Thermodynamics

Thermodynamics explains energy changes during chemical reactions and why some reactions release heat (like burning fuel) while others absorb heat (like ice melting). Understanding thermodynamics answers fundamental questions: Why do some reactions happen spontaneously while others need a push? What determines whether a chemical process is efficient or wasteful?

System and Surroundings: Defining Boundaries

Thermodynamics always starts by defining a system—the specific substance(s) being studied—and the surroundings—everything else in the universe. When fuel burns in an engine, the fuel is the system; the engine and air around it are surroundings.

Three types of systems exist:

Open systems exchange both matter and energy with surroundings. A kettle of boiling water: steam escapes (matter leaves), heat transfers out (energy leaves).

Closed systems exchange only energy, not matter. A sealed container with liquid: no gas escapes, but heat can flow in or out through walls.

Isolated systems exchange nothing with surroundings. Perfectly insulated containers exchange neither matter nor energy. In practice, truly isolated systems are theoretical—all real systems exchange something eventually.

The distinction matters because different thermodynamic laws apply. Many chemistry problems assume systems are closed (ignoring escaped gases) for simplicity.

Internal Energy: The Energy Stored in a Substance

Every substance contains internal energy (U)—the sum of all kinetic and potential energy of its particles. Kinetic energy comes from molecular motion (heat). Potential energy comes from bonds and attractive forces between particles.

Internal energy is a state function: it depends only on the current condition (temperature, pressure, composition), not how that state was reached. If you heat water from 20°C to 50°C, the change in internal energy is the same whether you use a hot flame or a slow burner—only the starting and ending states matter.

This seems obvious but is crucial: we can't measure absolute internal energy, only changes in it (ΔU). A substance at room temperature has internal energy we can't calculate absolutely, but we can say: "Heating from 20°C to 30°C increases internal energy by 1000 J."

Heat and Work: How Energy Transfers

Energy transfer between system and surroundings occurs via two mechanisms:

Heat (q) is energy transferred due to temperature difference. When you touch hot water, heat flows from water to your hand. In cooking, heat from fire transfers to food.

Work (w) is organized energy transfer, often mechanical. When a gas expands against external pressure, it does work (pushes against surroundings). When surroundings compress a gas, they do work on it. In internal combustion engines, expanding hot gases do work on pistons.

Chemistry focuses on PV work: pressure-volume work. A gas expanding must push against atmospheric pressure—that requires energy and is "work done by the system" (negative w, energy leaves the system).

First Law of Thermodynamics: Energy Conservation

The First Law is elegant: ΔU = q + w

Internal energy change equals heat plus work. If a system absorbs 100 J of heat and has 30 J of work done on it, its internal energy increases by 130 J. If it releases 100 J of heat and does 30 J of work, its internal energy decreases by 130 J.

Sign convention matters:

• q is positive if heat enters the system (endothermic)
• q is negative if heat leaves the system (exothermic)
• w is positive if work is done on the system
• w is negative if the system does work on surroundings

This law is really stating conservation of energy in a thermodynamic context: energy isn't created or destroyed, only transferred between system and surroundings.

Enthalpy: Heat at Constant Pressure

Most chemistry happens at constant atmospheric pressure (not in rigid containers). Under these conditions, a new quantity becomes useful: enthalpy (H), defined as H = U + PV.

The change in enthalpy (ΔH = ΔU + Δ(PV)) approximately equals heat absorbed at constant pressure. This simplifies calculations—instead of tracking both internal energy change and PV work separately, we track enthalpy.

Exothermic reactions (ΔH negative) release heat: combustion, neutralization reactions, freezing. Endothermic reactions (ΔH positive) absorb heat: melting, evaporation, many decompositions.

Enthalpy is why we talk about "heat of reaction" colloquially. Burning one mole of methane releases 890 kJ of heat (ΔH° = -890 kJ/mol). This tells you: that reaction is highly exothermic—useful for heating.

Heat Capacity: How Much Heat per Temperature Change?

Different substances require different amounts of heat to change temperature. Heat capacity (C) is the amount of heat required to raise one unit's temperature by 1°C (or 1 K).

Specific heat capacity (c) is heat capacity per unit mass: how much heat per gram per degree. Water's specific heat is 4.18 J/(g·°C)—very high. Metals like copper are 0.39 J/(g·°C). This explains why water heats slowly but copper heats quickly: water's high heat capacity means you need much more energy to raise its temperature.

Heat = mass × specific heat × temperature change: q = m × c × ΔT. This equation lets you calculate how much heat is needed to warm something or how much a substance released if it cooled.

Thermochemical Equations and Hess's Law

Thermochemical equations show both reactants/products and the enthalpy change:

C(graphite) + O₂(g) → CO₂(g) ΔH° = -393.5 kJ/mol

The -393.5 kJ means burning one mole of solid carbon produces 393.5 kJ of heat.

Hess's Law states: if a reaction can be written as a sum of other reactions, its enthalpy change equals the sum of those reactions' enthalpy changes. This is powerful—if you can't measure a reaction directly, combine other reactions to calculate it.

For example, if you know:

• H₂(g) + ½O₂(g) → H₂O(g) ΔH° = -242 kJ
• C(s) + O₂(g) → CO₂(g) ΔH° = -394 kJ

You can calculate enthalpy for forming other compounds by adding/subtracting these equations. Hess's Law works because enthalpy is a state function—it depends only on initial and final states, not the path taken.

Entropy and Spontaneity: The Second Law

The First Law says energy is conserved but doesn't explain why some reactions go forward while others don't. The Second Law introduces entropy (S): a measure of molecular disorder or randomness.

Spontaneous processes (like ice melting, salt dissolving) increase total entropy. The universe naturally progresses toward disorder. A hot object cooling to room temperature increases entropy (order decreases, energy spreads out).

Not all spontaneous reactions are exothermic: some endothermic reactions occur spontaneously if entropy increase is large enough. This introduces Gibbs free energy (G = H - TS): a single quantity combining enthalpy and entropy to predict spontaneity. If ΔG is negative, the reaction is spontaneous.

Understanding these concepts explains: Why does ice spontaneously melt in sunlight (absorbs heat—endothermic!) yet spontaneously freezes in winter? Because ΔG (combining enthalpy and entropy) drives the process, not just ΔH.

Socratic Questions

1. If ice melting is endothermic (absorbs heat), why does ice melt in sunlight? How can an endothermic process be spontaneous?

1. Why does water evaporate slowly at room temperature even though no heat is added? Where does the energy come from if not from heat?

1. In Hess's Law, why does the path taken to form a product not matter for enthalpy change, only initial and final states? What property of enthalpy makes this true?

1. Why is heat released when opposite ions attract in an ionic bond, yet forming that bond is often spontaneous? Doesn't releasing heat suggest disorder decreases?

1. If entropy always increases (Second Law), why doesn't a cup of tea spontaneously heat itself up? Why does it always cool down?