Chemistry Study Guide

1. What Is Chemistry?

Chemistry is the science of matter — what it is made of, how it behaves, and how it changes. From the oxygen you breathe to the medicines that save lives to the polymers in your phone case, chemistry underlies every material interaction in the universe.

Chemistry bridges physics (at the atomic scale) and biology (at the molecular/cellular scale). Physical chemistry explains why reactions occur and how fast. Organic chemistry focuses on carbon-based compounds — the molecules of life. Biochemistry studies chemical processes in living organisms. Analytical chemistry identifies and quantifies substances. Materials science develops new substances with desired properties. Understanding chemistry vocabulary is the first step into all of these fields.

2. Atoms: Structure and Subatomic Particles

An atom is the smallest unit of matter that retains the chemical properties of an element. Atoms consist of three types of subatomic particles: protons (positive charge, in the nucleus), neutrons (no charge, in the nucleus), and electrons (negative charge, orbiting the nucleus in electron shells).

The atomic number = number of protons, which uniquely identifies each element. A neutral atom has equal numbers of protons and electrons. The mass number = protons + neutrons. Isotopes are atoms of the same element with different numbers of neutrons (same atomic number, different mass number). Carbon-12 (6p + 6n) and Carbon-14 (6p + 8n) are isotopes of carbon.

Ions form when atoms gain or lose electrons. A cation (positive ion) has lost electrons (Na → Na⁺). An anion (negative ion) has gained electrons (Cl → Cl⁻). The number of electrons in the outermost shell (valence electrons) determines an element's chemical reactivity and bonding behavior.

3. The Periodic Table

The periodic table organizes all known elements by increasing atomic number into rows (periods) and columns (groups). The rows indicate the number of electron shells; elements in the same period have the same number of shells. The columns indicate valence electrons and correlate with similar chemical properties.

Elements are classified as metals (left and center: good conductors, malleable, lustrous), nonmetals (upper right: poor conductors, brittle as solids, diverse properties), and metalloids (staircase border: intermediate properties — silicon is the basis of computer chips because its semiconductor behavior sits between conductor and insulator).

The noble gases (Group 18: He, Ne, Ar, Kr, Xe, Rn) have full valence shells and are nearly chemically inert. The halogens (Group 17: F, Cl, Br, I) are highly reactive nonmetals with 7 valence electrons that readily gain one more to achieve a full shell. The alkali metals (Group 1: Li, Na, K) are highly reactive with one valence electron they readily lose.

4. Chemical Bonds

A chemical bond is the force of attraction that holds atoms together in a compound. Bonding occurs because combined atoms are at lower energy (more stable) than separate atoms. The three main bond types differ in how electrons are shared or transferred.

An ionic bond forms when one atom transfers electrons to another, creating oppositely charged ions that attract. This typically occurs between metals (electron donors) and nonmetals (electron acceptors). NaCl (table salt): Na donates one electron to Cl → Na⁺ and Cl⁻ attract electrostatically → crystal lattice.

A covalent bond forms when two atoms share electrons. This typically occurs between nonmetals. Single bonds share one pair (H−H, H₂O); double bonds share two pairs (O=O in O₂, C=O in CO₂); triple bonds share three pairs (N≡N in N₂, the strongest common bond). Polar covalent bonds share electrons unequally (as in water — oxygen attracts electrons more strongly than hydrogen), creating partial charges that give water its unique properties.

A metallic bond involves electrons delocalized (shared freely) across a lattice of positive metal ions — the "electron sea" model. This explains metals' electrical conductivity (electrons flow freely), thermal conductivity, and malleability (layers of atoms slide without breaking bonds).

5. Chemical Reactions

A chemical reaction transforms reactants into products by breaking and forming chemical bonds. Evidence of a reaction includes: temperature change, color change, gas production, precipitate formation, and light emission. The Law of Conservation of Mass requires that all atoms present in the reactants appear in the products — nothing is created or destroyed.

Chemical equations represent reactions symbolically. A balanced equation has equal numbers of each atom on both sides: 2H₂ + O₂ → 2H₂O. Coefficients (the large numbers) indicate molar ratios — here, 2 moles of hydrogen react with 1 mole of oxygen to produce 2 moles of water.

Types of reactions: Synthesis (A + B → AB); Decomposition (AB → A + B); Single displacement (A + BC → AC + B); Double displacement (AB + CD → AD + CB); Combustion (hydrocarbon + O₂ → CO₂ + H₂O + energy). Catalysts speed up reactions without being consumed — enzymes are biological catalysts enabling life's chemistry at body temperature.

6. States of Matter

Matter exists in four common states. Solid: fixed shape and volume; particles tightly packed in regular arrangements, vibrating in place. Liquid: fixed volume, variable shape; particles close together but able to flow past each other. Gas: variable shape and volume; particles far apart, moving rapidly and randomly. Plasma: ionized gas of free electrons and ions — the most common state of visible matter in the universe (stars, lightning, fluorescent lights).

State changes require energy input or release: melting (solid→liquid), freezing (liquid→solid), evaporation (liquid→gas), condensation (gas→liquid), sublimation (solid→gas directly, as in dry ice CO₂), and deposition (gas→solid). These changes are physical changes — no new substances form, only intermolecular forces change.

7. Acids and Bases

The Brønsted-Lowry definition: an acid is a proton (H⁺) donor; a base is a proton acceptor. Strong acids (HCl, H₂SO₄, HNO₃) dissociate completely in water; weak acids (acetic acid in vinegar, carbonic acid in soda) dissociate partially. Similarly for strong vs. weak bases.

The pH scale measures acidity/basicity from 0 (most acidic) to 14 (most basic), with 7 being neutral (pure water). pH = −log₁₀[H⁺]. Each unit represents a 10-fold change in hydrogen ion concentration: pH 5 is 10 times more acidic than pH 6. Indicators (like litmus or phenolphthalein) change color based on pH.

Acids and bases neutralize each other, producing a salt and water: HCl + NaOH → NaCl + H₂O. Buffer solutions resist pH changes — blood is buffered at pH ≈ 7.4 by carbonic acid and bicarbonate; even small pH changes impair enzyme function and threaten life.

8. Complete Key Terms Glossary

9. Frequently Asked Questions