Unit 1: Chemistry of Life

Atoms, Elements, Isotopes, and the Chemical Bonds of Life

Biology is “chemistry with consequences.” Every biological structure you learn later—cell membranes, enzymes, DNA, and even ecosystem-level processes—depends on how atoms interact. If you understand a few core chemical ideas (how electrons are arranged, what makes a bond polar, why water forms hydrogen bonds), the rest of AP Biology becomes much easier to reason through instead of memorize.

Matter, elements, and atoms

Matter is anything that has mass and occupies space. Living things are made of matter, but biology focuses on a relatively small subset of the periodic table.

An element is a pure substance made of only one type of atom and cannot be broken down into simpler substances by chemical means. An atom is the smallest unit of an element that still has that element’s properties; atoms are the building blocks of the physical world.

Biology emphasizes a handful of elements, especially oxygen (O), carbon (C), hydrogen (H), and nitrogen (N), because these are used to build carbohydrates, proteins, lipids, and nucleic acids (and thus contribute to cells and storage compounds across organisms). Some elements are required only in tiny amounts and are called trace elements, such as iron (Fe), iodine (I), and copper (Cu). Other elements are present in organisms in smaller quantities as well.

Subatomic particles and electron behavior

Inside atoms:

• Protons are positively charged and located in the nucleus.
• Neutrons are uncharged and located in the nucleus.
• Electrons are negatively charged and occupy regions around the nucleus.

A core biological idea is that electron behavior determines how atoms bond, and bonding helps determine molecular shape and properties.

Isotopes

Atoms of the same element always have the same number of protons, but they can differ in the number of neutrons. These variants are called isotopes.

Why valence electrons control bonding

Electrons are arranged in energy levels, and the outermost electrons are called valence electrons. Valence electrons are the ones involved in chemical bonding.

Atoms tend to form bonds that result in a more stable electron configuration (often a filled outer shell) and lower potential energy. It’s tempting to say atoms “want” an octet, but that wording can mislead you—atoms don’t have goals.

Compounds and chemical bonds (strong and weak)

A compound consists of two or more elements chemically combined. The atoms in compounds are held together by chemical bonds, especially covalent and ionic bonds. In biological systems, molecules and different parts of large molecules can also be stabilized by weaker attractions such as hydrogen bonds and van der Waals interactions.

Covalent bonds

A covalent bond forms when atoms share electrons. Covalent bonds are strong and form the backbone of biological molecules.

• Nonpolar covalent bond: electrons are shared equally.
• Polar covalent bond: electrons are shared unequally, creating partial charges.

Bond polarity depends on electronegativity, a measure of how strongly an atom attracts shared electrons. Oxygen and nitrogen are highly electronegative, so O–H and N–H bonds are usually polar.

Polarity matters because partial charges enable hydrogen bonding and other attractions. Many biological properties (water’s behavior, protein folding, DNA base pairing) depend on polarity.

Ionic bonds

An ionic bond forms when electrons are transferred from one atom to another, producing ions.

• A cation is positively charged (lost electrons).
• An anion is negatively charged (gained electrons).

In living systems, ionic compounds often dissociate in water and exist as separate ions because water stabilizes charges. So in biology you frequently think in terms of “ions in solution” rather than rigid ionic crystals.

Hydrogen bonds

A hydrogen bond is a weak attraction between a partially positive hydrogen (already covalently bonded to an electronegative atom like O or N) and a partially negative atom (like O or N) nearby.

Hydrogen bonds are individually weak, but many hydrogen bonds together are powerful—central to water’s cohesion, DNA base pairing, and protein secondary structure.

van der Waals interactions

van der Waals interactions are very weak, short-range attractions caused by temporary uneven electron distributions. They matter when many atoms pack closely, such as within folded proteins.

Chemical reactions in biology: rearranging matter

A chemical reaction rearranges atoms by breaking and forming chemical bonds. Atoms are conserved; they are not created or destroyed in ordinary biological reactions.

