Exam 1

CHAPTER 1

CHAPTER 1

Describe the 7 properties of life

1. Organization — life is highly organized, from atoms and molecules to cells and larger systems.

2. Metabolism / energy processing — organisms obtain and transform energy and matter.

3. Response to stimuli — organisms respond to changes in their environment.

4. Regulation / homeostasis — organisms regulate internal conditions.

5. Growth and development — organisms grow and undergo genetically programmed development.

6. Reproduction — organisms produce offspring and pass genetic information to them.

7. Evolution/adaptation — populations change over generations; natural selection can produce adaptations.

Important distinction: an individual organism grows and develops during its lifetime, but evolution is a population-level change across generations.

Levels of Biological Organization

Atom → molecule → organelle → cell → tissue → organ → organ system → organism → population → community → ecosystem → biosphere.

Know these especially well:

• Population = members of the same species living in the same area.

• Community = all populations of different species in an area.

• Ecosystem = a community plus its nonliving/abiotic environment.

• Biosphere = all regions of Earth where life exists.

Emergent Properties

Emergent properties result from the arrangement and interaction of parts as complexity increases. A system can have properties that its individual components do not have by themselves.

Example from lecture: a functioning bicycle emerges when all of its necessary parts are connected correctly.

Reductionism studies complex systems by breaking them into manageable components. Systems biology examines interactions among the parts of a biological system. Structure and function are also closely related: a structure’s form helps determine what it can do.

Themes: Organization, prokaryotic and eukaryotic

One of the five unifying themes is Organization. The cell is the smallest unit of organization that can perform all activities required for life.

Prokaryotic cells:

• Bacteria and Archaea

• lack a membrane-bound nucleus

• lack membrane-bound organelles

Eukaryotic cells:

• have a membrane-bound nucleus

• have membrane-bound organelles

• include plants, fungi, animals, and protists

Five themes to recognize: Organization, Information, Energy and Matter, Interactions, Evolution.

Themes: Genetic Information, gene expression

Genetic information is stored in DNA. Genes are units of inheritance that contain information used to make cellular products.

For protein-coding genes, gene expression can be summarized: DNA → RNA → protein

Transcription produces RNA from DNA information. Translation uses RNA information to produce a polypeptide/protein.

A major theme is that information is stored, transmitted, and expressed.

Themes: Energy and Matter, energy flow, chemical cycling

Energy flows through ecosystems. Producers capture energy and energy is transferred through organisms.

Chemical matter/nutrients cycle through ecosystems and can be reused.

EXAM TRAP: energy flows; matter/chemicals cycle.

Themes: Interactions, feedback regulation

Organisms interact with other organisms and with their physical environment.

Feedback regulation occurs when the output of a process influences that process:

• Negative feedback counteracts a change and is the most common form.

• Positive feedback reinforces/amplifies a change.

Themes: Evolution, decent with modification, natural selection

Evolution is descent with modification: living organisms are modified descendants of common ancestors.

Darwin’s natural-selection reasoning:

• Individuals in a population vary.

• Many variations are heritable.

• More offspring are produced than survive, creating competition.

• Individuals with traits better suited to their environment are more likely to survive and reproduce.

• Over generations, advantageous heritable traits become more common.

Natural selection is the mechanism that can produce adaptation. Natural selection acts on individuals, but evolution is the change in a population over generations.

Lecture examples include beach and inland mice whose coloration matches their environments, and Galápagos finches descended from a common ancestor.

Three Domains, kingdoms, prokaryote/eukaryote, and an example of organisms

Three domains:

• Bacteria — prokaryotic; example: bacterial cells.

• Archaea — prokaryotic; example: archaeal microorganisms.

• Eukarya — eukaryotic; includes plants, fungi, animals, and protists.

Within Eukarya:

• Plants generally make food by photosynthesis.

• Fungi absorb nutrients.

• Animals ingest food.

