Difficulty: Introductory | Prerequisites: Basic familiarity with atoms and elements from secondary school chemistry is helpful but not required.
This section covers the chemical foundations that make life possible. You need to understand chemical bonds because every biological molecule, from water to DNA, depends on them. Water's special properties (high specific heat, cohesion, solvent behaviour) create the environment in which all biochemistry happens. The pH scale and buffers explain how organisms maintain the narrow chemical conditions their enzymes require. This material underpins everything that follows in BIO 311C.
Covalent bonds (shared electrons) are the strongest bonds in biological systems. In water, ionic bonds are weaker than covalent bonds because water molecules pull the ions apart. Water's hydrogen bonds give it high specific heat, high heat of vaporisation, cohesion, and the ability to float as ice. The pH scale runs from 0 (strongly acidic) to 14 (strongly basic), with 7 as neutral. Buffers keep pH stable by absorbing or releasing H+ ions.
Covalent bond
A bond formed by the sharing of electron pairs between atoms. The strongest type of bond in biological (aqueous) systems. Think of two atoms holding hands: they share.
Ionic bond
A bond formed by the transfer of electrons from one atom to another, creating oppositely charged ions that attract. In water, ionic bonds are weakened because water molecules surround and pull apart the ions.
Hydrogen bond
A weak attraction between a slightly positive hydrogen atom (bonded to an electronegative atom like O or N) and a nearby electronegative atom. Individually weak, but collectively powerful: they give water its special properties.
Van der Waals interactions
The weakest intermolecular forces, arising from temporary fluctuations in electron distribution. In simple terms, these are fleeting attractions between molecules that happen to have their electrons unevenly distributed for a moment.
Electronegativity
The tendency of an atom to attract shared electrons toward itself. Oxygen is more electronegative than hydrogen, which is why water is polar.
Polar covalent bond
A covalent bond in which electrons are shared unequally because one atom is more electronegative. The O-H bonds in water are polar covalent: oxygen pulls the shared electrons closer, gaining a partial negative charge.
Cohesion
The attraction between molecules of the same substance. Water sticks to water via hydrogen bonds. This is what produces surface tension.
Adhesion
The attraction between molecules of different substances. Water sticking to the wall of a glass tube is adhesion.
Surface tension
The tendency of the surface of a liquid to resist disruption, caused by cohesion among surface molecules. This is why some insects can walk on water.
Specific heat
The amount of heat energy required to raise the temperature of 1 gram of a substance by 1 degree Celsius. Water's specific heat is high because hydrogen bonds must absorb energy before the molecules speed up.
Heat of vaporisation
The energy required to convert 1 gram of a liquid to gas. Water's is high because many hydrogen bonds must break for a molecule to escape into the air. This enables evaporative cooling (sweating).
pH
A scale from 0 to 14 measuring the concentration of hydrogen ions (H+) in a solution. pH 7 is neutral, below 7 is acidic, above 7 is basic (alkaline).
Buffer
A substance that minimises changes in pH by accepting H+ ions when the solution is too acidic and donating H+ ions when it is too basic. Think of it as a chemical shock absorber for pH.
Hydrophobic
Repelled by water. Nonpolar molecules are hydrophobic. In simple terms: oil and water do not mix because oil is hydrophobic.
Hydrophilic
Attracted to water. Polar and charged molecules are hydrophilic.
In a biological (aqueous) environment, the ranking of bond/interaction strength from strongest to weakest is:
Covalent bonds (shared electrons, not disrupted by water)
Ionic bonds (electrostatic attraction between ions; weakened in water because water molecules surround and separate the ions)
Hydrogen bonds (weak individually, strong collectively)
Van der Waals interactions (weakest, temporary dipole attractions)
Carbon is the backbone element of organic chemistry. It has four electrons in its outer shell, giving it a valence of 4 and the ability to form four covalent bonds. This makes carbon extraordinarily versatile: it can form long chains, branched structures, and rings.
In water, oxygen is more electronegative than hydrogen. The shared electrons spend more time near the oxygen, giving it a partial negative charge and the hydrogens a partial positive charge. The result is a polar covalent bond and a polar molecule.
Water's polar nature and hydrogen bonding give it several properties critical to life.
Cohesion and surface tension. Water molecules stick to each other (cohesion) through hydrogen bonds. At the surface, this creates surface tension, a film-like effect that allows small insects to walk on a pond.
Adhesion. Water sticks to other polar surfaces (e.g. the walls of a xylem vessel in a plant), helping pull water upward against gravity.
High specific heat. A large amount of heat energy is needed to raise water's temperature because some of that energy goes toward breaking hydrogen bonds rather than speeding up molecules. This stabilises temperatures in oceans and inside organisms.
High heat of vaporisation. Many hydrogen bonds must break for a water molecule to escape into the air as gas. When the most energetic (hottest) molecules evaporate from a surface, the surface cools. This is evaporative cooling, and it is why sweating works.
Ice floats. When water freezes, hydrogen bonds lock molecules into a crystalline lattice that spaces them farther apart than in liquid water. Ice is therefore less dense than liquid water and floats. Floating ice insulates the water below, preventing lakes and oceans from freezing solid and protecting aquatic life.
