Thermochemistry Mcat

(Note: This guide is part of our MCAT General Chemistry series.)

Thermochemistry for the MCAT: Everything You Need to Know

Learn key MCAT concepts about thermochemistry, plus practice questions and answers

(Note: This guide is part of our MCAT General Chemistry series.)

Table of Contents

Part 1: Introduction to thermochemistry

Part 2: Principles of thermochemistry

a) Laws of thermodynamics

b) Endothermic and exothermic reactions

c) Spontaneous and nonspontaneous reactions

d) Gibbs free energy

Part 3: Calorimetry

a) Heat transfer

b) Forms of heat transfer

Part 4: Phase change

a) Isothermal process

b) Phase change diagrams

c) Hess’ Law

Part 5: High-yield terms

Part 6: Passage-based questions and answers

Part 7: Standalone practice questions and answers

Part 1: Introduction to thermochemistry

Chemical reactions involve the cleavage and formation of chemical bonds. Physical reactions rearrange molecular interactions without changing the chemical identity of a substance—such as in the familiar phase change of freezing liquid water to form solid ice. Bond cleavage and formation, as well as intermolecular arrangements during phase change, are associated with the absorption and release of heat. Thermochemistry examines these quantitative changes in heat in the context of a variety of chemical reactions. In summary: thermochemistry is the quantitative study of heat that is released, absorbed, or evolved during chemical and physical reactions.

As a subdivision of thermodynamics, thermochemistry follows the zeroth, first, and second laws of thermodynamics. Thermodynamics studies how different forms of energy—such as mechanical and potential energy—are transferred. These overarching concepts will help you understand more about thermochemistry. For more information on this topic, be sure to refer to our guide on thermodynamics.

Throughout this guide, there are several important terms written in bold. At the end of this guide, there are also several MCAT-style practice questions for you to test your knowledge against.

Part 2: Principles of thermochemistry

a) Laws of thermodynamics

There are three laws of thermodynamics that are important for the MCAT. In the context of thermochemistry, a “system” refers to the molecules, bonds, and atoms involved in a chemical reaction. Its “environment” may refer to anything that is not involved in the reaction, including a solvent, the ambient air, or even the room the reaction is conducted in.

The zeroth law of thermodynamics states that if two thermodynamic systems are in thermal equilibrium with a third system, then all three systems are in thermal equilibrium with each other. Put another way, the zeroth law states that if systems A and B are in thermal equilibrium, and systems B and C are in thermal equilibrium, then systems A, B, and C are all in thermal equilibrium.

Even if the three systems were allowed to exchange energy, there would be no net heat exchange, and the three systems will possess the same temperature. Thus, the zeroth law introduces the concept of temperature as it relates to heat.

Note that the temperature and the heat possessed by an object are different quantities. Temperature is a measure of the average kinetic energy of a substance and is given by units Celsius (C), Kelvin (K), or Fahrenheit (F). Heat is the transfer of energy that results from a temperature difference between substances. Heat is given in the units of joules (J). Note that heat transfer is a prerequisite for a change in temperature.

Additionally, heat is a process function, while temperature is a state function. A state function results in a property whose value does not depend on the path taken to reach the final “state” at which the value is obtained. This is the opposite of a process function, in which the value of the property changes depending on the prior steps taken to achieve the final state.

The first law of thermodynamics states that energy is conserved. Energy is neither created nor destroyed; rather, it can only be transferred between individual objects and systems. For a closed system, a system that can exchange energy but not matter, the first law of thermodynamics can be stated as:

ΔU = Q-W

where ΔU = change in system’s internal energy,
Q = heat added to the system,
W = work done by the system

Thus, the internal energy of a system can be transferred into heat loss or gain, or into forms of work.

The second law of thermodynamics states that systems tend toward increasing entropy. This law introduces the concept of entropy as a measure of disorder. To better understand entropy, consider the relative entropies of gas, liquid, and solids. Gaseous molecules possess the highest entropy, followed by liquid, and then finally, solids.

Entropy can be described by the following equation:

ΔS = k * ln(W)

where ΔS = change in entropy,
k = Boltzmann’s constant,
ln(W) = natural logarithm of W

In this context, W does not represent work—but rather, the total number of possible microstates the system can adapt. A microstate refers to any combination of all possible orientations of particles within a system. Thus, since the number of microstates depends on the number of particles within a system, the entropy of a system tends to increase as more particles are added.

The Boltzmann constant, represented by k or k_B, is a fundamental constant used in several thermodynamic equations. While you may not need to memorize the exact quantity of this constant, it is useful to recognize that it is a defined constant.

b) Endothermic and exothermic reactions

Endothermic reactions are chemical reactions in which energy, often in the form of heat, is transferred from the environment into the system. As a result, the temperature of the system may rise while the temperature of the environment may drop.

Recall that the variable Q represents heat. As a result, ΔQ (final Q – initial Q) for an endothermic is positive. The change in enthalpy (ΔH) for an endothermic reaction is always positive (ΔH > 0). Enthalpy is a thermodynamic quantity that will be discussed later in this guide.

Exothermic reactions are chemical reactions in which heat is released from the system. As a result, the temperature of the system may drop while the temperature of the environment may rise. The ΔQ for an exothermic reaction is negative. The change in enthalpy (ΔH) is also negative (ΔH < 0).

What’s the difference between ΔQ and ΔH? Enthalpy is a state function, which means its value only depends on the starting and ending states of a process. Heat is a process function; its value depends on the particular process(es) taken to go from a starting to an ending state. Further, ΔH equals ΔQ when the reaction occurs under constant pressure.

c) Spontaneous and nonspontaneous reactions

The Gibbs free energy equation is given by:

ΔG = ΔH – TΔS

where ΔG is the change in Gibbs free energy,
ΔH is the change in enthalpy,
T is the temperature of the system,
ΔS is the change in entropy

Note that these thermodynamic properties are provided in terms of a relative value. In other words, the quantities of Gibbs free energy, enthalpy, and entropy will always be provided as an amount of change relative to some baseline value or previous measurement.

Thermochemistry Mcat

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Thermochemistry & Gases

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Heat, Entropy, and Free Energy

Entropy

Standard Gibbs Free Energy and Equilibrium

Heat and Heating Curves

Introduction to State Functions

Phases of Matter

Ideal Gases

Properties of Ideal Gases

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