Example: why salt dissolves in water (a bond vs. interaction story)

Table salt is held together by ionic attractions between Na⁺ and Cl⁻. Water is polar, so water molecules surround and stabilize each ion (oxygen end toward Na⁺, hydrogen end toward Cl⁻). The ionic lattice breaks apart because the water-ion attractions compensate for separating ions. This is a classic example of water’s polarity dominating biological chemistry.

Exam Focus

• Typical question patterns:
  • Identify key biological elements (especially CHON) and recognize that some are required only as trace elements (Fe, I, Cu).
  • Distinguish atoms vs. elements vs. compounds, and identify protons/neutrons/electrons in diagrams.
  • Predict whether a covalent bond is polar or nonpolar based on electronegativity and connect polarity to hydrogen bonding.
  • Interpret a diagram showing water molecules surrounding ions or polar groups.
  • Distinguish covalent vs. ionic vs. hydrogen bonds in macromolecule structure.
  • Recognize that isotopes are the same element (same protons) with different neutrons.
• Common mistakes:
  • Treating hydrogen bonds as “strong like covalent bonds.” They’re weaker individually, but numerous hydrogen bonds can stabilize large structures.
  • Claiming ionic bonds are “always strong” in cells. In aqueous environments, ions are often dissociated and stabilized by water.
  • Saying atoms “want” electrons. Better: atoms form bonds that lead to more stable electron configurations and lower potential energy.

Water and Hydrogen Bonding: The Molecule That Makes Life Possible

Water is not just the background for biology—it is an active participant. Hydrogen bonding among water molecules produces special properties that directly support life, including cohesion, adhesion, surface tension, temperature stability, and expansion upon freezing.

Water’s structure and polarity

A water molecule has polar covalent bonds between oxygen and hydrogen. Oxygen is more electronegative, so electron density is pulled toward oxygen.

Because of water’s bent shape, the partial charges do not cancel:

• the oxygen end is partially negative
• the hydrogen ends are partially positive

This polarity allows water molecules to form hydrogen bonds with each other and with other polar molecules.

Cohesion, adhesion, capillary action, and surface tension

Cohesion is the tendency of molecules of the same kind to stick together. In water, cohesion comes from hydrogen bonding.

Adhesion is the tendency of molecules of different kinds to stick together (for example, water sticking to plant cell walls).

Together, cohesion and adhesion explain capillary action, the movement of water through narrow spaces against gravity. This helps account for the ability of water to rise through roots, trunks, and branches.

Surface tension is a measure of how difficult it is to stretch or break the surface of a liquid. At the surface, water molecules form a tight hydrogen-bond network with neighbors beside and below, creating a “skin-like” surface. This is why light objects (and organisms like water striders) can sometimes sit atop water without sinking.

In action: transpiration and water transport in plants

When water evaporates from leaf surfaces (transpiration), it pulls on the column of water in xylem. Cohesion helps keep that column unbroken, and adhesion helps it cling to vessel walls.

A common misconception is that plants “pump” water upward like a heart pumps blood. Instead, much of the upward movement is driven by evaporation and cohesion-tension.

Water’s thermal properties: stabilizing temperature

Water resists temperature change, which matters because biochemical reactions are temperature-sensitive.

• Specific heat / heat capacity: the heat required to raise the temperature of a substance. Water has a high specific heat because hydrogen bonds absorb energy as they break and reform.
• Heat of vaporization: the heat needed to convert liquid water to gas. Water’s high heat of vaporization makes evaporative cooling effective.

In action: sweating and evaporative cooling

When sweat evaporates, it takes heat from your skin because breaking hydrogen bonds requires energy. This helps regulate body temperature.

Ice floats: water’s density anomaly (expansion on freezing)

Most substances become denser as they freeze. Water is unusual: solid water is less dense than liquid water.

When water freezes, molecules form a more stable hydrogen-bonded lattice that spaces them farther apart than in liquid water, so ice floats.

Why it matters: Lakes freeze from the top down. Ice on the surface insulates the water below, allowing aquatic life to survive winter.

Water as a solvent and the hydrophobic effect

A solution is a liquid mixture where a solvent dissolves a solute.

