Chemistry
General chemistry on the FE underpins corrosion, water treatment, combustion, and materials behavior. The tested material is compact: the mole concept and stoichiometry, acids and bases, oxidation–reduction, solution concentration, and the ideal-gas law — mostly single-step calculations once you set them up correctly.
Moles, molar mass, and stoichiometry
The mole is a fixed count of particles (Avogadro's number, 6.022×10²³). Molar mass, the sum of atomic masses in g/mol, converts between mass and moles: CaCO₃ = 40 + 12 + 3·16 = 100 g/mol, so 36 g of water (18 g/mol) is 36/18 = 2 mol. Balanced chemical equations conserve mass and charge, and their coefficients give mole ratios that convert reactant amounts into product amounts. The limiting reactant is fully consumed first and caps how much product forms. Setting up a clean mole ratio is the heart of every stoichiometry problem.
Acids, bases, and pH
pH measures acidity as pH = −log₁₀[H⁺], where [H⁺] is the hydrogen-ion molarity. A strong acid such as HCl dissociates completely, so a 0.001 M solution has [H⁺] = 10⁻³ and pH = 3; likewise pOH = −log₁₀[OH⁻] and pH + pOH = 14 at 25°C. The scale is logarithmic, so each pH unit is a tenfold change in [H⁺]. Strong acids and bases dissociate fully while weak ones only partially; buffers resist pH change, and neutralization pairs an acid with a base to form water and a salt.
Oxidation states and redox
Oxidation states track electron ownership by rule: free elements are 0, oxygen is usually −2, hydrogen +1, and the states in a neutral molecule sum to zero. In H₂SO₄, hydrogen contributes 2(+1) and oxygen 4(−2) = −8, so sulfur must be +6. In a redox reaction, one species is oxidized (loses electrons, state increases) and another is reduced (gains electrons, state decreases) — remember 'OIL RIG': Oxidation Is Loss, Reduction Is Gain. Redox drives corrosion, batteries, and combustion; galvanic (cathodic) protection uses a sacrificial anode to shield the metal being protected.
Solutions and concentration
Molarity is moles of solute per liter of solution, M = n/V. Dissolving 0.5 mol NaCl to a total 0.5 L gives 1.0 M. Dilution conserves the amount of solute, so M₁V₁ = M₂V₂: diluting 100 mL of 6.0 M HCl to 1.0 M requires V₂ = (6.0·100)/1.0 = 600 mL total. Other measures include molality (per kilogram of solvent), mass percent, and parts per million (ppm), the last common in environmental work. Keeping volume in liters and tracking whether the question asks for final volume or added volume prevents slips.
The ideal-gas law and gas behavior
The ideal-gas law P·V = n·R·T links pressure, volume, moles, and absolute temperature; with R = 8.314 J/mol·K (or 0.08206 L·atm/mol·K), 1 mol in a 24.6 L vessel at 300 K exerts P = nRT/V = (1·8.314·300)/0.0246 ≈ 101 kPa. At standard temperature and pressure, one mole of ideal gas occupies about 22.4 L, so 2 mol occupy 44.8 L. Temperature must always be absolute (kelvin). The related gas laws (Boyle's, Charles's) are special cases holding one variable fixed.
Equilibrium, kinetics, and the periodic table
Reversible reactions reach equilibrium when forward and reverse rates match, described by an equilibrium constant K from the ratio of product to reactant activities. Le Chatelier's principle predicts the response to a disturbance: adding reactant, raising pressure on the side with more gas moles, or changing temperature shifts equilibrium to partially oppose the change. Reaction kinetics — how fast a reaction proceeds — depends on concentration, temperature, and catalysts, which lower the activation energy without being consumed. Periodic trends (electronegativity, atomic radius, ionization energy) explain reactivity and bonding across the table.
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