Chemistry of Life: Bonds, Polarity, and Water in Biology

Covalent bonds and bond polarity

Carbon can form up to four bonds: it can share electrons to make up to four covalent bonds.
The simplest covalent compound is methane, where carbon bonds with hydrogens to form CH₄.
Covalent bonds are not ionic; they have discrete bond angles and specific shapes.
When carbon is bound to four hydrogens, it has a particular shape (e.g., methane).
Linking many molecules (e.g., CH₄ units) in specific orientations can create new shapes and new functions; for example, a macromolecule like cholesterol has new shape and new function compared to a single methane molecule.
Covalent bonds differ from ionic bonds in how electrons are distributed around the two atoms involved.
In a polar covalent bond, electrons are shared unequally, and one atom may hog electrons, creating partial charges: partial negative on the more electronegative atom and partial positive on the less electronegative atom.
These partial charges are denoted as δ⁻ and δ⁺.
The shape and properties of a molecule depend on how electrons are distributed in its bonds.
A polar covalent bond involves unequal sharing of electrons, leading to polar regions within a molecule that can interact differently with other molecules.

Electronegativity and bond polarity

Electronegativity describes an atom’s affinity for electrons in a bond.
There is a trend: electronegativity increases as you move to the right on the periodic table and generally decreases down a group; the course notes describe an increase “diagonally” as you go from left to right.
Fluorine has the highest electronegativity among the elements, but it is not common in biology.
Oxygen is the most electronegative biologically relevant element discussed here.
The purpose of memorizing electronegativity values for key atoms is to determine bond polarity and, thus, bond behavior.
Example numbers used in class: Hydrogen ≈ $2.1$, Oxygen ≈ $3.5$.
The difference in electronegativity drives bond polarity:
- Let \Delta EN = ENA - ENB.
- If 0 \le \Delta EN \le 0.4\, → nonpolar covalent bond.
- If 0.5 \le \Delta EN \le 1.7\, → polar covalent bond.
- If \Delta EN > 1.7\, → ionic bond.
For a hydrogen-oxygen bond,
\Delta EN = EN(O) - EN(H) = 3.5 - 2.1 = 1.4,
which falls in the polar covalent range.
Why memorize? The differences in electronegativity determine bond polarity and thus the chemistry and behavior of molecules in biological contexts.

Polar vs nonpolar molecules; hydrophilic vs hydrophobic

Polar covalent bonds create regions of partial charge that can interact with other polar regions or with water.
Polar molecules tend to associate with water (hydrophilic).
Nonpolar molecules tend to associate with other nonpolar molecules and tend to exclude water (hydrophobic).
Many biomolecules have both polar and nonpolar regions, giving them multiple properties and interactions.
The orientation and type of bonding in a molecule influence its solubility and interactions in aqueous environments.

Hydrogen bonds

Hydrogen bonds are electrostatic interactions that occur when a hydrogen atom, covalently bonded to a partially negative atom (e.g., O or N), experiences attraction from another partially negative atom’s lone pair.
Example: an O–H bond (polar covalent) creates a partial positive on H and partial negative on O; this H can form a hydrogen bond with another electronegative atom (e.g., another O or N with lone pairs).
In biological systems, hydrogen bonds can occur within a molecule or between molecules.
These bonds are individually weak but collectively can be very strong; they are essential for the stability and properties of many biomolecules.
Hydrogen bonds are a key factor in water’s behavior and in the structure of macromolecules.

Van der Waals forces

Van der Waals forces involve attractions between nonpolar molecules when they come very close to one another.
They arise from transient dipoles and interactions between nuclei and electron clouds.
These forces are short-range and relatively weak on their own, but when many such interactions occur together, they can provide significant stability.
Van der Waals forces help explain processes like surface interactions and the ability of organisms to move over surfaces.

Water: structure and hydrogen bonding

Water is H₂O, with two polar covalent bonds between hydrogen and oxygen.
The O–H bonds are polar covalent because of electronegativity difference: O is more electronegative than H, giving oxygen a partial negative charge and hydrogen a partial positive charge.
Water molecules can hydrogen bond with each other due to these partial charges, forming a lattice in ice and a dynamic network in liquid water.
Water can hydrogen bond with itself (cohesion) and with other molecules (adhesion).
Water’s unusual properties stem from hydrogen bonding:
- High cohesion and surface tension.
- High specific heat capacity (many hydrogen bonds must be broken to increase temperature).
- High heat of vaporization (a lot of energy is required to break bonds so that molecules can escape as vapor).
Water has anomalous density behavior: liquid water is more dense than ice over a substantial temperature range.
This density anomaly is important for life on Earth: ice floats on liquid water, insulating the water below and helping sustain aquatic life.
Water is essential for metabolism: organisms rely on water for their chemical reactions; cells are roughly ~60% water; life depends on water to enable metabolic processes and the formation of biomolecules.
Water can act as a solvent and participate directly in chemical reactions; it enables the linking and interaction of biomolecule building blocks and macromolecules.
Water’s role in biology includes maintaining stable temperatures in organisms and environments due to its high heat capacity and heat of vaporization.
The lecture emphasizes that water’s properties (cohesion, adhesion, surface tension, density, specific heat, vaporization) arise from hydrogen bonding.

Implications and connections

Structure determines function: the orientation and types of bonds lead to different shapes and biological roles (e.g., methane vs. cholesterol macrostructure).
Bond polarity and electronegativity differences underpin many biological interactions, such as solubility, protein folding, membrane dynamics, and molecular recognition.
Water’s unique properties are foundational to biochemistry and life on Earth: metabolism, temperature regulation, and the stability of biological systems rely on water’s hydrogen-bond network.
The course links chemistry concepts (bonds, polarity, electronegativity) to real-world biology, including why certain molecules interact with water and how macromolecules achieve their functions.

Quick recap and assessment hints

Key bond types: covalent (including polar covalent) vs ionic; hydrogen bonds as a special case of intermolecular interaction.
Electronegativity differences determine bond polarity and bond classification:
- 0 \le \Delta EN \le 0.4: nonpolar covalent
- 0.5 \le \Delta EN \le 1.7: polar covalent
- \Delta EN > 1.7: ionic
For H–O: \Delta EN = 3.5 - 2.1 = 1.4 → polar covalent.
Water properties to remember: density anomaly (ice floats), high specific heat and high heat of vaporization, cohesion, adhesion, and surface tension.
Real-world connections: membrane dynamics, protein structure, and metabolism all hinge on these chemical principles.
Quiz reminder: there is a graded content quiz covering Wednesday’s and today’s lecture; open until Sunday at midnight.