Water is an excellent solvent for ionic and polar substances because it can surround charged or polar groups and form stabilizing interactions.

Nonpolar substances (like oils) do not dissolve well in water. This leads to the hydrophobic effect: nonpolar molecules aggregate to minimize disruption of water’s hydrogen-bond network.

Hydrophobic molecules are not “repelled by water” like magnets repel. The driving force is that water can maintain a more stable hydrogen-bond network when nonpolar surface area exposed to water is minimized.

In action: membrane formation

Cell membranes form because amphipathic phospholipids (polar head, nonpolar tails) self-assemble in water so that nonpolar tails are shielded from water.

Exam Focus

• Typical question patterns:
  • Explain a biological phenomenon (transpiration, capillary action, sweating, ice floating, membrane formation) using hydrogen bonding and polarity.
  • Predict how a change (temperature rise, adding salt, changing solute polarity) affects water behavior.
  • Interpret diagrams of hydrogen bonding between water molecules or between water and solutes.
• Common mistakes:
  • Confusing cohesion (water-water) with adhesion (water-other surface).
  • Saying ice floats “because it’s cold” rather than because hydrogen bonding creates a less dense lattice.
  • Treating hydrophobic interactions as a “bond.” The hydrophobic effect is an emergent property of water’s behavior.

pH, Acids and Bases, and Buffers in Biological Systems

Many biomolecules only work properly within a narrow pH range because pH affects charge, shape, and therefore function (for example, enzyme activity depends on the ionization states of amino acid side chains).

Acids, bases, and the meaning of pH

Reactions are influenced by whether the solution is acidic, basic, or neutral.

In aqueous solutions:

• An acid increases the concentration of hydrogen ions by releasing (donating) H⁺.
• A base decreases the concentration of hydrogen ions by accepting H⁺ and/or increases hydroxide concentration. Many bases release OH⁻ in water, and those OH⁻ then reduce free H⁺ by forming water.

The pH scale is commonly represented from 1 to 14, with 7 as neutral. pH is a logarithmic measure of hydrogen ion concentration:

pH = -\log[H^+]

A change of 1 pH unit corresponds to a tenfold change in hydrogen ion concentration.

Why pH matters in biology

Many biological molecules contain groups that can gain or lose H⁺. When those groups gain/lose H⁺, their charge changes. Charge changes can alter hydrogen bonding patterns, disrupt ionic interactions, change protein shape, and affect how substrates bind to active sites.

Water can act as an acid or a base

Water can dissociate into ions (to a very small extent). The product of the hydrogen and hydroxide ion concentrations in water is constant at a given temperature:

K_w = [H^+][OH^-]

At 25°C, pure water has equal [H⁺] and [OH⁻], giving pH 7.

Worked examples: interpreting pH quantitatively

Example 1: Calculate pH from hydrogen ion concentration

If a solution has [H^+] = 1 \times 10^{-3} M:

pH = -\log(1 \times 10^{-3}) = 3

So the solution is acidic.

Example 2: Compare acidity by pH difference

A solution with pH 5 has 100 times more H⁺ than a solution with pH 7, because each pH unit is a factor of 10.

A common misconception is to treat pH as linear (“pH 6 is just a little more acidic than pH 7”). It’s actually ten times more concentrated in H⁺.

Buffers: resisting pH change

A buffer is a solution that minimizes changes in pH when acids or bases are added. Buffers typically consist of a weak acid (can donate H⁺) and its conjugate base (can accept H⁺).

Mechanistically:

• If you add acid (H⁺), the conjugate base binds some of it, preventing a large increase in free H⁺.
• If you add base (which removes H⁺), the weak acid donates H⁺, preventing a large drop in free H⁺.

In action (biology example): blood buffering

Human blood maintains a narrow pH range using buffer systems (including carbonic acid/bicarbonate). You do not need every reaction detail here, but you do need the core logic: buffers protect biological systems from dangerous pH swings.

