Lesson Explainer: Energy Changes in Reactions Chemistry

In this explainer, we will learn how to identify types of energy and relate changes in energy to chemical bonding and chemical reactions.

Energy is a quantity that exists in many categories. We may easily recognize kinetic energy, the energy of an object in motion, but there are other categories of energy that are less recognizable because they are held by an object. This type of stored energy is broadly classified as potential energy, the energy held by an object due to its position relative to other objects. An important type of potential energy is chemical energy, which is very important when considering energy changes in chemical reactions. Some energy categories are listed in the table below.

There are many categories of energy, and energy of one category can be transferred to a different category. For example, a ball on the top of a hill has potential energy. As the ball rolls down the hill, the potential energy is transferred into kinetic energy (motion), thermal energy (heating of the ball, ground, and air due to friction), and sound energy. Throughout the interconversion of energy, the total amount of energy must remain the same. This is summarized by the law of conservation of energy, which states that the total amount of energy in the universe is assumed to be constant. Thus, energy can neither be created nor destroyed only transferred from one category to another.

Law: Law of Conservation of Energy

The Law of conservation of energy says that energy cannot be created or destroyed; it is only ever transferred from one category to another.

All of the types of energy can be measured with the unit joule (J). For reference, it takes 4.184 joules of energy to increase the temperature of one millilitre of water by one degree Celsius. We can also measure energy using kilojoules (kJ), calories (cal), or kilocalories (kcal or Cal). The unit joules can be converted into any of these units by using the following conversion factors: 1000=1,4.184=1,4184=1.JkJJcalJkcal

The types of energy of particular interest to chemists are chemical, thermal, and radiant energy; you may see radiant energy referred to in some sources as light energy. Chemical potential energy is often explained as the energy stored in a chemical bond. However, atoms that are not bonded together still have chemical potential energy. Chemical potential energy is due to the relative position of particles and their electrostatic attractions.

In a molecule of hydrogen (H2) there exists a combination of attractive and repulsive forces between the protons and electrons of the two atoms.

When the forces are balanced, the atoms will be at a particular distance from one another. At this distance, the hydrogen molecule has its lowest possible chemical potential energy. To separate the atoms, the force of attraction must be overcome by supplying the molecule with more energy, typically from the categories of thermal energy or radiant energy. When heating or using the radiant energy, in for example the form of UV light, to separate atoms, additional energy is transferred into chemical potential energy. Thus, when the atoms are separated, they have a higher chemical potential energy than when they are bonded.

If the atoms recombine to form a bond, the chemical potential energy will decrease. The additional chemical potential energy must be transferred into other categories of energy through heat, light, or sound. Therefore, bond breaking always requires energy and bond forming always releases energy.

Atoms, molecules, and particles that are not at absolute zero exhibit motion. This means that each has kinetic energy. The thermal energy of a substance is the sum of all the kinetic energies of the particles in the substance. When we measure the temperature of a substance, we are measuring the average kinetic energy of the particles. It is important to recognize that the thermal energy of a substance does not indicate how hot or cold the substance is, nor does the temperature indicate the total thermal energy of a substance.

For example, a cup of boiling water has a higher temperature than a swimming pool because, on average, the molecules in the cup are moving faster than the molecules in the swimming pool. However, the swimming pool has a greater quantity of thermal energy. This is because the swimming pool contains more molecules. Thus, the sum of the kinetic energies in the swimming pool is higher than that of the cup of boiling water.

The term heat is often used interchangeably with thermal energy; however, heat specifically refers to energy that has been transferred from a substance with higher average kinetic energy (higher temperature) to a substance with a lower average kinetic energy (lower temperature). An object with a greater temperature will heat an object with a lesser temperature until the two substances have the same temperature.

Radiant energy is the energy of photons traveling in waves. X-rays, ultraviolet radiation, microwaves, and gamma radiation all have radiant energy. When the wavelength of the photon path is between 400 and 700 nanometres, we can see the radiant energy in the form of visible light.

Example 1: Identifying the Types of Energy Released by a Reaction

Diphenyl oxalate, hydrogen peroxide, and rhodamine A are mixed together in a plastic tube. After a few seconds, the plastic tube begins to feel warm, and the resulting mixture starts to glow a bright-red color. Which types of energy are produced by the mixture?

Answer

The tube feels warm because thermal energy is being released. This is commonly referred to as heat. We see a bright-red glow as visible electromagnetic radiation or radiant energy is being released. This is commonly referred to as light. The types of energy released by the mixture are heat and light.

To help us keep track of the energy conversions and categories of energy in a reaction, it is useful to define a system and its surroundings. A system is a specific part of the universe that we wish to observe, while the surroundings include anything that is not a part of the system.

Definition: System

A system is the specific part of the universe we wish to observe.

Definition: Surroundings

Surroundings include anything that is not a part of the system.

In the example below, the solution is defined as the system. This means that the beaker, the air, and the surface the beaker sits on are all part of the surroundings.

When performing a chemical reaction, we may wish to consider all species involved in the reaction as the system and everything else as the surroundings.

