Thermochemistry Basics

Thermochemistry is the study of heat flow into and out of chemical reactions and physical changes. Three words get mixed up constantly, so pin them down first: - Heat (q) is energy that moves between two things because they're at different temperatures. It's a transfer, not a substance — you can't have "a heat," only "heat flowing." - Temperature is a measure of the *average kinetic energy* of the particles in a substance. It tells you how hot something is, not how much energy it holds. - Energy is the capacity to do work or transfer heat. Heat is one way energy moves; it's measured in joules (J) or kilojoules (kJ), same as any other energy. A cup of water and a swimming pool can be at the same temperature, but the pool holds far more total thermal energy because it has so much more mass. Temperature is intensive (doesn't depend on amount); total thermal energy is extensive (it does).

Every chemical or physical change either releases heat to the surroundings or absorbs heat from them. Exothermic ("exo-" = out) reactions release heat. The system loses energy, so the surroundings get warmer. Combustion is the classic example: CH₄ + 2O₂ → CO₂ + 2H₂O, ΔH = −890 kJ/mol Burning methane releases 890 kJ of heat per mole into the surroundings — that's why a stove burner feels hot. Endothermic ("endo-" = in) reactions absorb heat. The system gains energy, so the surroundings get colder. Photosynthesis is the classic example: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂, ΔH = +2803 kJ/mol A plant absorbs light energy to drive this reaction forward — without an outside energy source, it doesn't happen.

Enthalpy (H) is the heat content of a system at constant pressure — most reactions run in open beakers or flasks exposed to the atmosphere, so constant pressure is the normal lab condition. You can't measure H directly, but you can measure the change in enthalpy, ΔH, which equals the heat absorbed or released by a reaction at constant pressure: ΔH = H(products) − H(reactants) ΔH is written for the reaction as it's balanced, in kJ or kJ/mol, and its sign tells you the direction of heat flow for the system.

At constant pressure, the heat measured in a calorimetry experiment IS the enthalpy change: q = ΔH. That's why q and ΔH share the same sign convention above, and why the next section's q = mcΔT equation is really how ΔH gets measured in the lab.

Specific heat capacity (c), usually just called specific heat, is the amount of heat needed to raise the temperature of 1 gram of a substance by 1 °C. It's a property of the material — metals have low specific heat (they heat up and cool down fast), water has unusually high specific heat (it resists temperature change), which is why water is used to regulate temperature in everything from car radiators to your own body. The governing equation connects heat, mass, specific heat, and temperature change:

Worked example — heating water. How much heat is needed to raise 50.0 g of water from 20.0 °C to 80.0 °C? q = m c ΔT = (50.0 g)(4.184 J/g·°C)(80.0 °C − 20.0 °C) q = (50.0)(4.184)(60.0) = 12 552 J ≈ 12.6 kJ ΔT is positive (temperature rose), so q is positive — the water absorbed heat, consistent with something being heated.

Specific heat values for common substances, in J/(g·°C):

Calorimetry problems put two substances at different temperatures in contact and use conservation of energy: heat lost by the hotter substance equals heat gained by the cooler one, so q(lost) = −q(gained). Worked example. A 20.0 g piece of aluminum at 100.0 °C is dropped into 100.0 g of water at 20.0 °C. Assuming no heat escapes to the surroundings, what is the final temperature? Set q(Al) = −q(water): m(Al) c(Al) (Tf − 100.0) = −m(water) c(water) (Tf − 20.0) (20.0)(0.897)(Tf − 100.0) = −(100.0)(4.184)(Tf − 20.0) 17.94(Tf − 100.0) = −418.4(Tf − 20.0) 17.94Tf − 1794 = −418.4Tf + 8368 436.34Tf = 10 162 Tf ≈ 23.3 °C Notice the final temperature lands much closer to the water's starting temperature than the aluminum's — water's high specific heat and larger mass both dominate the outcome.

A heating curve plots temperature (y-axis) against heat added (x-axis) as a substance is heated steadily from solid through liquid to gas. It has five segments that alternate between sloped and flat:

During the two flat plateaus, energy is still being added, but it goes into breaking intermolecular forces (overcoming attractions between particles) rather than increasing kinetic energy — so temperature holds steady even though heating continues. The energy needed for these phase changes is called latent heat: the heat of fusion (ΔHfus) for melting/freezing, and the heat of vaporization (ΔHvap) for boiling/condensing. For water: ΔHfus = 334 J/g (6.01 kJ/mol) at 0 °C, and ΔHvap = 2260 J/g (40.7 kJ/mol) at 100 °C. Vaporization always takes far more energy than fusion for the same substance, because boiling completely separates particles while melting only loosens their fixed arrangement.

Using the values in the file (q = mcΔT, with ΔT = final temperature minus initial temperature), answer these. Round q to 3 significant figures. 1. Calculate q for trial 6. Is the process exothermic or endothermic for the substance? 2. Which trial released the greatest amount of heat? 3. Trials 1, 3, and 5 are all endothermic. Rank those three trials from most heat absorbed to least.