Scientific Method and Chemistry Essentials for Biology

Biology is the scientific study of life; science is an evidence-based approach to understanding the natural world, based on inquiry and verifiable evidence.
Scientific knowledge arises from applying the scientific method to answer questions, explain phenomena, and test hypotheses.
Science aims for true understanding, not personal opinions or spiritual beliefs.

Observations and questions: what we can sense, plus prior knowledge or results.
Consult prior knowledge and existing results.
Formulate a hypothesis: a testable, general statement that leads to predictions.
Make predictions: usually written as an if-then statement.
Design and perform experiments under controlled conditions; discovery/observational science as an alternative.
Collect and interpret data; draw conclusions.
Peer review and publishing: manuscript evaluated by editors, then anonymous external reviewers; ensures quality.

A hypothesis is testable and data can support or reject it.
A hypothesis cannot be proven true beyond all doubt; future data may challenge it.
Widely accepted when:
- Several sources of evidence support it $\Rightarrow$ no trustworthy data refute it $\Rightarrow$ other alternative hypotheses rejected.

Initiation steps begin with observations and questions; observations encompass senses or existing knowledge/results.

Investigators draw conclusions based on data.
Data can come from discovery science (descriptive/observational) or controlled experiments testing a hypothesis.

After data analysis, decide whether evidence supports or falsifies the hypothesis.
Sufficient evidence leads to manuscript submission for peer review; editors assess, then referees evaluate methods, data, conclusions.
Peer review helps ensure high-quality journal articles.

Biology rests on chemistry; matter and molecules underpin metabolic processes and life.
Metabolism involves chemical reactions; living systems are built from and organized around chemical bonds and interactions.

Matter is anything that occupies space and has mass; composed of chemical elements.
There are $92$ naturally occurring elements; elements have symbols; arranged in the periodic table.
Elements combine to form compounds; compounds contain two or more different elements in fixed ratios.

Four elements make up $96\%$ of body weight: $O$ , $C$ , $H$ , $N$ with approximate percentages: $O: 65.0\%$ , $C: 18.5\%$ , $H: 9.5\%$ , $N: 3.3\%$ .
Trace elements: < $0.01\%$ of body weight; essential examples include iron, iodine, zinc, selenium, copper, manganese, cobalt, etc.

An atom consists of a nucleus (protons and neutrons) and an electron cloud.
Atomic number = number of protons; example: Hydrogen $1$ , Chlorine $17$ , Uranium $92$ .
Atomic mass is the average mass of all isotopes; expressed in atomic mass units (Da).
1 Da ≈ mass of a hydrogen atom; 1 Da = $1.67\times10^{-24}\text{ g}$ .
Mass number = protons + neutrons; varies by isotope (e.g., $^{12}C$ has mass number $12$ ).

Isotopes have same atomic number but different numbers of neutrons; mass numbers differ; isotopes behave similarly in chemical reactions.
Common carbon isotopes: $^{12}C$ (~ $99\%$ ), $^{13}C$ (~ $1\%$ ), $^{14}C$ (~ $1\%$ ).
Radioactive isotopes have half-lives; half-life is the time for half the atoms to decay.

Elements combine to form compounds: two or more different elements in fixed ratios (e.g., NaCl).
All compounds are molecules, but not all molecules are compounds.
Molecule: two or more atoms bonded together; can be the same element or different elements.
Ionic bonds: transfer of electrons; results in oppositely charged ions (e.g., Na⁺ and Cl⁻ in NaCl).
Covalent bonds: sharing of electrons; can be single, double, or triple; strongest in biology.
Polar covalent bonds arise when electronegativity difference yields partial charges; water is a key example.

Electronegativity differences produce polar covalent bonds; electrons spend more time near the more electronegative atom (δ−) and less near the less electronegative atom (δ+).
Polar covalent bonds are important for hydrogen bonding and biomolecule interactions.

Water is H₂O: two polar O–H covalent bonds; oxygen is δ−, hydrogens are δ+.
Water molecules form hydrogen bonds: electrostatic attraction between a partially positive H and a neighboring electronegative atom (O, N, or F).
Hydrogen bonds are weaker than covalent bonds; typical bond energies:
- Hydrogen bond: $5\text{–}10\%$ of a covalent bond; ~ $23\text{ kJ/mol}$ in liquid water.
- O–H covalent bond: ~ $470\text{ kJ/mol}$ .
- C–C covalent bond: ~ $348\text{ kJ/mol}$ .
Water’s polarity and hydrogen bonding give water’s unusual properties and drive biomolecular interactions.

Water’s cohesive forces enable liquid water at room temperature and contribute to its high surface tension.
Flickering clusters: hydrogen-bond networks constantly break and reform in liquid water (roughly $5\times10^{11}\ \text{s}^{-1}$ ).
Ice lattice: each water molecule forms up to 4 hydrogen bonds; ice is less dense than liquid water and floats, insulating aquatic life.

Water dissolves many substances; a solution contains a solute(s) dissolved in the solvent.
Aqueous solution: solvent is water.
“Like dissolves like”: polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar substances.