Exam Focus

• Typical question patterns:
  • Identify whether a solution is acidic/basic/neutral from a pH value or from relative [H⁺]/[OH⁻].
  • Calculate pH from a given hydrogen ion concentration (often powers of 10).
  • Explain why a small pH change can cause major biological effects.
  • Predict how a buffer responds when acid or base is added.
• Common mistakes:
  • Forgetting the log nature of pH (1 unit equals a tenfold change in [H⁺]).
  • Mixing up acids and bases by “strength” rather than effect on [H⁺].
  • Thinking buffers keep pH perfectly constant. Buffers resist change, but they can be overwhelmed.

Carbon and Functional Groups: The Backbone of Organic Molecules

Organic vs. inorganic (as used in biology)

In biology, molecules that contain carbon are often categorized as organic molecules, while many compounds that do not contain carbon are called inorganic. (There are exceptions in chemistry, but this classification is a useful AP Biology rule of thumb.)

Why carbon is uniquely suited for life

Life on Earth is carbon-based because carbon is a versatile atom: it can bind to other carbon atoms and also to elements such as hydrogen, oxygen, and nitrogen. Carbon has four valence electrons, so it can form four covalent bonds, enabling:

• long chains
• branched structures
• rings
• single and double bonds (and rarely triple bonds in biology)

This flexibility allows organisms to build diverse molecules from a small set of elements.

Hydrocarbons and energy storage

Hydrocarbons contain only carbon and hydrogen. They are largely nonpolar and therefore hydrophobic.

Hydrocarbon-rich molecules are often energy dense because they contain many C–H bonds. Many lipids have large hydrocarbon regions and serve as long-term energy storage.

Isomers: same formula, different structure

Isomers are compounds with the same molecular formula but different structures.

• Structural isomers: differ in the arrangement of covalent bonds.
• Geometric isomers: differ in spatial arrangement around a double bond (cis/trans).
• Enantiomers: mirror-image isomers.

Biology is intensely shape-dependent: small changes in 3D structure can dramatically change function.

Functional groups: predictable chemistry attached to carbon skeletons

A functional group is a specific group of atoms attached to a carbon skeleton that gives a molecule characteristic properties. Functional groups strongly influence polarity, acidity/basicity, and reactivity.

Functional groups to recognize:

• Hydroxyl (–OH): polar; common in sugars and alcohols
• Carbonyl (>C=O): found in aldehydes and ketones; affects shape/reactivity
• Carboxyl (–COOH): acidic; can donate H⁺ and become negatively charged
• Amino (–NH₂): basic; can accept H⁺ and become positively charged
• Phosphate (–PO₄²⁻ or related): often negatively charged; key in energy transfer and nucleic acids
• Sulfhydryl (–SH): can form disulfide bonds in proteins

Memory aid (when useful): “-ol” often signals hydroxyl; “amino” implies nitrogen and basicity; “phospho-” often signals negative charge and energy transfer.

In action: why phosphorylation changes protein behavior

Adding a phosphate group introduces negative charges and changes how a protein interacts with water and other molecules. This can alter shape and activity, which is why phosphorylation is common in cell signaling.

Exam Focus

• Typical question patterns:
  • Explain why carbon can form diverse molecules (four valence electrons; chains/rings).
  • Identify functional groups in a diagram and predict properties (polarity, acidity, ability to form hydrogen bonds).
  • Reason about how isomers could have different biological effects.
• Common mistakes:
  • Treating “functional group” as just a vocabulary label rather than a predictor of chemical behavior.
  • Assuming isomers behave similarly because they share a formula; in biology, shape differences often dominate.
  • Confusing phosphate (often charged and hydrophilic) with hydrocarbon regions (nonpolar and hydrophobic).

Building and Breaking Macromolecules: Dehydration, Hydrolysis, and Structure-Function

Four classes of organic compounds are central to life on Earth:

1. carbohydrates
2. proteins
3. lipids
4. nucleic acids

Monomers and polymers

A monomer is a small molecular subunit that can be linked into a larger structure. A polymer is a long molecule made of repeating monomer units.

