Unit 6: Thermochemistry and Thermodynamics
In this unit, we study thermodynamics and thermochemistry. Thermodynamics is the study of heat transfer. Thermochemistry is specifically the study of heat transfer in chemical reactions. We were introduced to thermodynamics in Unit 5 when we learned about the energy associated with phase changes. Thermodynamics and thermochemistry allow us to predict whether a reaction will produce heat, such as the burning of a candlewick, or if a reaction will require heat to proceed, such as the reaction that occurs inside a disposable cold pack. In this unit we also learn about Gibbs Free Energy, which tells us whether a reaction is spontaneous, meaning the reaction will occur without external "help".
Completing this unit should take you approximately 6 hours.
Upon successful completion of this unit, you will be able to:
- define temperature and heat, and state the units used for each;
- perform enthalpy change, enthalpy of reaction, enthalpy of combustion, and enthalpy of formation calculations;
- define entropy;
- describe the driving force of a chemical reaction and relate it to Gibbs free energy; and
- use Hess' law to solve thermodynamic problems.
6.1: Energy
Before we begin our study of thermodynamics and thermochemistry, we need some definitions.
Thermodynamics incorporates many everyday words, such as energy, heat, work, system, and surroundings, which have precise meanings in this scientific context. In thermodynamics, the system refers to whatever you are interested in learning about. For example, the system could be a chemical reaction, or a substance undergoing a phase change. The surroundings refer to everything the system comes into contact with. For example, if you are doing a chemical reaction in a beaker, the system is the reaction and the surroundings are the beaker and air directly surrounding the beaker.
There are two overall types of energy: kinetic and potential. Kinetic energy is the energy of motion. Any object in motion has kinetic energy. Potential energy refers to stored energy that can be due to an object's position, or the energy stored in chemical bonds. Heat refers to the disordered flow of energy between a system and its surroundings when there is a temperature difference. Work refers to the ordered flow of energy between a system and its surroundings when there is not a temperature difference.
Read this text, which offers more detailed definitions of these topics. Do not focus on the chart that shows different energy units. In chemistry, we almost exclusively use the joule as our energy unit. The other somewhat commonly used energy unit is the calorie.
Read this text, which provides more detail about chemical energy. Here, we learn that the total energy of a system is the sum of all of the kinetic energy and all of the potential energy of that system. The energy of a molecular system is complex because molecules are constantly in many different forms of motion contributing to kinetic energy and many different interactions contributing to potential energy. Therefore, we can only measure the change in energy of a process rather than absolute energy.
This text also introduces the term enthalpy (H), meaning energy at constant pressure. If a process produces heat, we call it an exothermic process. If a process uses heat, we call it an endothermic process. Pay attention to the last section of the text, which discusses the phase changes we learned about in Unit 5.
Watch this video, which goes into more detail about the difference between internal energy (U) and enthalpy (H). The presenter uses a pressure-volume graph to show why internal energy is defined at constant volume and enthalpy is defined at constant pressure.
Watch this video, which defines the heat of formation (△Hf) as the energy change involved in forming one mole of a compound from its elements in their natural form. The presenter tabulates the heat of formations for most common compounds, which can be used to calculate the enthalpy change of a reaction if you have a balanced chemical equation. You can use the sign of the enthalpy of reaction to determine if the reaction is endothermic (positive enthalpy) or exothermic (negative enthalpy).
Watch this video, which shows another worked example of how to use heats of formation to calculate the enthalpy of reaction. Note that many chemists often use the terms "heat" and "enthalpy" interchangeably.
Scientists use calorimetry to measure the change in enthalpy of reactions in a laboratory setting. In calorimetry, we conduct the reaction in an isolated setting and measure the temperature change. We can then use the equation
– where 
is heat, 
is mass, 
is specific heat of the substance, and 
is change in temperature – to find the heat of the reaction.There are two main types of calorimetry: constant pressure, or coffee cup calorimetry, and constant volume, or bomb calorimetry.
Read this text, which describes the two types of calorimetry, and shows worked examples of how to calculate heat of reaction from calorimetry data. Pay attention to the sign conventions here. In calorimetry, we directly measure the temperature change of the surroundings, not the system, and first calculate the heat of the surroundings. Then, we need to convert this to heat of the system.
Watch this video to see another worked example of a coffee cup calorimetry problem.
