Difficulty: Intermediate | Prerequisites: Chapter 1 (biological hierarchy, atoms and molecules)
This chapter moves from the broad view of life down to its chemical foundations. You need to understand how atoms bond, why water is so critical to biology, and how functional groups determine the behaviour of organic molecules. The concepts here, particularly bond types, polarity, and pH, underpin everything in Chapter 3 (biological molecules) and beyond. If electronegativity and the octet rule feel unfamiliar, review a general chemistry primer before diving in.
Atoms bond through covalent, ionic, and hydrogen bonds, and the type of bond depends on electronegativity differences between atoms. Water's unique properties (cohesion, high specific heat, solvent ability) stem from hydrogen bonding and are essential for life. Organic molecules get their chemical behaviour from functional groups such as hydroxyl, carboxyl, amino, and phosphate groups.
Proton
A positively charged subatomic particle found in the nucleus of an atom. The number of protons defines the element (its atomic number).
Neutron
A subatomic particle with no charge, found in the nucleus alongside protons. Neutrons contribute to mass but not to charge.
Electron
A negatively charged subatomic particle that orbits the nucleus in energy shells. Electrons in the outermost shell (valence electrons) participate in bonding.
Mass number
The sum of protons and neutrons in an atom's nucleus. Think of it as the atom's "headcount" of heavy particles.
Molecular weight (molecular mass)
The total mass of all atoms in a molecule, expressed in daltons or g/mol. In simple terms, add up the atomic weights of every atom in the molecular formula.
Valence electrons
Electrons in the outermost energy shell of an atom. These are the electrons that interact with other atoms to form bonds.
Valence (bonding capacity)
The number of covalent bonds an atom can form, determined by how many unpaired valence electrons it has out of 8 (octet rule). H = 1, O = 2, N = 3, C = 4.
Octet rule
Atoms tend to gain, lose, or share electrons until they have eight electrons in their outer shell, achieving a stable configuration. Think of it as every atom wanting a "full house" of eight.
Electronegativity
The relative ability of an atom to attract shared electrons when bonded to another atom. In simple terms, it measures how greedy an atom is for electrons. Increases across a period (left to right) and up a group on the periodic table.
Nonpolar covalent bond
A bond where two atoms share electrons equally because they have the same or very similar electronegativity. Example: O2, H2.
Polar covalent bond
A bond where electrons are shared unequally because one atom is more electronegative than the other, creating partial positive and partial negative charges. Example: the O-H bonds in water.
Ionic bond
A bond formed when one atom completely transfers electrons to another, typically between a metal and a nonmetal. The electron donor becomes a cation (+) and the acceptor becomes an anion (-).
Hydrogen bond
A weak attraction between a hydrogen atom covalently bonded to an electronegative atom (like O or N) and another electronegative atom nearby. Individually weak, but collectively very important (they hold water together and stabilise protein structure).
Van der Waals interactions
Weak, transient attractions between molecules caused by the constantly shifting distribution of electron clouds. Think of them as fleeting "static cling" between molecules.
Hydrophilic
Polar or charged molecules that dissolve readily in water. "Water-loving."
Hydrophobic
Nonpolar molecules that do not dissolve in water and tend to cluster together. "Water-fearing." Example: lipids.
pH
A measure of hydrogen ion concentration: pH = -log[H+]. Lower pH = more acidic; higher pH = more basic. Neutral water has a pH of 7.
Kw (ion product of water)
The equilibrium constant for water's auto-ionisation: Kw = [H+] x [OH-] = 10^-14 at 25 degrees C. This relationship always holds in aqueous solution.
pK
The negative logarithm of any equilibrium constant. Used to express dissociation constants on a convenient scale.
Acid
A substance that donates protons (H+), increasing [H+] in solution and lowering pH.
Base
A substance that accepts protons or donates hydroxide ions (OH-), decreasing [H+] and raising pH.
Buffer
A substance that resists changes in pH by absorbing excess H+ (when pH drops) or donating H+ (when pH rises). Think of it as a chemical shock absorber for pH.
Functional group
A specific group of atoms within a molecule that determines its chemical properties and reactivity. The same functional group behaves similarly regardless of the molecule it is attached to.
Every atom has three subatomic particles: protons (+, in the nucleus), neutrons (no charge, in the nucleus), and electrons (-, orbiting in shells). The atomic number equals the number of protons and defines the element. The mass number equals protons + neutrons.
Atomic weight is the weighted average of all isotopic masses of an element based on their natural abundance. Molecular weight (molecular mass) is the sum of the atomic weights of all atoms in a molecule.