• Protists are a diverse group, mostly single-celled in the lecture.

EXAM TRAP: Bacteria + Archaea = prokaryotes; Eukarya = eukaryotes.

Scientific method, control, variables, hypothesis

Scientific inquiry is an idealized process, not a rigid recipe. A common pattern is: observation → question → hypothesis → prediction → experiment/test → data → conclusion.

Definitions:

• Observation — something noticed or measured.

• Data — recorded observations. Qualitative data are descriptive; quantitative data are numerical measurements.

• Hypothesis — an explanation based on observations/assumptions that leads to a testable prediction.

• Prediction — what should happen if the hypothesis is correct.

• Experiment — a scientific test carried out under controlled conditions.

• Independent variable — factor deliberately changed/manipulated by the researcher.

• Dependent variable — measured response.

• Control group — comparison group that does not receive the experimental condition.

• Experimental group — group exposed to the condition being tested.

Reasoning:

• Inductive reasoning = specific observations → generalization.

• Deductive reasoning = general premise → specific prediction.

A hypothesis must be testable. Scientists can increase confidence in a hypothesis through repeated testing, but cannot prove a hypothesis true with absolute certainty. Scientific inquiry may backtrack or change direction as new observations appear.

Scientific theory

A scientific theory is a broad, well-supported explanatory framework. It is supported by a large body of evidence and can generate many testable hypotheses.

EXAM TRAP: In science, “theory” does NOT mean a guess or an unsupported opinion. A hypothesis is a testable explanation; a theory is broader and supported by extensive evidence.

Science vs technology

Science seeks to understand natural phenomena and how the natural world works. Technology applies scientific knowledge to accomplish a specific purpose or solve a problem.

They are interdependent: scientific discoveries can lead to new technologies, and new technologies can give scientists better ways to investigate nature.

CHAPTER 2

CHAPTER 2

Which 4 elements are most common in living things

Carbon (C), hydrogen (H), oxygen (O), and nitrogen (N) — CHON — make up about 96% of living matter according to the course lecture.

The lecture notes that most of the remaining 4% consists of calcium, phosphorus, potassium, and sulfur. Trace elements are required only in minute quantities but may still be essential; iron, for example, is required for hemoglobin.

Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.

• Atomic number = number of protons.

• Mass number = protons + neutrons.

• Changing the number of neutrons changes the isotope, not the element.

• Radioactive isotopes are unstable and decay spontaneously, releasing particles and energy.

• Radioactive isotopes can be used as tracers and in radiometric dating; a parent isotope decays into a daughter isotope at a predictable rate called its half-life.

Valence shell, “Octet rule”

The valence shell is the outermost electron shell. Valence electrons largely determine chemical behavior.

The octet rule describes the tendency of atoms to gain, lose, or share electrons to obtain a stable outer shell, commonly eight electrons.

An element with a full valence shell is relatively chemically inert. The number of unpaired valence electrons is generally related to an atom’s valence, or bonding capacity.

Covalent bond, single, double

A covalent bond forms when atoms share electrons.

• Single covalent bond = one shared pair of electrons.

• Double covalent bond = two shared pairs of electrons.

A structural formula uses lines to represent covalent bonds. A molecule consists of two or more atoms held together by covalent bonds.

Polar and nonpolar covalent bonds

Nonpolar covalent bond = electrons are shared equally or nearly equally.

Polar covalent bond = electrons are shared unequally because one atom attracts the shared electrons more strongly.

Unequal sharing produces partial charges (δ− and δ+). Water is a major example of a molecule containing polar covalent bonds.

Electronegativity

Electronegativity is an atom’s attraction for the electrons in a covalent bond. The more electronegative an atom is, the more strongly it pulls shared electrons toward itself.

A greater difference in electronegativity means more unequal electron sharing and greater bond polarity.

Course lecture bond guide:

• 0.0–0.3 = nonpolar covalent

• 0.4–1.9 = polar covalent

• 2.0–3.3 = ionic

Use the course ranges if a numerical bond-polarity question appears.