Many biological macromolecules are polymers:

• carbohydrates (many sugars)
• proteins (many amino acids)
• nucleic acids (many nucleotides)

Lipids are an important exception: they are large biological molecules but are not true polymers with repeating monomer units in the same way.

Dehydration synthesis (condensation): building polymers

Dehydration synthesis forms a covalent bond between monomers while removing a molecule of water. A larger compound is formed and water is produced.

A common misconception is that dehydration here means “drying out.” In this context, it specifically means water is lost as a product during bond formation.

Hydrolysis: breaking polymers

Hydrolysis breaks a polymer into monomers by adding water. The water helps break the bond between monomers. Many digestion reactions are hydrolysis reactions.

Structure determines function (a reasoning tool)

In AP Biology, “structure and function” is a cause-and-effect framework:

• The sequence of monomers changes a polymer’s shape.
• Shape changes the interactions it can form.
• Interactions determine biological function.

You will see this repeatedly in carbohydrates (alpha vs. beta linkages), proteins (sequence → folding → active site shape), and nucleic acids (sequence encodes information).

In action: why cellulose is structural while starch is storage

Cellulose and starch are both glucose polymers, but the type of glycosidic linkage differs, changing 3D shape and therefore function.

Exam Focus

• Typical question patterns:
  • Identify whether a diagram shows dehydration synthesis (condensation) or hydrolysis.
  • Predict how changing monomer order changes polymer properties.
  • Use structure-function logic to explain why two similar molecules behave differently.
• Common mistakes:
  • Mixing up dehydration (builds, produces water) and hydrolysis (breaks, uses water).
  • Calling lipids “polymers” in the same sense as proteins or nucleic acids.
  • Treating structure-function as memorization rather than connecting bonding and shape to biological behavior.

Carbohydrates: Energy, Structure, and Cell Recognition

Carbohydrates are often introduced as “sugars and starches,” but in biology they also serve structural roles and act as identification tags on cell surfaces.

The basic chemistry of carbohydrates

Carbohydrates are organic compounds containing carbon, hydrogen, and oxygen, often in an approximate ratio of 1:2:1. What matters most biologically is how those atoms are arranged and linked.

Most carbohydrates are categorized as:

• Monosaccharides (one sugar; monomers)
• Disaccharides (two sugars)
• Polysaccharides (many sugars; polymers)

Monosaccharides often form rings in water, and ring structure influences how they link together.

Monosaccharides: quick energy and building blocks

Monosaccharides are an important energy source for cells and also serve as building blocks.

Common examples include glucose, fructose, and galactose. The two most commonly emphasized simple sugars are glucose and fructose, which share the formula C6H12O6. Glucose is a major fuel for cells, is an important part of many foods, and is produced by plants during photosynthesis.

Glucose and fructose can be depicted as either “straight” chains or rings, with OH and H groups attached.

Disaccharides and glycosidic linkages

When two monosaccharides are joined, the bond is a glycosidic linkage, formed by dehydration synthesis. The disaccharide formed from two glucose molecules is maltose.

To break a disaccharide back into two monosaccharides, add water (hydrolysis).

Glycosidic linkage geometry: alpha vs. beta

Small changes in linkage orientation can drastically change properties. In glucose polymers:

• alpha linkages often create helical structures and are commonly digestible by humans.
• beta linkages create straighter chains that can hydrogen bond to neighboring chains, producing strong fibers.

This is a key structure-function pattern: same monomer (glucose), different bond geometry, different polymer shape, different function.

Polysaccharides: storage vs. structure

Polysaccharides are made of many repeated monosaccharide units and can be branched or unbranched.

Polysaccharides you should know:

• Starch: plant energy storage; alpha-linked glucose.
• Glycogen: animal energy storage; alpha-linked glucose and typically more highly branched than starch, supporting rapid glucose addition/removal.
• Cellulose: plant structural support in cell walls; beta-linked glucose. Straight chains align and form many hydrogen bonds, creating strong fibers.
• Chitin: structural polysaccharide in arthropod exoskeletons and fungal cell walls; built from beta-linked modified glucose subunits (a nitrogen-containing derivative), which helps explain its strength.

