General Chemistry: Quantitative Relationships
General Chemistry: quantitative relationships
The OAT General Chemistry section tests stoichiometry, atomic structure and periodic trends, bonding, gases, solutions, thermodynamics, kinetics, equilibrium, acids and bases, redox, and electrochemistry. Many questions are calculation-based, so keep units consistent and track significant figures.
Core rules
- Mole map: balance the equation first; convert mass to moles, apply mole ratios, and convert back. The limiting reagent is the reactant that produces the least product.
- Periodic trends: atomic radius decreases across a period and increases down a group. Ionization energy and electronegativity generally increase across a period and decrease down a group.
- Gases: PV = nRT with temperature in kelvin. At constant T and n, P1V1 = P2V2. Dalton's law says total pressure is the sum of partial pressures.
- Equilibrium: the K expression excludes pure solids and liquids. Le Chatelier's principle predicts shifts; a catalyst changes rate, not K.
- Acids and bases: pH = -log[H+]; strong acids dissociate completely. Buffer capacity is best near the pKa.
Worked example
How many grams of water form when 4.0 g H2 reacts with 32.0 g O2? Balanced equation: 2H2 + O2 -> 2H2O. Moles H2 = 4.0/2.0 = 2.0 mol; moles O2 = 32.0/32.0 = 1.0 mol. The equation needs 2 mol H2 per 1 mol O2, so the reactants are in exact stoichiometric ratio and neither is limiting. Moles H2O = 2.0 mol. Mass = 2.0 x 18.0 = 36.0 g.
Traps
- Using Celsius in gas-law calculations.
- Forgetting to balance equations before applying mole ratios.
- Including pure solids in K or in equilibrium concentration calculations.
- Treating a catalyst as a shift in equilibrium position.