Ionic bond, ions

An ion is a charged atom or molecule.

• Cation = positively charged; forms when electrons are lost.

• Anion = negatively charged; forms when electrons are gained.

An ionic bond is the attraction between a cation and an anion. For example, sodium transfers an electron to chlorine, producing Na+ and Cl−; their attraction forms sodium chloride.

Ionic compounds/salts are often found as crystals.

Hydrogen bonds

A hydrogen bond forms when a hydrogen atom covalently bonded to an electronegative atom is attracted to another electronegative atom.

In living systems, the electronegative partners are usually oxygen or nitrogen.

Important: hydrogen bonds are weaker than covalent bonds. Covalent bonds hold atoms together within molecules; hydrogen bonds are attractions between appropriate partial charges. Many weak interactions together can strongly affect biological structure.

Van der Waals

Van der Waals interactions are weak attractions between molecules that are close together because electrons can temporarily become unevenly distributed, producing temporary charges.

They are individually weak but can contribute significantly when many occur together.

Molecular shape, orbitals

Molecular size and shape are important to function. An orbital is a three-dimensional region where an electron is found most of the time.

The positions of atoms’ orbitals determine molecular shape. In covalent bonding, s and p orbitals may hybridize to create specific molecular shapes.

EXAM IDEA: structure/shape influences how a molecule interacts with other molecules, which influences function.

Chemical reactions- reactants/products, equilibrium

Chemical reactions rearrange atoms; atoms are not created or destroyed.

• Reactants = starting substances.

• Products = substances formed by the reaction.

• Chemical equilibrium = forward and reverse reactions occur at equal rates, so there is no net change in concentrations.

EXAM TRAP: equilibrium does NOT mean the reactions have stopped. Both directions continue, but at equal rates.

CHAPTER 3

CHAPTER 3

Polar molecules, hydrogen bonds

Water is a polar molecule. Oxygen is more electronegative than hydrogen, so the shared electrons are pulled closer to oxygen. Oxygen has a partial negative charge (δ−), while the hydrogens have partial positive charges (δ+).

Because water is polar, water molecules form hydrogen bonds with one another.

Water’s polarity and hydrogen bonding are responsible for many of its life-supporting properties.

Cohesion, adhesion

Cohesion = attraction between molecules of the same substance. Water molecules stick to other water molecules.

Adhesion = attraction between molecules of different substances. Water can stick to other materials.

Cohesion helps water move as a continuous column through plants. Adhesion helps water interact with the walls of water-conducting cells. Hydrogen bonding also gives water unusually high surface tension.

Specific heat, water/hydrogen bonds, benefits for life

Specific heat = amount of heat that must be absorbed or lost for 1 g of a substance to change its temperature by 1°C.

The specific heat of water is about 1 cal/(g·°C). Water has a high specific heat because hydrogen bonds absorb heat when they break and release heat when they form.

Therefore, water can absorb/release a large amount of heat with only a small temperature change. This moderates temperatures and minimizes fluctuations that could harm organisms.

The lecture also defines:

• Thermal energy = kinetic energy associated with random molecular motion.

• Temperature = average kinetic energy of molecules.

• Heat = thermal energy transferred from one body of matter to another.

• 1 cal = 4.184 J.

Evaporative cooling, heat of vaporization

Evaporation/vaporization = transformation from liquid to gas.

Heat of vaporization = heat a liquid must absorb for 1 g to become a gas.

During evaporation, the higher-energy molecules escape first. The remaining liquid has lower average kinetic energy, so its temperature falls. This is evaporative cooling.

Evaporative cooling helps stabilize temperatures in organisms and bodies of water.

Ice, density, benefits

When water freezes, hydrogen bonds become more ordered and create an open structure. This makes ice less dense than liquid water, so ice floats.

Water reaches its greatest density at about 4°C.

