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

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. The greater the difference in electronegativity between two atoms, the more polar the bond.

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

• Fluorine is the most electronegative element.

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, resulting in partial charges (δ+ and δ−). The dipole moment quantifies the separation of charge in a molecule and is a vector quantity.

• Water is polar due to the difference in electronegativity between oxygen and hydrogen and its bent molecular geometry.

• The dipole moment ($\mu$) of a molecule is calculated by vector addition of individual bond dipoles:

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

• Molecules like CO2 have polar bonds but are nonpolar overall due to their linear symmetry, resulting in a net dipole moment of zero.

Non-Covalent Interactions

Types of Non-Covalent Interactions

Non-covalent interactions are essential for stabilizing the structure of biomolecules such as proteins and nucleic acids. They are generally weaker than covalent bonds but are crucial for biological function.

• Ionic Bonds: Electrostatic attraction between oppositely charged ions.

• Salt Bridges: Specific ionic interactions within or between biomolecules.

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

• Van der Waals Forces: Weak attractions due to transient dipoles, including dipole-dipole, dipole-induced dipole, and induced dipole-induced dipole interactions.

Relative Strengths of Bonds in Biochemistry:

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

Hydrogen Bonding

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

• Hydrogen bonds are critical for the structure of DNA, proteins, and other biomolecules.

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., oils, fats).

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

Hydrophilic	Hydrophobic
Polar covalent compounds (alcohols, ketones)	Nonpolar covalent compounds (hydrocarbons, fats)
Amino acids, phosphates, carbohydrates	Fatty acids, cholesterol

The Hydrophobic Effect

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

• Critical for the formation of biological membranes and protein folding.

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 quantified by its dissociation constant ($K_a$):

$HA \rightleftharpoons H^+ + A^-$

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

Ion Product Constant 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 Scale

pH is a measure of hydrogen ion concentration:

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

Henderson-Hasselbalch Equation

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

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

• 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, each corresponding to a different $pK_a$.

Buffer Solutions

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 $pK_a$.

• Buffer capacity increases with higher concentrations of buffer components.

If pH equals	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 Buffers

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

• Bicarbonate buffer (H2CO3/HCO3-): Important in blood plasma.

Example of the bicarbonate buffer system:

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

• 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

• No interference with biological reactions or detection methods

• Suitable ionic strength and solubility

• Biological compatibility

Summary: Understanding the properties of water, non-covalent interactions, and acid-base chemistry is fundamental in biochemistry, as these principles govern the structure, stability, and function of biomolecules in aqueous environments.