Thermodynamics and kinetics

Recognizing the type speeds product prediction: a hydrocarbon plus O₂ always combusts to CO₂ and water; an acid plus a base neutralizes to a salt and water; two ionic solutions may precipitate if a product is insoluble (use solubility rules). Many reactions belong to two types at once (a single-displacement reaction is also redox). The MCAT tests recognition and product prediction, not exhaustive classification.

Convert each given amount to moles, compare against the balanced mole ratio to find which reactant is limiting (produces the least product), and convert that product back to the requested unit. The other reactant is in excess. The theoretical yield is the limiting-reagent prediction; the actual yield is always less (side reactions, losses), so percent yield quantifies the gap. Rests entirely on the mole.

Specific heat c is the heat to raise one gram by one degree; water's is famously high (4.18 J/g·°C), so it resists temperature change. Heat capacity is the same idea for a whole object (per °C, not per gram) — don't conflate the two. Crossing a phase boundary takes latent heat (q = mL) at constant temperature, and the heat of vaporization is much larger than the heat of fusion. A heating curve is therefore sloped segments (q = mcΔT) joined by flat plateaus (q = mL).

The first law is conservation of energy for thermodynamic systems (heat into the system and work done by the system are the conventional positives). For a gas, expansion against pressure does work equal to the area under the P–V curve, so process type matters: isothermal (constant T, ΔU = 0), adiabatic (no heat, Q = 0), isobaric (constant P), isochoric (constant V, no work). Thermal expansion (ΔL = αLΔT) explains why materials grow when heated — bimetallic strips and expansion gaps in bridges.

Always use absolute temperature (kelvin). The combined gas law P₁V₁/T₁ = P₂V₂/T₂ handles "before/after" changes without n. The ideal gas law also gives gas density (ρ = PM/RT) and drives gas stoichiometry (volume ↔ moles via 22.4 L/mol at STP), connecting back to the mole. Choose the R whose units match the problem's pressure/energy.

In a gas mixture each component contributes a partial pressure proportional to its mole fraction — the principle behind alveolar gas exchange in 4B. Because average kinetic energy depends only on temperature, at a given T a lighter molecule has a higher speed (equal ½mv² ⇒ smaller m, larger v); Graham's law makes this quantitative for effusion (lighter gases effuse faster, ∝ 1/√M).

The ideal gas law assumes point particles with no attractions — valid at low pressure and high temperature, where molecules are far apart and fast. At high pressure the particles' own volume becomes significant, and at low temperature intermolecular attractions pull molecules together (lowering the real pressure below ideal). The van der Waals equation adds correction terms for both. Gases behave most ideally when small, nonpolar, hot, and dilute.

Adding a species shifts the system away from it; compressing a gas mixture shifts toward the side with fewer moles of gas. Only temperature changes K itself — concentration and pressure changes shift the position but leave K unchanged. For an exothermic reaction, raising temperature decreases K (shifts left); for endothermic, raising temperature increases K. A catalyst changes neither K nor the position — only the rate of reaching it.

When a reaction can give more than one product, which dominates depends on conditions. Under mild conditions the faster-forming (kinetic) product accumulates; given enough energy and time to reach equilibrium, the more stable (thermodynamic) product wins. This crystallizes the chapter's theme: the fastest path and the most favorable destination are not always the same.