Enthalpy is a state function. This means that enthalpy is path independent: it does not matter how you determine enthalpy experimentally or through calculation, you always get the same answer. This means you can use whatever information you have about a reaction to solve for enthalpy change. This is called Hess' Law.
One way to apply Hess' Law is to break a complex reaction up into the sum of smaller reactions. If you know the enthalpies of the smaller reactions, you can simply sum them to determine the enthalpy of the larger reaction.
Watch this video to see how to apply Hess' law. Pay close attention to what happens when a small reaction is flipped or multiplied by a factor. Changing a reaction in any way changes the enthalpy of that reaction.
Read this short section which supplements the video you just watched on Hess's Law.
Another way to determine the enthalpy of reaction is by using tabulated bond dissociation energies. You can apply this in a similar manner to how we used standard enthalpies of formation. Using bond energy to calculate the enthalpy of reaction is not particularly accurate because energies are affected by each molecule's unique surroundings, such as intermolecular forces. For this reason, energies are only an average across many different reactions.
Watch this video, which shows a worked example of this type of problem.
6.2: Thermodynamics
Thermodynamics is the general study of heat and energy transfer. It is governed by the three laws we will examine in this section. Some chemists call the first law of thermodynamics the law of conservation of energy. It states that energy can be neither created nor destroyed, but it may change form. For example, imagine a campfire: the energy is stored in chemical bonds in the wood and is released as light and heat.
Read this text. Note the highlighted equation that defines the first law of thermodynamics. This equation shows that the change of internal energy of a system is the sum of heat and work. Pay attention to the sign convention box. This shows that when heat or work is transferred from system to surroundings, it has a negative sign. When heat or work is transferred from the surroundings to the system, it has a positive sign.
This text also introduces pressure-volume work, which refers to the work involved in changing the pressure or volume of a system. Problem Example 1 shows a typical pressure-volume work problem. Later parts of this text use calculus to derive useful equations. If you have not learned calculus, you can skip the calculus parts and still learn the important material. Focus on the worked problems and the problem examples.
The second law of thermodynamics involves a thermodynamic quantity we call entropy (S). Entropy is a measure of the disorder of a system, measured in joules (J). The second law of thermodynamics states that the entropy of the universe is always increasing.
One consequence of the second law of thermodynamics is that in any engine there will be some energy lost as heat that cannot be harnessed to do work. We observe this in our everyday lives. If you touch the hood of your car while the engine is running, the hood of the car will feel hot. This is because some of the energy from your car engine is lost as heat. Because of this, the second law of thermodynamics explains why a perpetual motion machine can never exist.
Read this text. The first section explains the difference between reversible and irreversible processes. A reversible process can be modeled as a series of tiny steps, while an irreversible process must be modeled as a single large change. The second section discusses the meaning of entropy, and what disorder means on a microscopic level.
Entropy is a state function, which means we can apply Hess' Law to it. Absolute entropies of most common substances are tabulated, allowing us to calculate the entropy of a reaction in the same way we can calculate enthalpy of reaction from standard enthalpies of formation.
Read this text, which introduces the second law of thermodynamics. Pay close attention to the green box, which shows entropy calculations for the process of water freezing. It shows that while the system (water becoming ice) decreases in entropy, the entropy change of the universe is still positive.
The third law of thermodynamics states that the entropy of a perfect crystal at absolute zero (0K) is zero. This means that all molecular motion ceases in a perfect crystal at absolute zero. Pay attention to the definition of the third law of thermodynamics.
Gibbs Free Energy uses enthalpy, temperature, and entropy to predict whether a reaction or process will happen spontaneously or not. Recall that a spontaneous process is one that happens on its own, without needing "help" to go. A non-spontaneous process needs constant "help" to occur. If the Gibbs Free Energy for a process is negative, the process is spontaneous. If the Gibbs Free Energy is positive, the process is non-spontaneous.
It is important to point out that spontaneous does not mean fast; in fact, many spontaneous reactions are extremely slow. For example, the reaction of diamond turning to graphite is a spontaneous reaction, but it is extremely slow. So, we do not need to worry about diamond jewelry turning into graphite!
Watch this video, which describes the Gibbs Free Energy equation and shows how enthalpy, temperature, and entropy work together to determine if a process is spontaneous.
Read this text which discusses free energy, the thermodynamic function which predicts the direction of a chemical reaction and the composition of the system at equilibrium.
This video shows a worked example of using the Gibbs Free Energy equation to determine the spontaneity of a chemical reaction.