What Are The Chemical Reactions Of Energy Changes?

Chemical reactions involve the rearrangement of atoms as chemical bonds between atoms are broken and formed. This rearrangement is accompanied by an energy change as energy is either released or absorbed. Understanding the energy changes involved in chemical reactions provides insight into how and why reactions occur.

The energy changes in reactions are important because the direction and extent to which a reaction proceeds depends on the energy changes. Reactions that release energy are thermodynamically favored and tend to proceed rapidly and completely, while reactions that absorb energy tend to proceed slowly and incompletely unless energy is input. Quantifying the energy changes allows prediction of reaction direction, calculation of equilibrium constants, and determination of reaction rates.

This article will provide an in-depth examination of the different types of energy changes in chemical reactions, how these can be measured and calculated, and applications in thermodynamics, reaction rates, and more.

Exothermic vs Endothermic Reactions

The key distinction between exothermic and endothermic reactions is whether energy is released or absorbed during the chemical reaction. In an exothermic reaction, energy is released to the surroundings in the form of heat during the reaction. In contrast, an endothermic reaction absorbs heat energy from the surroundings.

A simple example of an exothermic reaction is combustion – the burning of fuels like wood or gasoline. Combustion reactions involve the reaction of a fuel with oxygen, forming products like carbon dioxide and water. As the bonds in the fuel molecules are broken and new bonds form to create the products, energy is released in the form of heat and light. Other common exothermic reactions include metal oxidation reactions like rusting, neutralization reactions between acids and bases, and polymerization reactions that form large molecules from small monomers.

Endothermic reactions absorb heat energy from the surroundings. A classic example is photosynthesis, where plants use the energy from sunlight to convert carbon dioxide and water into glucose and oxygen. The absorption of heat helps the reaction proceed toward products as the plant creates energy-rich sugars. Other endothermic reactions include thermal decomposition processes, where heat is used to break down a larger molecule into smaller components, and many phase change processes like the melting of ice or evaporation of water require heat input.

Whether a reaction is exothermic or endothermic depends on the relative energy levels of the reactants versus the products. Exothermic reactions result in more stable, lower energy products, releasing the excess energy. Endothermic reactions require energy input to push the reaction toward less stable, higher energy products. Understanding the exothermic or endothermic nature of a reaction is key for scientists and engineers designing safe and efficient chemical processes and devices.

Measuring Energy Changes

There are several ways to measure the energy changes that occur during chemical reactions. One of the most common methods is through calorimetry. Calorimetry involves measuring the heat absorbed or released during a reaction. This allows you to quantify the enthalpy changes associated with the reaction.

A simple calorimeter consists of an insulated container with two chambers, one for the reaction and one for a reference material. By measuring the temperature change in the reference material, you can calculate the heat absorbed or released by the reaction. More advanced calorimeters allow for more precise measurements under different conditions of temperature and pressure.

Calorimetry requires a controlled environment and precise instrumentation to get accurate measurements. It provides quantitative data on a reaction’s enthalpy changes. This can be used to distinguish endothermic reactions that absorb heat from exothermic ones that release heat during the reaction.

Other methods like spectroscopy can also be used to measure energy changes during chemical reactions. Spectroscopic techniques look at changes in electromagnetic radiation from reactants and products to gather insight into bond breaking and forming. This provides additional thermodynamic data beyond just heat flow measurements.

Having multiple complementary techniques for measuring energy changes gives chemists a detailed understanding of the thermodynamics at play during chemical reactions. Calorimetry remains a foundational and widely used approach in directly quantifying the heating and cooling associated with reactions.


The laws of thermodynamics describe the relationships between thermal energy, heat, and work. They explain how energy can be transferred between a system and its surroundings.

The first law of thermodynamics states that energy cannot be created or destroyed in an isolated system. The total energy of the universe remains constant. This means that energy can only be transferred or transformed from one form to another. For example, chemical energy in a reaction can be converted into thermal energy and light.

The second law of thermodynamics states that the entropy, or disorder, of an isolated system will always increase over time. Entropy is a measure of randomness and possibilities. In a chemical reaction, entropy increases as energy spreads out and becomes more disordered. Exothermic reactions, which release energy, result in an increase in entropy. Endothermic reactions, which absorb energy, result in a decrease in entropy.

The third law of thermodynamics states that the entropy of a perfect crystal at absolute zero is exactly equal to zero. This provides a reference point for the measurement of entropy.

These laws demonstrate that energy changes inevitably occur during chemical reactions. Exothermic reactions release energy, causing an increase in entropy and disorder in the system. Endothermic reactions absorb energy, resulting in a decrease in entropy. The transfers and transformations of energy described by thermodynamics govern all chemical reactions involving energy changes.

Activation Energy

Activation energy is the minimum amount of energy needed for a chemical reaction to occur. Reactions require an initial input of energy to get started. This initial energy input is called the activation energy (Ea).

For a reaction to proceed, the reactant molecules must collide with enough energy to break their bonds and form new bonds to create the product molecules. The activation energy is like the energy barrier that must be overcome for the reaction to happen.

The higher the activation energy, the slower the reaction rate, as more energy is required for the reactant molecules to have effective collisions. Reactants with more energy than the activation energy will react, while those with less energy will just bounce off each other.

Catalysts work by providing an alternative reaction pathway with a lower activation energy. This speeds up the reaction rate as more reactant molecules now have enough energy to react.

The activation energy is a key concept in the kinetics and thermodynamics of chemical reactions. Understanding activation energies allows chemists to predict and control reaction rates.