Why floating ice matters: the ice layer can insulate the water underneath, allowing liquid water and aquatic life to persist below the frozen surface.

Solvent/solute, aqueous solution, hydration shell

• Solvent = substance doing the dissolving.

• Solute = substance being dissolved.

• Solution = homogeneous mixture of substances.

• Aqueous solution = solution in which water is the solvent.

• Hydration shell = sphere of water molecules surrounding a dissolved ion.

Water is a versatile solvent because its polarity allows it to interact with charged and polar substances. Water can even dissolve large polar molecules such as proteins if they contain ionic/polar regions.

Hydrophilic/hydrophobic

Hydrophilic = interacts readily with water; usually polar or charged.

Hydrophobic = mixes poorly with water; usually nonpolar.

Oil is hydrophobic because it has relatively nonpolar bonds. Hydrophobic molecules related to oils are important components of cell membranes.

Amphipathic molecules have both hydrophilic and hydrophobic regions.

Acids/bases

An acid increases H+ concentration in a solution. A base decreases H+ concentration and/or increases OH− concentration.

Water can dissociate into H3O+ (hydronium, often represented as H+) and OH−. In pure water, H+ and OH− concentrations are equal.

Strong acids and bases dissociate essentially completely in water. Weak acids and bases reversibly release/accept H+ and can shift the balance away from neutrality.

pH, buffers

pH describes how acidic or basic a solution is and is related to H+ concentration.

• pH < 7 = acidic

• pH = 7 = neutral

• pH > 7 = basic

Lower pH = higher H+ concentration.

Buffers resist sudden changes in pH by accepting or donating H+. A common buffer system contains a weak acid and its corresponding base.

At 25°C, pure water has [H+] = [OH−] = 10−7 M, and [H+][OH−] = 10−14.

Ocean acidification and effects, carbonate

CO2 from the atmosphere can dissolve into seawater and contribute to carbonic acid formation. This increases H+ concentration and lowers ocean pH.

The extra H+ reacts with carbonate ions, reducing the amount of carbonate available. Carbonate is important for making calcium carbonate structures.

Therefore, ocean acidification can make calcification more difficult for organisms such as reef-building corals and other organisms that form calcium-carbonate structures.

CHAPTER 4

CHAPTER 4

Carbon- versatility, electron configuration

Carbon is exceptionally versatile because it has four valence electrons and can form four covalent bonds.

Carbon can bond with itself and with many other elements, especially hydrogen, oxygen, and nitrogen. Carbon skeletons can form:

• long chains

• branched chains

• rings

• combinations of these

Carbon-carbon single, double, and triple bonds can create different molecular geometries. With four single bonds, a carbon generally has a tetrahedral arrangement; a carbon-carbon double bond produces a planar arrangement around the double bond.

The electron configuration of carbon determines the kinds and number of bonds it can form, which contributes to the enormous diversity of biological molecules.

Organic compound

Organic chemistry is the study of compounds that contain carbon. Organic compounds range from simple molecules to very large, complex biological molecules.

The major biological macromolecule classes—carbohydrates, lipids, proteins, and nucleic acids—are carbon-based organic molecules.

Hydrocarbons, non-polar

Hydrocarbons are organic molecules consisting only of carbon and hydrogen.

They are generally nonpolar because C–H bonds do not create large charge differences. Therefore, hydrocarbons are generally hydrophobic.

Hydrocarbon regions can store substantial chemical energy and are common components of lipids such as fats.

Isomers, structural, cis-trans, enantiomers

Isomers have the same molecular formula but different structures and properties.

• Structural isomers — atoms have different covalent arrangements.

• Cis-trans (geometric) isomers — same covalent connections but different spatial arrangements, commonly because rotation around a double bond is restricted. Cis = groups on the same side; trans = groups on opposite sides.

• Enantiomers — mirror-image isomers.