In action: why you can digest starch but not cellulose

Human digestive enzymes can hydrolyze alpha glycosidic linkages in starch, but we lack enzymes to break beta linkages in cellulose. Some animals rely on symbiotic microbes to digest cellulose.

Carbohydrates on cell surfaces: recognition and communication

Carbohydrates can attach to proteins or lipids on the cell surface:

• Glycoproteins: proteins with carbohydrates attached
• Glycolipids: lipids with carbohydrates attached

These surface carbohydrates act like “ID badges” for cell-cell recognition and signaling. Carbohydrates are not only energy molecules; their informational role on membranes is also important.

Exam Focus

• Typical question patterns:
  • Recognize carbohydrate categories (mono-, di-, polysaccharides) and interpret ring vs. chain drawings.
  • Compare starch, glycogen, and cellulose based on linkage type and function; connect branching to rapid access.
  • Identify glycosidic linkages and predict digestibility/structure based on alpha vs. beta.
  • Identify maltose as a disaccharide made from two glucose units.
  • Explain how carbohydrates contribute to cell recognition (glycoproteins/glycolipids).
• Common mistakes:
  • Assuming all polysaccharides are used for energy storage.
  • Missing that alpha vs. beta linkages change 3D structure and hydrogen bonding patterns.
  • Confusing “branched” with “stronger”; branching often relates to storage and rapid access, not structural strength.

Lipids: Membranes, Energy Storage, and Signaling

Lipids are defined less by one shared structure and more by a shared property: they are largely hydrophobic (nonpolar). Their nonpolar character makes them essential for membranes, long-term energy storage, insulation, and signaling.

Why lipids behave differently in water

Most lipids have extensive hydrocarbon regions, making them nonpolar. In water, nonpolar molecules cluster due to the hydrophobic effect, which is foundational to membrane formation.

Fats (triacylglycerols): concentrated energy storage

A common fat molecule is a triacylglycerol (triglyceride):

• one glycerol (the glycerol backbone)
• three fatty acids

Humans store fat in adipose tissue, which contains cells called adipocytes filled with triglycerides.

A fatty acid is a long hydrocarbon chain; one end includes a carboxyl group.

Because triglycerides contain many energy-rich C–H bonds, they store a lot of chemical energy.

Saturated vs. unsaturated fatty acids (lipid saturation)

• Saturated fatty acids have no carbon-carbon double bonds; they tend to be straight and pack tightly.
• Unsaturated fatty acids have one or more double bonds (often cis), creating kinks that prevent tight packing.

Lipid saturation refers to the extent of saturation in a lipid: the more double bonds present, the more unsaturated the lipid is.

Biologically, packing affects whether fats are solid or liquid at room temperature and can influence membrane fluidity when similar tails are present in membrane lipids.

Phospholipids: the architecture of cell membranes

A phospholipid has:

• a hydrophilic (polar) head that includes a negatively charged phosphate group
• two hydrophobic (nonpolar) fatty acid tails

Because they have both hydrophilic and hydrophobic regions, phospholipids are amphipathic. In water, they spontaneously form bilayers with tails inward and heads outward, driven by thermodynamics (water minimizes disruption to its hydrogen-bond network).

In action: why membranes form without “glue”

Membranes do not require a protein scaffold to hold them together. The bilayer is stabilized largely by the hydrophobic effect plus van der Waals interactions among tails.

Steroids: structure and signaling

Steroids are lipids with a characteristic four-ring carbon skeleton.

• Some steroids function as hormones.
• Cholesterol is a steroid found in membranes. It generally increases membrane fluidity (while helping reduce excessive fluidity at very high temperatures) and is also used to make certain hormones and vitamin D.

Waxes and other lipids

Waxes are long-chain lipids that are highly hydrophobic and often function in waterproofing (plant leaf surfaces, some animal coatings).