Valence electrons sit in the outermost shell and are the ones involved in bonding. The octet rule states that atoms are most stable with eight electrons in their outer shell (hydrogen is stable with two).
Valence (bonding capacity) is determined by how many bonds an atom needs to reach eight:
H = 1, O = 2, N = 3, C = 4, P = 3 (or 5 in some contexts)
Na = 1 (tends to lose one electron), Cl = 1 (tends to gain one electron)
Atoms with three or fewer valence electrons tend to lose them entirely (forming cations). Atoms with five or more tend to gain electrons (forming anions).
Electronegativity increases from left to right across a period and from bottom to top within a group on the periodic table. It determines how bonds form:
Nonpolar covalent bond: two atoms with the same electronegativity share electrons equally. Example: H-H in H2.
Polar covalent bond: electrons shared unequally because one atom is more electronegative. Partial charges develop (delta+ and delta-). Example: O-H bonds in water.
Ionic bond: one atom is so much more electronegative that it pulls electrons away completely. One atom becomes a cation, the other an anion. Typically metal + nonmetal. Example: NaCl.
Hydrogen bond: a weak attraction between H (bonded to an electronegative atom like O or N) and a nearby electronegative atom. Individually weak, collectively powerful.
Van der Waals interactions: very weak, transient attractions from fluctuating electron clouds. They depend on molecular shape and proximity.
Hydrophobic interactions: nonpolar molecules cluster together in aqueous environments, not because they attract each other, but because water excludes them. Example: lipid tails of phospholipids.
Bond strength ranking for this course: covalent > ionic > hydrogen > Van der Waals.
Water's properties arise from its polar covalent O-H bonds and the hydrogen bonds between water molecules:
Cohesion: water molecules stick to each other via H-bonds. Important for water transport in plants (pulling water up through xylem).
Adhesion: water sticks to other polar surfaces. Helps with absorption and capillary action.
High specific heat: water absorbs and releases large amounts of heat with little temperature change. Stabilises temperatures in organisms and environments.
High heat of vaporisation: it takes a lot of energy to evaporate water. This allows evaporative cooling (sweating, transpiration in plants).
Ice floats: when water freezes, H-bonds lock molecules into a crystalline structure that is less dense than liquid water. Ice floats, insulating the water below and protecting aquatic life.
Versatile solvent: water's polarity lets it dissolve polar and ionic substances readily.
Medium for biochemical reactions: water is the environment in which most cellular reactions occur and is itself a reactant in many (e.g. hydrolysis).
To prepare a molar solution from a solid:
MW (g/mol) x M (mol/L) x volume (L) = grams of solute needed
Example: to make 100 mL of 0.5 M NaOH (MW = 40): 40 g/mol x 0.5 mol/L x 0.1 L = 2 g of NaOH, dissolved in water to a final volume of 100 mL.
To dilute a stock solution:
C1V1 = C2V2
C1 = concentration of stock, V1 = volume of stock to use, C2 = desired concentration, V2 = desired final volume.
Example: to make 100 mL of 2% NaCl from a 10% stock: 10% x V1 = 2% x 100 mL, so V1 = 20 mL of stock + 80 mL of water.
pH = -log[H+]. The scale runs from 0 (most acidic) to 14 (most basic), with 7 as neutral.
Key relationships:
[H+] x [OH-] = Kw = 10^-14 (always, at 25 degrees C)
As [H+] increases, pH decreases (more acidic)
As [OH-] increases, pH increases (more basic)
Each pH unit represents a tenfold change in [H+]
Acids donate H+ (lower pH). Bases accept H+ or donate OH- (raise pH). Buffers resist pH change by absorbing or releasing H+ as needed.
Hydroxyl (-OH): polar, classified as an alcohol. Found in carbohydrates, proteins, nucleic acids.
Carbonyl (-C=O): slightly polar. At the end of a carbon chain = aldehyde; in the middle = ketone. Found in sugars, proteins, nucleotides. Can react with -OH to form ring structures.
Carboxyl (-COOH): acidic (donates H+). Found in all amino acids and fatty acids.
Amino (-NH2): basic (accepts H+). Found in all amino acids and proteins.
Sulfhydryl (-SH): reactive. Two -SH groups can form a disulfide bridge (-S-S-), important for protein structure. Found in active sites of some enzymes.
Phosphate (-OPO3): highly polar, acidic, reactive. Found in ATP, DNA, RNA, phospholipids. Critical for energy transfer.
Methyl (-CH3): nonpolar, affects solubility and bioactivity. Found in fatty acids, some amino acids, nucleotides.
