Water, Non-Covalent Interactions, and Acid-Base Chemistry in Biochemistry Ch 2

Electronegativity, Polarity, Dipole Moment, and Polar Compounds

Electronegativity

Electronegativity is a measure of an atom’s ability to attract shared electrons in a chemical bond. It is a key factor in determining the polarity of molecules.

• Oxygen and nitrogen are more electronegative than carbon and hydrogen.

• Fluorine is the most electronegative element (value = 4.0).

Element	Electronegativity
Oxygen	3.5
Nitrogen	3.0
Sulfur	2.5
Carbon	2.5
Phosphorus	2.1
Hydrogen	2.1

Polarity and Dipole Moment

A polar bond forms when electrons are unequally shared between atoms with different electronegativities, resulting in partial charges (δ+ and δ−).

• Dipole moment (μ) quantifies the separation of charge in a molecule and is a vector quantity.

• For molecules with multiple bonds, the total dipole moment is the vector sum of individual bond dipoles:

$\mu = \sqrt{\mu_1^2 + \mu_2^2 + 2\mu_1\mu_2 \cos \theta}$

• Symmetry affects molecular polarity: e.g., CO2 has polar bonds but is a nonpolar molecule due to its linear geometry.

Example: Water (H2O) is polar due to its bent shape and difference in electronegativity between O and H.

Non-Covalent Interactions in Biomolecules

Types of Non-Covalent Interactions

Non-covalent interactions are essential for stabilizing the structure of biomolecules such as proteins and nucleic acids.

• Ionic Bonds (Salt Bridges): Attraction between oppositely charged ions or groups.

• Hydrogen Bonds: Attraction between a hydrogen atom bonded to an electronegative atom (usually O or N) and another electronegative atom with a lone pair.

• Van der Waals Forces: Weak interactions including dipole-dipole, dipole-induced dipole, and dispersion (London) forces.

Example: The tertiary structure of proteins is stabilized by hydrogen bonds, ionic interactions, and van der Waals forces.

Hydrogen Bonding

• Hydrogen bonds are non-covalent and directional.

• Water can form up to four hydrogen bonds (two as donor, two as acceptor) due to its tetrahedral geometry.

• Hydrogen bonds are crucial in DNA base pairing and protein secondary structure.

Relative Strengths of Bonds

Bond Type	Strength (kcal/mol)
Covalent (C–H)	~105
Covalent (O–H)	~110
Hydrogen Bond	1–20
Van der Waals	~1

Hydrophilic, Hydrophobic, and Amphipathic Compounds

Definitions

• Hydrophilic: Water-loving; substances that dissolve easily in water (e.g., salts, sugars).

• Hydrophobic: Water-fearing; substances that do not dissolve in water (e.g., hydrocarbons, oils).

• Amphipathic: Molecules with both hydrophilic and hydrophobic regions (e.g., phospholipids, soaps).

The Hydrophobic Effect

The hydrophobic effect is the tendency of nonpolar substances to aggregate in aqueous solution, minimizing their exposure to water. This effect is driven by entropy: water molecules form ordered cages (clathrates) around hydrophobic molecules, decreasing entropy. Aggregation releases some water molecules, increasing entropy and stabilizing the system.

• Critical for membrane formation and protein folding.

Example: Formation of micelles by amphipathic molecules in water.

Acid/Base Dissociation, Ion Product Constant, pH, and the Henderson-Hasselbalch Equation

Acids and Bases

• Acid: Proton donor.

• Base: Proton acceptor.

The strength of an acid is given by its dissociation constant ($K_a$):

$HA \rightleftharpoons H^+ + A^-$ $K_a = \frac{[H^+][A^-]}{[HA]}$

Ion Product of Water

Water self-ionizes to a small extent:

$2H_2O \rightleftharpoons H_3O^+ + OH^-$ $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ (at 25°C)

pH is defined as:

$pH = -\log_{10}[H^+]$

Henderson-Hasselbalch Equation

This equation relates pH, pKa, and the ratio of conjugate base to acid:

$pH = pK_a + \log_{10}\left(\frac{[A^-]}{[HA]}\right)$

• When $[A^-] = [HA]$, $pH = pK_a$.

• If $pH < pK_a$, the acid form predominates (protonated).

• If $pH > pK_a$, the base form predominates (deprotonated).

Titration of Weak Acids/Bases and Buffer Solutions

Titration Curves

• Equivalence point: The point at which the amount of acid equals the amount of base during titration.

• Inflection point: The point on the titration curve where pH changes most rapidly; corresponds to $pH = pK_a$ for monoprotic acids.

• Polyprotic acids have multiple inflection points.

Buffers

A buffer is a solution that resists changes in pH upon addition of small amounts of acid or base. Buffers consist of a weak acid and its conjugate base.

• Effective buffering occurs within ±1 pH unit of the acid’s pKa.

• Buffer capacity increases with higher concentrations of buffer components.

pH relative to pKa	Base/Acid Ratio
pKa - 3	1/1000
pKa - 2	1/100
pKa - 1	1/10
pKa	1/1
pKa + 1	10/1
pKa + 2	100/1
pKa + 3	1000/1

Biological Buffer Systems

• Phosphate buffer (H2PO4-/HPO42-): Principal buffer in cells.

• Bicarbonate buffer (CO2/HCO3-): Important in blood plasma.

Example of bicarbonate buffering:

$CO_2(g) + H_2O(l) \rightleftharpoons H_2CO_3(aq) \rightleftharpoons H^+(aq) + HCO_3^-(aq)$

• Hyperventilation decreases blood CO2 and increases pH (alkalosis).

• Hypoventilation increases blood CO2 and decreases pH (acidosis).

Criteria for Good Buffers in Biochemistry

• Appropriate pKa value near the desired pH.

• No interference with biological reactions or detection methods.

• Suitable ionic strength and solubility.

• Biological or non-biological origin as required.

Summary Table: Properties of Water, Ammonia, and Methane

Property	Water (H2O)	Ammonia (NH3)	Methane (CH4)
Polarity	Polar	Polar	Nonpolar
Hydrogen Bonding	Yes	Yes	No
Boiling Point (°C)	100	-33	-161

Additional info:

• Some content was inferred and expanded for clarity and completeness, such as the explanation of the hydrophobic effect and buffer systems.

• Tables were reconstructed based on standard biochemistry knowledge.