Exam Focus

• Typical question patterns:
  • Identify lipid types from structural diagrams (triglyceride vs. phospholipid vs. steroid).
  • Explain how phospholipid structure leads to bilayer formation and membrane properties.
  • Compare saturated vs. unsaturated fatty acids and predict packing/fluidity effects.
  • Connect cholesterol’s structure (four rings) to its roles in membranes and as a precursor for hormones/vitamin D.
• Common mistakes:
  • Thinking phospholipids are “held together by covalent bonds” across the bilayer; the bilayer is stabilized mainly by hydrophobic interactions.
  • Confusing triglycerides (energy storage, common in adipocytes) with phospholipids (membranes).
  • Treating “unsaturated = unhealthy” as a biology rule; AP Biology emphasizes structure-function, not nutrition myths.

Proteins: Amino Acids, Folding, and Functional Diversity

Proteins are among the most structurally and functionally diverse macromolecules. They act as enzymes, transporters, receptors, structural fibers, motors, and signaling molecules. Proteins can do so many jobs because they can fold into precise 3D shapes—and shape controls function.

Amino acids: the monomers of proteins

Amino acids are the building blocks of proteins and contain carbon, hydrogen, oxygen, and nitrogen. There are 20 different amino acids.

Each amino acid has a central carbon bonded to:

• an amino group (–NH₂)
• a carboxyl group (–COOH)
• a hydrogen
• an R group (side chain)

Amino acids differ only in their R groups, and R-group chemistry drives protein folding and function.

Spotting an amino acid in a diagram: look for both an amino group and a carboxyl group attached to the same central carbon.

Side-chain polarity and AP-style categories

Side-chain polarity affects whether an amino acid is more hydrophobic or hydrophilic. A useful AP grouping is:

• Hydrophobic (nonpolar and uncharged)
• Hydrophilic (polar and uncharged)
• Ionic (polar and charged)

Among commonly emphasized amino acids:

• Glutamic acid and aspartic acid can donate a proton, making them negatively charged.
• Lysine and arginine can accept a proton at physiological pH, making them positively charged.
• Two amino acids contain sulfur: methionine and cysteine.

Peptide bonds, dipeptides, and polypeptides

When two amino acids join, they form a dipeptide via dehydration synthesis: the carboxyl group of one reacts with the amino group of another.

The bond formed is a peptide bond.

A chain of amino acids is a polypeptide. Once a polypeptide twists and folds, it forms a functional 3D structure called a protein (some proteins consist of multiple polypeptides).

Levels of protein structure and stabilizing forces

Protein structure is described at four levels:

1. Primary structure: linear amino acid sequence (held by covalent peptide bonds).
2. Secondary structure: local folding (alpha helices and beta-pleated sheets), stabilized mainly by hydrogen bonds between backbone atoms.
3. Tertiary structure: overall 3D shape of one polypeptide, stabilized by R-group interactions (hydrogen bonds, ionic interactions, hydrophobic interactions, van der Waals interactions, and sometimes disulfide bonds).
4. Quaternary structure: arrangement of multiple polypeptide subunits, stabilized by similar interactions as tertiary structure.

In action example: hemoglobin

Hemoglobin, which helps distribute oxygen in blood, is a classic quaternary structure protein formed by the interaction of four polypeptide chains.

The hydrophobic effect in protein folding

Many proteins fold so that nonpolar R groups are buried inside and polar/charged R groups are exposed to water. This is largely driven by the hydrophobic effect.

It’s more accurate to say proteins fold into shapes that let surrounding water maintain a stable hydrogen-bond network, rather than saying proteins fold because of “attraction to water.”

Denaturation: when proteins lose structure

Denaturation is the loss of a protein’s native shape due to changes in temperature, pH, or salt concentration.

• Heat increases molecular motion and can disrupt weak interactions.
• pH changes can alter charges on R groups, disrupting ionic interactions and hydrogen bonding.

Denaturation often destroys function (especially for enzymes and receptors) even if the primary structure remains intact.

In action: pH and protein activity

Digestive enzymes function in very different pH environments (acidic stomach vs. more neutral intestines). Proteins working in each location are adapted to those conditions.

