Introduction

"It is the only physical theory of universal content concerning which I am convinced that, within the framework of the applicability of its basic concepts, it will never be overthrown." — Albert Einstein

Burning a fuel like methane, cooking gas, or coal releases the chemical energy stored inside its molecules as heat. That same chemical energy can just as easily be redirected — burned inside an engine, it becomes mechanical work; discharged through a galvanic cell like a dry battery, it becomes electrical energy. Different forms of energy are constantly being converted into one another, and thermodynamics is the study of exactly these transformations.

Importantly, thermodynamics is not concerned with how a transformation happens or how fast — it only looks at the system's initial and final states. That restricts its laws to systems that are already in equilibrium, or moving from one equilibrium state to another. (A system is "in equilibrium" when its macroscopic properties — pressure, temperature, and so on — stop changing with time.) This restriction is what makes thermodynamics powerful: instead of tracking the messy details of a reaction as it happens, it lets us predict outcomes just from where a system starts and where it ends up.

This chapter builds the machinery to answer three questions that recur throughout chemistry: How much energy does a given reaction release or absorb, and can we work that out in advance? What determines whether a reaction happens on its own at all? And once a reaction does proceed, how far does it go before stopping?

We start by making precise what we mean by a system, its surroundings, and the state it is in — the vocabulary every later argument depends on. From there we develop the first law of thermodynamics (energy is conserved), and use it to define two of the most useful quantities in all of chemistry: internal energy ($U$) and enthalpy ($H$). Enthalpy, in particular, is what lets us talk about the “heat of a reaction” in a way that can be measured, tabulated, and combined algebraically (Hess's law) to predict enthalpies we could never measure directly.

Knowing how much heat a reaction releases is only half the story, though. A reaction can be strongly exothermic and still never happen by itself — and some reactions that absorb heat happen spontaneously anyway (an ice cube melting on a warm day is one). The second half of this chapter builds the tools that answer the real question chemists care about: will this reaction happen on its own, and in which direction? That means introducing entropy ($S$), the measure of disorder that the second and third laws of thermodynamics are built around, and combining it with enthalpy into Gibbs energy ($G$) — the single quantity whose sign tells you whether a process is spontaneous at a given temperature. The chapter closes by connecting $\Delta G$ to the equilibrium constant $K$, linking thermodynamics directly to the chemical equilibrium you will study next.

By the end of this chapter you should be able to: classify a system as open, closed, or isolated; state and apply the first law of thermodynamics; calculate work and heat for simple gas processes; distinguish internal energy from enthalpy and calculate one from the other; use Hess's law and bond enthalpies to find unknown reaction enthalpies; and use $\Delta H$, $\Delta S$, and $\Delta G$ together to decide whether a reaction is spontaneous.