal membranes.

Eukaryotic Organelles

• Lysosomes: Degrade macromolecules.
• Peroxisomes: For oxidative reactions.
• Vacuoles: For storage.
• Nucleus: Has a double membrane.

Human Body

• Contains ~10 trillion cells and 4-100 trillion microorganisms (microbiome), mostly in the intestine.

Water

• Fundamental for life; ~60% of human body weight.
• Molecular geometry: bent; Electron geometry: tetrahedral.

Molecular Forces

• Intramolecular (e.g., covalent bonds) and intermolecular forces (e.g., hydrogen bonds).
• Complementarity of bases in DNA/RNA: A-T (2 H-bonds), G-C (3 H-bonds).

Water as a Solvent

• High dielectric constant (DC) diminishes electrostatic interactions between ions.
• $\$\uparrow\$DC = $\$\downarrow\$ ability of ions to associate with each other.

Amphiphilic Molecules

• Have polar and nonpolar components; may form spherical micelles or sheets.
• Examples: Palmitate.

Sweating and Sports Drinks

• Sweating causes evaporation, which has a cooling effect (2.5 kJ heat released per gram/mL of water lost).
• Exercise leads to water and electrolyte (Na+, Cl-, K+) losses.
• Short workouts (<90min): water is enough; Long/intense: sports drinks needed.

Cell Membrane Permeability

• Carbon dioxide, oxygen gas, and steroids can freely cross the cell membrane.

Acid-Base Chemistry

• H+ combines with water to form hydronium ion (H3O+).
• Proton jumping: Protons move and are shared among nearby water molecules.
• K_w = [H^+][OH^-] = 10^{-14}
• Neutral: [H+] = [OH-] = 10^{-7}\$ M
• Acidic: [H+] > [OH-]
• Basic: [H+] < [OH-]
• pH = -log [H+]
• Neutral: pH = 7; Acidic: pH < 7; Basic: pH > 7; Physiologic: pH = 7.4
• BAAD: Bases Accept H+, Acids Donate H+

Weak Acids/Bases

• Higher Ka = Lower pKa = Acid more likely to ionize (higher tendency to donate a proton).

Henderson-Hasselbalch Equation

• pH = pK_a + log \frac{[A^-]}{[HA]}
• When pH = pKa, the acid is half dissociated ([A-] = [HA]).

Buffers

• Resist changes in pH by:
• Combining added protons with the conjugate base.
• Donating protons from the acid to neutralize added hydroxide ions.
• Effective buffering capacity: generally within one pH unit of its pKa.

Bicarbonate Buffer System

• Principal buffer in interstitial fluid and blood plasma.
• CO2 + H2O \rightleftharpoons H2CO3 \rightleftharpoons H^+ + HCO_3^-$$

Le Chatelier's Principle

• Normal, excess, or insufficient acid conditions shift the equilibrium accordingly.

Clinical Correlation: Acid-Base Homeostasis

• Hyperventilation: ↓ CO2 in blood → ↓ carbonic acid → ↓ H+ ions, pH rises (respiratory alkalosis).
• Paper bag helps rebreathe CO2.
• Increased CO2 in blood: Shifts to right to produce more H+ (e.g., diabetic ketoacidosis and metabolic acidosis).