Exam Focus

• Typical question patterns:
  • Identify amino acid structure (amino group, carboxyl group, R group) and categorize amino acids by side-chain properties (hydrophobic, hydrophilic, ionic).
  • Predict how temperature, pH, or salt changes affect protein structure and function.
  • Connect a change in amino acid sequence to altered folding and function.
  • Identify which interactions stabilize which levels of structure (especially hydrogen bonds for secondary structure).
  • Recognize hemoglobin as an example of quaternary structure.
• Common mistakes:
  • Saying primary structure is held by hydrogen bonds; primary structure is held by covalent peptide bonds.
  • Assuming denaturation always breaks peptide bonds; denaturation usually disrupts secondary/tertiary/quaternary structure without breaking primary structure.
  • Treating “hydrophobic bonds” as a real bond type; hydrophobic clustering is driven by water’s properties.

Nucleic Acids: Information Storage, Transmission, and Energy Transfer

Nucleic acids store and transmit biological information. DNA contains hereditary instructions, while RNA is essential for protein synthesis and broader gene expression/regulation. A nucleotide derivative, ATP, is central to cellular energy transfer.

Nucleotides: the monomers of nucleic acids

Nucleic acids contain carbon, hydrogen, oxygen, nitrogen, and phosphorus. They are made of monomers called nucleotides, each with:

• a phosphate group
• a five-carbon sugar (ribose in RNA, deoxyribose in DNA)
• a nitrogenous base

Nitrogenous bases include:

• Purines (two-ring): adenine, guanine
• Pyrimidines (one-ring): cytosine, thymine (DNA), uracil (RNA)

DNA vs. RNA (core structural differences)

DNA (deoxyribonucleic acid):

• sugar: deoxyribose
• bases: A, T, C, G
• usually double-stranded
• long-term information storage (hereditary “blueprints”)

RNA (ribonucleic acid):

• sugar: ribose
• bases: A, U, C, G
• usually single-stranded (but can fold)
• essential roles in protein synthesis and gene expression/regulation

How nucleotides link: the sugar-phosphate backbone

Nucleotides connect via covalent bonds between the sugar of one nucleotide and the phosphate of the next, forming a sugar-phosphate backbone. The base sequence extending from the backbone carries information.

DNA double helix: complementary base pairing

In DNA, two strands align with bases pairing via hydrogen bonds:

• A pairs with T
• C pairs with G

The strands run in opposite directions (antiparallel). Base pairing depends on hydrogen bonding and shape complementarity.

In action: predicting a complementary DNA sequence

If one strand contains:

• A G T C C A

The complementary strand is:

• T C A G G T

This works because pairing rules reflect compatible hydrogen-bonding patterns and shapes; mismatches are less stable.

ATP: a nucleotide with an energy role

ATP (adenosine triphosphate) is a nucleotide-related molecule used for energy transfer in cells. A key Unit 1 idea is that ATP’s structure—especially the phosphate groups and their charge interactions—makes ATP useful for coupling energy-requiring processes to energy-releasing reactions.

Biology often describes ATP as having “high-energy bonds,” but the deeper idea is that ATP hydrolysis leads to products that are more stable in aqueous conditions (including reduced electrostatic repulsion and resonance stabilization). You are not expected to master the physical chemistry, but you should connect ATP’s function to its chemical structure.

Exam Focus

• Typical question patterns:
  • Identify parts of a nucleotide and distinguish DNA vs. RNA from diagrams.
  • Use base-pairing rules to predict complementary sequences.
  • Explain how hydrogen bonding enables DNA base pairing and contributes to double-helix stability.
  • Connect ATP’s role in energy transfer to its phosphate-rich structure.
• Common mistakes:
  • Mixing up ribose vs. deoxyribose and thymine vs. uracil.
  • Thinking DNA strands are held together by covalent bonds; the two strands are held together primarily by hydrogen bonds between bases.
  • Treating ATP as “stored energy in a bond” without linking to molecular stability and interaction changes after hydrolysis.