Versatile solvent. Water dissolves polar and ionic substances (hydrophilic compounds) by surrounding their molecules or ions. Nonpolar substances (hydrophobic) do not dissolve in water.
The pH scale measures hydrogen ion concentration on a logarithmic scale from 0 to 14. Each whole-number step represents a tenfold change in H+ concentration.
pH below 7: acidic (higher H+ concentration)
pH 7: neutral
pH above 7: basic/alkaline (lower H+ concentration, higher OH- concentration)
A solution with a pH of 12 is strongly basic. A solution with a pH of 4 is acidic.
Buffers are substances (typically a weak acid and its conjugate base) that resist changes in pH. When acid is added, the buffer's base component accepts the extra H+ ions. When base is added, the buffer's acid component donates H+ ions. In human blood, the bicarbonate buffer system keeps pH near 7.4, which is critical for enzyme function.
pH = -log[H+]
[H+] x [OH-] = 10^-14 (the ion product constant of water at 25 C)
Worked example: if [H+] = 10^-9 M, then pH = 9. To find [OH-]: 10^-9 x [OH-] = 10^-14, so [OH-] = 10^-5 M.
Worked example: if pH = 4, then [H+] = 10^-4 M. [OH-] = 10^-14 / 10^-4 = 10^-10 M.
Students often think ionic bonds are stronger than covalent bonds. In a vacuum or a crystal, ionic bonds can be very strong. In an aqueous (biological) environment, water weakens ionic bonds significantly, making covalent bonds the strongest.
Students sometimes confuse cohesion (water-to-water attraction) with adhesion (water-to-other-surface attraction). Both involve hydrogen bonding, but cohesion is between like molecules and adhesion is between different molecules.
A common error is thinking a low pH number means fewer H+ ions. The opposite is true: pH 2 has far more H+ ions than pH 7. The scale is logarithmic and inverse.
Students forget that each pH unit represents a tenfold change. A solution at pH 3 has ten times more H+ than a solution at pH 4, and a hundred times more than pH 5.
⚠️ The bond-strength ranking question (covalent > ionic > hydrogen > van der Waals in aqueous environments) appears in both multiple-choice and short-answer format.
⚠️ Expect a calculation question: given [H+], find pH and [OH-], or vice versa. Know the formulas cold.
⚠️ Surface tension (cohesion) is tested via a real-world example, often insects walking on water.
⚠️ The ice-floating question asks you to explain why (hydrogen bonds create a less-dense lattice) and why it matters biologically (insulates water below, protects aquatic life).
⚠️ Buffers are tested both as true/false (definition) and as a long-answer question (explain what happens to H+ when acid is added to a buffered solution).
True or false: In an aqueous environment, ionic bonds are stronger than covalent bonds. (False.)
Fill in the blank: Water molecules stick to other water molecules due to ______. (Cohesion.)
True or false: A buffer minimises changes in pH by accepting or donating H+ ions. (True.)
Fill in the blank: A solution with a pH of 12 is considered strongly ______. (Basic / alkaline.)
If pH = 4, what is [OH-]? (10^-10 M.)
Q: Which property of water is most directly responsible for the ability of some insects to walk on a pond?
A: Cohesion and surface tension. Hydrogen bonds between water molecules create a cohesive surface that resists disruption.
Q: Explain why ice floats on liquid water and how this supports life.
A: When water freezes, hydrogen bonds lock molecules into a crystalline lattice that is less dense than liquid water. The floating ice insulates the water below, preventing the entire body of water from freezing solid and allowing aquatic organisms to survive.
Q: Rank the following from strongest to weakest in an aqueous biological environment: hydrogen bonds, Van der Waals interactions, covalent bonds, ionic bonds.
A: Covalent bonds > ionic bonds > hydrogen bonds > Van der Waals interactions.
Q: Calculate the pH of a solution where [H+] = 10^-9 M. Determine [OH-]. If a buffer is added and then a small amount of acid is introduced, what happens to the H+ ions?
A: (a) pH = 9. (b) [OH-] = 10^-5 M. (c) The buffer's base component accepts the added H+ ions, preventing the pH from dropping significantly.
Q: Briefly explain the role of a buffer in a biological system like human blood.
A: Buffers resist changes in pH by accepting excess H+ ions when the solution becomes too acidic and donating H+ ions when it becomes too basic. In blood, the bicarbonate buffer system maintains pH near 7.4, which is critical for enzyme function and homeostasis.
Water properties connect directly to the macromolecules unit: proteins fold in water, and hydrophobic interactions in an aqueous environment drive much of tertiary protein structure. pH and buffers connect to enzyme function (enzymes denature outside their optimal pH range) and to the functional groups on amino acids, which can gain or lose protons depending on pH.
covalent bond, ionic bond, hydrogen bond, van der Waals, electronegativity, polar covalent, nonpolar covalent, water properties, cohesion, adhesion, surface tension, specific heat, heat of vaporisation, evaporative cooling, ice density, universal solvent, hydrophobic, hydrophilic, pH scale, acid, base, alkaline, buffer, bicarbonate, ion product constant, carbon valence, BIO 311C, intro to biology, UT Austin