Reaction Rates

The rate of a chemical reaction describes how quickly reactants are converted into products. Reaction rates depend heavily on energy changes during the reaction. Exothermic reactions, which release energy in the form of heat, tend to have faster reaction rates. This is because the release of heat provides more kinetic energy to the molecules, allowing them to collide more frequently. The extra energy helps overcome the activation energy barrier faster.

In contrast, endothermic reactions absorb heat energy. This cooling effect causes the molecules to move more slowly, decreasing collision frequency. More energy must be added to reach the activation energy, slowing down the rate. However, increasing temperature can provide this extra energy needed to speed up endothermic reactions.

Catalysts also work to lower activation energy and accelerate reaction rates. They provide an alternate pathway or mechanism for the reaction that has a smaller energy barrier. Less energy input is required for reactions to occur. Enzymes in the human body act as catalysts to enable necessary chemical reactions to proceed rapidly at body temperature.

Energy Diagrams

Reaction energy diagrams, also known as reaction profiles, illustrate how the total energy of a chemical system changes during the course of a chemical reaction. They provide a visual representation of the energy transitions that take place as reactants are converted into products.

On the diagram, the x-axis represents the progression of the reaction from reactants to products, while the y-axis represents the total energy of the system. The starting point on the left represents the energy of the reactants. As the reaction proceeds, energy must be absorbed to break bonds in the reactants. This is represented by an upward slope on the diagram and is known as the activation energy barrier. Once this peak is reached, the system can slide downhill into products, releasing energy as new bonds are formed. The final point on the right represents the total energy of the products.

The magnitude of the energy change between reactants and products indicates whether the reaction is exothermic (releases energy) or endothermic (absorbs energy). The activation energy also provides insight into reaction kinetics, as higher activation energies mean slower reaction rates. Overall, energy diagrams provide a clear visualization of the energetic landscape that must be traversed for a chemical reaction to occur. They are powerful tools for understanding reaction thermodynamics and kinetics.


Enthalpy is a thermodynamic property that is used to describe energy changes in chemical reactions and physical transformations. It refers to the total heat content of a system at constant pressure. Enthalpy provides insight into whether a chemical reaction is endothermic (absorbs heat) or exothermic (releases heat).

The enthalpy change (ΔH) of a reaction indicates the amount of heat absorbed or released when the reactants form products at constant pressure. An endothermic reaction has a positive ΔH, meaning it absorbs heat. An exothermic reaction has a negative ΔH, indicating that heat is released.

Enthalpy changes are usually measured in units of kilojoules per mole (kJ/mol). A larger absolute value for ΔH corresponds to a more energetic reaction. Exothermic reactions like combustion have very negative ΔH values, while endothermic reactions like photosynthesis have positive ΔH values.

By convention, the enthalpy change is reported as the difference between the products and reactants. So for the general reaction:
Reactants → Products
ΔH = Hproducts – Hreactants

Enthalpy changes provide valuable insights into the energy profile and favored direction of a chemical reaction. This makes enthalpy an important thermodynamic property for understanding energy transfers in chemistry.


Entropy is a measure of the disorder or randomness of a system. It relates to the number of ways the atoms or molecules of a system can be arranged. A highly ordered system has low entropy, while a disordered system has high entropy.

According to the second law of thermodynamics, entropy in an isolated system always increases over time as the system becomes more disordered. Chemical reactions that increase the entropy of the universe are more likely to occur spontaneously.

Gibbs free energy (G) is a thermodynamic function that combines entropy and enthalpy. The change in Gibbs free energy (ΔG) indicates whether a chemical reaction will occur spontaneously:

  • ΔG < 0 - Reaction proceeds spontaneously
  • ΔG > 0 – Reaction is non-spontaneous
  • ΔG = 0 – System is at equilibrium

Reactions with a positive entropy change (ΔS > 0) tend to be spontaneous, while those with a negative entropy change (ΔS < 0) tend to be non-spontaneous. Entropy favors disordered states, so spontaneous reactions tend to produce more disorder and randomness.

By analyzing the entropy change, chemists can determine if a reaction will occur spontaneously under certain conditions. Entropy is a key factor along with enthalpy change in predicting chemical reactivity.


Energy changes in chemical reactions have many applications in the real world. Here are some examples:

Combustion Reactions: The burning of fuels like wood, coal, natural gas and gasoline involve exothermic reactions that release energy in the form of heat. This energy can be harnessed to generate electricity, power vehicles, or heat homes. The combustion of rocket and jet fuels provides the thrust needed for aviation and space exploration.

Batteries: Batteries rely on reduction-oxidation reactions to generate electricity. In rechargeable batteries like lithium-ion batteries, the reactions are reversible allowing the batteries to be recharged. The energy released powers everything from small electronics to electric vehicles.

Photosynthesis: Plants utilize sunlight energy to power endothermic reactions that convert carbon dioxide and water into glucose and oxygen. This process provides energy for plants and also releases oxygen needed by other organisms. It is the basis for almost all life on Earth.

Metabolism: The complex series of chemical reactions that occur in cells rely on energy changes to power life processes. Catabolic reactions break down nutrients and release energy that is then captured to drive anabolic reactions that build cell components.

Cooking: Processes like boiling water, baking breads and roasting meats involve endothermic and exothermic reactions. Understanding reaction kinetics helps optimize cooking by controlling the rate of heat transfer and chemical changes.

In summary, energy changes in chemical reactions power many essential processes and innovations in our world. Controlling and harnessing these energy transfers has broad applications for technology, industry, transportation, food, medicine and more.

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