GACE Science (6–12) (717)

All matter consists of atoms, the smallest unit of an element retaining that element's chemical properties. Atoms contain three subatomic particles: protons (positive charge, located in nucleus), neutrons (neutral charge, located in nucleus), and electrons (negative charge, orbiting nucleus in energy levels).

Atomic number equals the number of protons and defines the element. Mass number equals protons plus neutrons. Isotopes are atoms of the same element with different numbers of neutrons. Ions form when atoms gain or lose electrons, becoming negatively charged anions or positively charged cations.

Electrons occupy energy levels (shells) around the nucleus. The first shell holds maximum 2 electrons, second holds 8, third holds 18. Valence electrons occupy the outermost shell and determine chemical reactivity. Elements in the same group share similar valence electron configurations and chemical properties.

Orbital notation describes electron distribution using s, p, d, and f subshells. The aufbau principle states electrons fill lowest energy orbitals first. Hund's rule states electrons occupy orbitals singly before pairing. The Pauli exclusion principle limits each orbital to two electrons with opposite spins.

The periodic table arranges elements by increasing atomic number. Periods are horizontal rows; elements in the same period have the same number of electron shells. Groups are vertical columns; elements in the same group share similar chemical properties due to identical valence electron counts.

Major group classifications include alkali metals (Group 1, highly reactive, one valence electron), alkaline earth metals (Group 2, two valence electrons), halogens (Group 17, seven valence electrons, highly reactive nonmetals), and noble gases (Group 18, full valence shells, chemically inert).

Atomic radius decreases left to right across a period (increasing nuclear charge pulls electrons closer) and increases top to bottom within a group (additional electron shells). Ionization energy increases left to right and decreases top to bottom. Electronegativity increases left to right and decreases top to bottom. Metallic character decreases left to right and increases top to bottom.

Ionic bonds form through electron transfer between metals and nonmetals. Metals lose electrons to form cations; nonmetals gain electrons to form anions. Electrostatic attraction between oppositely charged ions creates the bond. Ionic compounds form crystal lattice structures, have high melting points, conduct electricity when dissolved or molten, and are generally soluble in water.

Covalent bonds form through electron sharing between nonmetals. Single bonds share one electron pair; double bonds share two pairs; triple bonds share three pairs. Bond strength increases with bond order. Polar covalent bonds form when electrons are shared unequally due to electronegativity differences. Nonpolar covalent bonds involve equal sharing.

Metallic bonds involve a "sea of electrons" shared among metal cations. This electron mobility explains metallic properties: electrical conductivity, thermal conductivity, malleability, ductility, and luster.

Intermolecular forces are attractions between molecules, weaker than chemical bonds. London dispersion forces exist in all molecules, arising from temporary dipoles; strength increases with molecular size. Dipole-dipole forces occur between polar molecules. Hydrogen bonding occurs when hydrogen bonds to nitrogen, oxygen, or fluorine and interacts with another electronegative atom; this is the strongest intermolecular force and explains water's unique properties.

The law of conservation of mass requires equal numbers of each atom on both sides of an equation. Balance equations by adjusting coefficients, not subscripts. Balance metals first, then nonmetals, then hydrogen, and finally oxygen.

Stoichiometry uses mole ratios from balanced equations to calculate quantities of reactants and products. One mole equals 6.022 × 10²³ particles (Avogadro's number). Molar mass equals the mass of one mole of a substance in grams.

Limiting reactants are completely consumed and determine maximum product yield. Excess reactants remain after reaction completion. Percent yield equals actual yield divided by theoretical yield, multiplied by 100.

Reaction rate measures how quickly reactants convert to products. Factors affecting rate include temperature (higher temperature increases rate), concentration (higher concentration increases rate), surface area (greater surface area increases rate), and catalysts (lower activation energy without being consumed).

Chemical equilibrium occurs when forward and reverse reaction rates are equal. Le Chatelier's principle states that systems at equilibrium shift to counteract applied stresses. Adding reactants shifts equilibrium toward products; removing products has the same effect. Temperature changes favor endothermic or exothermic direction depending on whether heat is added or removed.

Acids donate hydrogen ions (H⁺), taste sour, have pH below 7, turn litmus red, and react with metals to produce hydrogen gas. Bases accept hydrogen ions, taste bitter, feel slippery, have pH above 7, and turn litmus blue.

Strong acids and bases dissociate completely in water. Weak acids and bases partially dissociate. Neutralization reactions between acids and bases produce water and a salt.

The pH scale ranges from 0 to 14. pH equals the negative logarithm of hydrogen ion concentration. Each pH unit represents a tenfold change in H⁺ concentration. Neutral solutions have pH 7. Acidic solutions have pH below 7. Basic solutions have pH above 7.

Solutions are homogeneous mixtures of solute dissolved in solvent. Solubility is the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature. Factors affecting solubility include temperature (solubility of solids generally increases with temperature; gases decrease), pressure (affects gas solubility per Henry's law), and nature of solute and solvent ("like dissolves like").

Concentration measures include molarity (moles solute per liter solution), molality (moles solute per kilogram solvent), and percent composition.

Colligative properties depend on solute particle number, not identity. Boiling point elevation occurs when solute particles raise the boiling point. Freezing point depression lowers the freezing point. Osmotic pressure drives solvent movement across semipermeable membranes.