ChemTalk

What is Thermochemistry?

Concept of Heat

Core Concepts

Thermochemistry focuses on understanding how heat is involved in chemical reactions. Important thermochemical laws guide the study of heat and its interactions with the world. In this article, you will learn about the central concepts in thermochemistry, some crucial formulas, and how to solve thermochemistry problems.

Topics Covered in Other Articles

What is Thermochemistry?

Thermochemistry is a branch of chemistry that studies the heat involved in various reactions. Heat is a form of energy. The energy from heat can drive a chemical reaction, be released from a reaction, or both! Thermochemistry attempts to understand and explain the transformations the energy in a reaction undergoes.

“Thermo” means relating to heat. And “chemistry” is the study of molecules and reactions. So thermochemistry is the study of heat and molecules and reactions or the study of heat in chemical systems.

Food chemistry is one branch of chemistry where thermochemistry is a common study. When our body breaks down sugars and fats, it creates energy. Thermochemistry can help explain this. It can also help explain how an ice cube cools down a warm glass of water or how coal can supply energy to run our power grids.

First Law of Thermodynamics

The first law of thermodynamics is also referred to as the law of conservation of energy. That means energy can not be created or destroyed, it just changes forms. In more precise terms, the first law of thermodynamics says that the total energy change in a system is related to the heat added to the system minus the work done by the system. The study of heat in this system would be a thermochemistry study.

Heat of Reaction Equation (aka Enthalpy of Reaction)

The heat of reaction, also known as the enthalpy of reaction, looks at the overall heat energy that occurs in a reaction. This involves the heat absorbed into the matter as a reaction occurs, endothermic, or the heat released as a reaction occurs, exothermic. (Description of Endothermic vs Exothermic Reactions)

    \begin{gather*} {\Delta H = \sum \Delta H_{\text{products}} - \sum \Delta H_{\text{reactants}}} \end{gather*}

  • \Delta H: enthalpy
  • \sum \Delta H_{\text{products}} : sum of the heat absorbed/released by the products
  • \sum \Delta H_{\text{reactants}} : sum of the heat absorbed/released by the reactants

Example:

Table of enthalpy values
Table of Enthalpy Values for Some Substances

Consider the following:

    \begin{align*} {NaOH& + HCl \rightarrow NaCl + H_{2}O} \\ {\Delta H &= ?} \\ {\Delta H &= \sum \Delta H_{\text{products}} - \sum \Delta H_{\text{reactants}}} \\ {\Delta H &= (-411.0 + -241.8) - (-426.7 + -92.3)} \\ {\Delta H &= -682.8 + 519 = -163.8\text{kJ/mol}} \end{align*}

Calorimetry

Calorimetry studies the energy changes that occur between a system and its surroundings. The system represents the substance that is undergoing a chemical reaction. While the surroundings on the hand, represent everything else that is interacting with the system. A calorimeter, a calibrated tool, measures these energy changes. A calorimeter isolates the reaction occurring inside from the wider environment. A typical at-home example includes using a Styrofoam cup, where the material reduces the heat interactions between the contents of the cup from the outside.

Calorimetry is one way to study thermochemistry
*q represents heat. q becomes negative in a exothermic process when heat leaves the system, therefore the temperature of the surrounding should be greater than the temperature of the system. q is positive in an endothermic process. For that, the temperature of the surroundings will decrease since the system is absorbing the heat from the surroundings.

In the understanding from the image above, the heat absorbed in an ideal perfect world will always be equivalent to the heat released. By the law of conservation of energy, the total heat exchange within a closed calorimeter is equal to 0.

In other words,

    \begin{gather*} {q_{\text{released}} (-q) = q_{\text{absorbed}} (q)} \\ { -q + q = 0} \end{gather*}

Important Thermochemistry Equations

Heat Capacity (C)

The heat capacity of any matter describes the amount of heat (q) it can absorb or release during a temperature change.

    \begin{gather*} {C = \frac{q}{\Delta T}} \end{gather*}

Specific Heat Capacity (c)

The specific heat capacity refers to the amount of heat required to raise the temperature of 1 gram of a matter by 1 degree.

    \begin{gather*} {c = \frac{q}{ m \Delta T}} \end{gather*}

The specific heat capacity of a substance determines how good of a heat sink the substance will be. For example, water has a high specific heat capacity. Even though the sun beats down on the ocean all day, the ocean water does not change temperature dramatically. If water had a low specific heat capacity, the ocean would get hot during the day with the sun’s energy, and then cool off dramatically overnight as that energy dissipated into the air. This property of water is key to life functioning on earth.

Thermochemistry Practice Problems

Problem 1

If a substance has a high heat capacity, does that mean it is more or less resistant to temperature change?

Problem 2

You combust 0.100 \text{mol} of ethanol in a calorimeter, which has a jacket of 1.00 \text{kg} of water. Ethanol combusts according to the following reaction equation:

    \begin{gather*} {H_{3}CCH_{2}OH (\text{l}) + 3O_{2} (\text{g}) \rightarrow 2CO_{2} (\text{g}) + 3H_{2}O (\text{g})} \end{gather*}

When you combust all of the ethanol, you observe the temperature of the water increase by 5.16 \degree \text{C}. Using the heats of formation of the other chemicals involved in the reaction and the specific heat of water ( 4.18 \text{J/gK}), calculate the heat of formation of ethanol. Assume all heat from the combustion was absorbed by the water.

Thermochemistry Practice Problem Solutions

1: More resistant to temperature change

2: \Delta H_{f, \text{ethanol}} = -277 \text{kJ/mol}

Further Reading

Thermochemistry: Heat and Enthalpy