Making a molar solution from a solid:
MW (g/mol) x M (mol/L) x Volume (L) = grams of solute
Dilution formula:
C1 x V1 = C2 x V2
pH calculation:
pH = -log[H+]
[H+] = 10^(-pH)
Kw = [H+] x [OH-] = 10^-14
Dimensional analysis tip: write each conversion factor as a fraction, multiply the tops, divide by the bottoms, and cancel units. Example: (40 g / 1 mol) x (0.5 mol / 1 L) x (1 L / 1000 mL) x 100 mL = 2 g.
Students often think hydrogen bonds are a type of covalent bond. They are not. Hydrogen bonds are weak intermolecular attractions, far weaker than covalent or ionic bonds.
Students confuse polar molecules with ions. A polar molecule has unequal electron sharing but no complete electron transfer. An ion has gained or lost electrons entirely.
Students assume a low pH means a weak acid. Low pH means high [H+], which means strongly acidic. Strength and concentration are related but distinct concepts.
Students forget that each pH unit is a tenfold difference. A solution at pH 3 has ten times more H+ than one at pH 4, and a hundred times more than one at pH 5.
True or False: Carbon can form four covalent bonds.
True. Carbon has 4 valence electrons, so its valence is 4.
Fill in the blank: A bond where electrons are shared unequally is called a _____ covalent bond.
Polar.
True or False: A solution with a pH of 2 has a higher [H+] than one with a pH of 5.
True. pH 2 has 1000 times more H+ than pH 5.
Fill in the blank: The functional group -COOH is called the _____ group and acts as an _____.
Carboxyl; acid.
True or False: Ice sinks in liquid water because it is denser.
False. Ice floats because its crystalline H-bond structure makes it less dense than liquid water.
⚠️ Know the valence of C, N, O, H, P, Na, and Cl.
⚠️ Be able to distinguish between nonpolar covalent, polar covalent, ionic, and hydrogen bonds given a molecular scenario.
⚠️ Know all seven properties of water and connect each to hydrogen bonding.
⚠️ Be able to calculate grams of solute for a molar solution and use C1V1 = C2V2 for dilutions.
⚠️ Know the seven functional groups, their polarity, whether they are acidic/basic, and where they appear in biological molecules.
Q: Explain why water has a high specific heat and why this matters for living organisms.
A: Hydrogen bonds between water molecules absorb a large amount of heat before the temperature rises. This means water resists rapid temperature changes, which helps organisms maintain stable internal temperatures and buffers aquatic environments from extreme temperature swings.
Q: A molecule has an -NH2 group and a -COOH group. Predict whether each group will act as an acid or a base in solution.
A: The -COOH (carboxyl) group acts as an acid by donating H+. The -NH2 (amino) group acts as a base by accepting H+.
Q: What is the [H+] of a solution with pH 4? What is the [OH-]?
A: [H+] = 10^-4 M. Since Kw = [H+] x [OH-] = 10^-14, [OH-] = 10^-14 / 10^-4 = 10^-10 M.
Q: You need to make 250 mL of 0.2 M glucose (MW = 180). How many grams of glucose do you weigh out?
A: 180 g/mol x 0.2 mol/L x 0.25 L = 9 g of glucose.
Q: Compare hydrogen bonds and Van der Waals interactions in terms of strength and biological importance.
A: Hydrogen bonds are stronger than Van der Waals interactions and play major roles in water's properties, DNA base pairing, and protein structure. Van der Waals interactions are weaker and more transient, arising from temporary electron cloud fluctuations, but they contribute to molecular shape recognition and membrane structure when many occur together.
The bond types and functional groups from this chapter directly determine how the four classes of biological molecules (Chapter 3) behave. Peptide bonds in proteins are covalent. DNA's double helix is held together by hydrogen bonds. Phospholipid membranes depend on hydrophobic interactions. Understanding polarity here is essential for grasping membrane transport and enzyme function later in the course.
Atomic particles, proton, neutron, electron, mass number, molecular weight, valence electrons, octet rule, electronegativity, covalent bond, polar covalent, nonpolar covalent, ionic bond, hydrogen bond, Van der Waals, hydrophilic, hydrophobic, water properties, cohesion, adhesion, specific heat, heat of vaporisation, solvent, pH scale, acid, base, buffer, Kw, pK, molarity, dilution formula, C1V1=C2V2, functional groups, hydroxyl, carbonyl, carboxyl, amino, sulfhydryl, phosphate, methyl, BIO 311C, introductory biology, University of Texas