Biological importance: even subtle differences in molecular arrangement can change biological effects. The lecture notes that different enantiomers of a drug may have different effects, and often only one is biologically active.

Functional groups, amino, carboxyl, phosphate

Functional groups are specific groups of atoms attached to carbon skeletons that give organic molecules characteristic properties and are commonly involved in chemical reactions.

The course specifically lists:

• Amino group — −NH2; contains nitrogen and commonly acts as a base/accepts H+. • Carboxyl group — −COOH; contains oxygen and commonly acts as an acid/donates H+.

• Phosphate group — commonly represented as −OPO3²−; often contributes negative charge and is important in energy transfer and nucleic acids.

The broader lecture identifies seven major functional groups: hydroxyl, carbonyl, carboxyl, amino, sulfhydryl, phosphate, and methyl. Know the three specifically named on the Fall 2026 study guide especially well.

ATP

ATP = adenosine triphosphate.

ATP contains adenine, a sugar, and three phosphate groups. Hydrolysis of ATP can release free energy that cells use to drive energy-requiring processes.

ATP is an energy-transfer molecule used for cellular work, not the cell’s main long-term energy-storage molecule.

CHAPTER 5

CHAPTER 5

Monomers/polymers

A monomer is a building-block subunit. A polymer is a long molecule made by linking many similar building blocks.

The four major biological macromolecule classes are:

• Carbohydrates

• Lipids

• Proteins

• Nucleic acids

Carbohydrates, proteins, and nucleic acids are polymers. Lipids are not true polymers.

A huge variety of polymers can be made from a relatively small set of monomers. The order and arrangement of monomers contribute to molecular diversity.

Dehydration/Hydrolysis reaction

Dehydration reaction = joins molecules by removing water.

Hydrolysis = breaks bonds by adding water.

Memory trick:

DEHYDRATION → BUILD

HYDROLYSIS → BREAK

For a polymer made of 10 monomers, complete hydrolysis requires 9 water molecules because 9 bonds connect the 10 monomers.

For each type: monomers, examples, functions

CARBOHYDRATES

• Monomers: monosaccharides

• Examples: glucose, starch, glycogen, cellulose, chitin

• Functions: fuel and building material/structure

LIPIDS

• Not true polymers

• Examples: fats/triglycerides, phospholipids, steroids

• Functions: long-term energy storage, insulation/cushioning, membranes, and signaling

PROTEINS

• Monomers: amino acids

• Examples: enzymes, transport proteins, structural proteins

• Functions: highly diverse; see the eight functions below

NUCLEIC ACIDS

• Monomers: nucleotides

• Examples: DNA and RNA

• Functions: store, transmit, and help express hereditary information

Carbohydrates — monomers

Monosaccharides are the simplest carbohydrates and the monomers/building blocks of larger carbohydrates.

Glucose (C6H12O6) is the most common monosaccharide and an important cellular fuel.

Monosaccharides can be classified by the location of the carbonyl group (aldose vs ketose) and by the number of carbons in the carbon skeleton. In aqueous solutions, many sugars form rings.

Carbohydrates — disaccharide

A disaccharide consists of two monosaccharides joined by a dehydration reaction.

The covalent bond joining the two sugars is a glycosidic linkage.

Carbohydrates — glycosidic linkage

A glycosidic linkage is the covalent bond that joins monosaccharides.

The type and position of glycosidic linkages help determine the structure and function of a polysaccharide.

Carbohydrates — Structural carbohydrates

Cellulose:

• Major component of plant cell walls.

• Polymer of glucose.

• Uses β glycosidic linkages.

• Cellulose molecules are straight and unbranched.

• Hydrogen bonding between parallel cellulose molecules contributes to strength.

• Human digestive enzymes cannot hydrolyze its β linkages, so cellulose passes through as insoluble fiber.

Chitin:

• Structural polysaccharide.

• Found in arthropod exoskeletons.

• Also provides structural support in many fungal cell walls.