INTRODUCTION AND SECTION 5.1 Thermodynamics is the study of energy and its transformations. In this chapter we have focused on thermochemistry, the transformations of energy—especially heat—during chemical reactions.

An object can possess energy in two forms: (1) kinetic energy is the energy due to the motion of the object, and (2) potential energy is the energy that an object possesses by virtue of its position relative to other objects. An electron in motion near a proton, for example, has kinetic energy because of its motion and potential energy because of its electrostatic attraction to the proton. The SI unit of energy is the joule (J): 1 J = 1 kg-m 2/s2. Another common energy unit is the calorie (cal), which was originally defined as the quantity of energy necessary to increase the temperature of 1 g of water by 1 °C: 1 cal = 4.184 J.

When we study thermodynamic properties, we define a specific amount of matter as the system. Everything outside the system is the surroundings. When we study a chemical reaction, the system is generally the reactants and products. A closed system can exchange energy, but not matter, with the surroundings. Energy can be transferred between the system and the surroundings as work or heat. Work is the energy expended to move an object against a forceHeat is the energy that is transferred from a hotter object to a colder one. Energy is the capacity to do work or to transfer heat.

SECTION 5.2 The internal energy of a system is the sum of all the kinetic and potential energies of its component parts. The internal energy of a system can change because of energy transferred between the system and the surroundings. According to the first law of thermodynamics, the change in the internal energy of a system, ΔE, is the sum of the heat, q, transferred into or out of the system and the work, w, done on or by the system: ΔE = q + w. Both q and w have a sign that indicates the direction of energy transfer. When heat is transferred from the surroundings to the system, q > 0. Likewise, when the surroundings do work on the system, w > 0. In an endothermic process the system absorbs heat from the surroundings; in an exothermic process the system releases heat to the surroundings.

The internal energy, E, is a state function. The value of any state function depends only on the state or condition of the system and not on the details of how it came to be in that state. The heat, q, and the work, w, are not state functions; their values depend on the particular way by which a system changes its state.

SECTIONS 5.3 AND 5.4 When a gas is produced or consumed in a chemical reaction occurring at constant pressure, the system may perform pressure-volume (P-V) work against the prevailing pressure of the surroundings. For this reason, we define a new state function called enthalpy,H, which is related to energy: H = E + PV. In systems where only pressure-volume work is involved, the change in the enthalpy of a system, ΔH, equals the heat gained or lost by the system at constant pressure: ΔH = qp (the subscript p denotes constant pressure). For an endothermic process, ΔH > 0; for an exothermic process, ΔH < 0.

In a chemical process, the enthalpy of reaction is the enthalpy of the products minus the enthalpy of the reactants: ΔHrxn = H (products) – H (reactants). Enthalpies of reaction follow some simple rules: (1) The enthalpy of reaction is proportional to the amount of reactant that reacts. (2) Reversing a reaction changes the sign of ΔH. (3) The enthalpy of reaction depends on the physical states of the reactants and products.

SECTION 5.5 The amount of heat transferred between the system and the surroundings is measured experimentally by calorimetry. A calorimeter measures the temperature change accompanying a process. The temperature change of a calorimeter depends on its heat capacity, the amount of heat required to raise its temperature by 1 K. The heat capacity for one mole of a pure substance is called its molar heat capacity; for one gram of the substance, we use the term specific heat. Water has a very high specific heat, 4.18 J/g-K. The amount of heat, q, absorbed by a substance is the product of its specific heat (cs), its mass, and its temperature change: q = Cs × m × ΔT.

If a calorimetry experiment is carried out under a constant pressure, the heat transferred provides a direct measure of the enthalpy change of the reaction. Constant-volume calorimetry is carried out in a vessel of fixed volume called a bomb calorimeter. Bomb calorimeters are used to measure the heat evolved in combustion reactions. The heat transferred under constant-volume conditions is equal to ΔE. Corrections can be applied to ΔE values to yield enthalpies of combustion.

SECTION 5.6 Because enthalpy is a state function, ΔH depends only on the initial and final states of the system. Thus, the enthalpy change of a process is the same whether the process is carried out in one step or in a series of steps. Hess's law states that if a reaction is carried out in a series of steps, ΔH for the reaction will be equal to the sum of the enthalpy changes for the steps. We can therefore calculate ΔH for any process, as long as we can write the process as a series of steps for which ΔH is known.

SECTION 5.7 The enthalpy of formation, ΔHf, of a substance is the enthalpy change for the reaction in which the substance is formed from its constituent elements. The standard enthalpy change of a reaction, ΔH°, is the enthalpy change when all reactants and products are at 1 atm pressure and a specific temperature, usually 298 K (25 °C). Combining these ideas, the standard enthalpy of formation, Δf, of a substance is the change in enthalpy for the reaction that forms one mole of the substance from its elements in their most stable form with all reactants and products at 1 atm pressure and usually 298 K. For any element in its most stable state at 298 K and 1 atm pressure, ΔH°f = 0. The standard enthalpy change for any reaction can be readily calculated from the standard enthalpies of formation of the reactants and products in the reaction:

SECTION 5.8 The fuel value of a substance is the heat released when one gram of the substance is combusted. Different types of foods have different fuel values and differing abilities to be stored in the body. The most common fuels are hydrocarbons that are found as fossil fuels, such asnatural gaspetroleum, and coal. Coal is the most abundant fossil fuel, but the sulfur present in most coals causes air pollution. Renewable energy sources include solar energy, wind energy, biomass, and hydroelectric energy. Nuclear power does not utilize fossil fuels but does create controversial waste-disposal problems. The challenge of providing energy for the world has significant political and social implications in the areas of food supply and the environment.


• Interconvert energy units. (Section 5.1)

• Distinguish between the system and the surroundings in thermodynamics. (Section 5.1)

• State the first law of thermodynamics. (Section 5.2)

• Understand the concept of a state function and be able to give examples. (Section 5.2)

• Express the relationships among the quantities q, w, Δ E, and ΔH. Learn their sign conventions, including how the signs of q and ΔH relate to whether a process is exothermic or endothermic. (Sections 5.2 and 5.3)

• Use thermochemical equations to relate the amount of heat energy transferred in reactions at constant pressure (ΔH) to the amount of substance involved in the reaction. (Section 5.4)

• Calculate the heat transferred in a process from temperature measurements together with heat capacities or specific heats (calorimetry). (Section 5.5)

• Use Hess's law to determine enthalpy changes for reactions. (Section 5.6)

• Use standard enthalpies of formation to calculate ΔH° for reactions. (Section 5.7)



Kinetic energy


The change in internal energy


Relates the change in internal energy to heat and work (the first law of thermodynamics)


The work done by an expanding gas at constant pressure


Enthalpy change at constant pressure


Heat gained or lost based on specific heat, mass, and temperature change


Standard enthalpy change of a reaction