Systems may be open, closed, or isolated. An open system allows for the transfer of matter and energy with the surroundings. An example of an open system is a beaker.

Definition: Open System

An open system is a system that can exchange matter and energy with the surroundings.

A closed system only allows for the transfer of energy, not matter. An example of a closed system is an Erlenmeyer flask with a stopper.

Definition: Closed System:

A closed system is a system that does not exchange matter but can exchange energy with its surroundings.

An isolated system does not allow for the transfer of matter or energy. An insulated container is an approximation of an isolated system; however, truly isolated systems are only theoretical.

Definition: Isolated System

An isolated system is a system that does not allow for the transfer of matter or energy with its surroundings.

When we perform a chemical reaction, energy may be gained and/or lost by the system through heat, light, and/or sound. As the total amount of energy in the universe must remain constant, the amount of energy lost by the system must be gained by the surroundings and vice versa.

Example 2: Identifying the Space That Exchanges Heat Energy with a System

The space in which a chemical reaction takes place is known as the system. What is the name of the space where heat energy is absorbed from or released into?

Answer

Energy cannot be created nor destroyed. When we say a system gains heat energy, this energy must have come from another source. Likewise, when a system loses energy, the energy is not destroyed but is transferred to another space. We call this other space the surroundings.

Below is the reaction of methane gas (CH4) and oxygen: CH()+2O()CO()+2HO()4222gggg

Methane and oxygen have chemical potential energy due to the relative position of the atoms in the molecules and thermal energy due to the random motion of the molecules. The sum of the chemical potential energy and thermal energy of the methane and oxygen is the total internal energy of the system. This is called enthalpy (𝐻).

Definition: Enthalpy (𝐻)

Enthalpy (𝐻) is the total internal energy of a system.

Separating the atoms in a molecule of methane or oxygen requires energy. This energy can be supplied by heating the system. Some of the thermal energy supplied is converted into chemical potential energy as the bonds are broken.

Once the atoms have separated, they quickly recombine to form the products that have a lower chemical potential energy than the separated atoms. A decrease in the chemical potential energy of the system directly corresponds to an increase in the energy of the surroundings, most often as heat and/or light.

When we monitor a chemical reaction, we can neither measure the energy required to break the bonds nor the energy released when bonds are formed. However, we can measure the net change in the enthalpy (Δ𝐻) of the reactants and products by measuring the amount of heat absorbed or released by the reaction.

Example 3: Describing the Conservation of Energy in a Chemical Reaction

Which of these statements does not describe the conservation of energy in a chemical reaction?

  1. Energy is neither created nor destroyed during a chemical reaction.
  2. The energy contained in the bonds of reactant molecules always equals the energy contained in the bonds of product molecules.
  3. If the energy of a system increases, then the energy of the surroundings decreases by the exact same amount.
  4. Energy can only be transferred from one form to another.
  5. If the energy of a system decreases, then the energy of the surroundings increases by the exact same amount.

Answer

According to the law of conservation of energy, the total amount of energy in the universe is constant. This means that energy cannot be created nor destroyed. However, energy can be transferred from one category of energy to another. Thus, answers A and D are true statements.

To monitor energy transfer, we can define a system and its surroundings. A system is a particular part of the universe that we wish to monitor and the surrounding constitute the rest of the universe.

When performing a chemical reaction, we define the system as the species involved in the reaction. If the reaction releases energy, the system loses energy to the surroundings. If the reaction absorbs energy, the system gains energy from the surroundings.

As energy can neither be created nor destroyed, the amount of energy gained by a system must be equal to the amount of energy lost by the surroundings and vice versa. Thus, answers C and E are true statements.

This means that answer B must be false. The energy contained in the bonds of the reactants does not have to equal the energy contained in the bonds of the products. The energy difference represents the energy gained or lost by the system during the reaction. The statement that does not describe the conservation of energy in a chemical reaction is the statement in answer B.

When the products have a lower enthalpy than the reactants, the sign of Δ𝐻 will be negative, indicating that the overall reaction results in energy being released from the system to the surroundings: CH()+2O()CO()+2HO()kJmol4222ggggΔ𝐻=891/

Reactions that result in a net release of energy are exothermic.

Definition: Exothermic Process

An exothermic process is a process that releases energy to its surroundings.

As energy is released by the reaction, the change in enthalpy may be included in the chemical equation as a product: CH()+2O()CO()+2HO()+kJ4222gggg891

When the products have a higher enthalpy than the reactants, the sign of Δ𝐻 will be positive, indicating that the reaction requires a net increase in energy: 2NH()N()+3H()kJmol322gggΔ𝐻=+92/

Reactions that require a net increase in energy are endothermic.

Definition: Endothermic Process

An endothermic process is a process that absorbs energy from its surroundings.

As energy is absorbed by the reaction, the change in enthalpy may be included in the chemical equation as a reactant: 92kJ+2NH()N()+3H()322ggg

Example 4: Understanding a Thermochemical Equation

The chemical equation for the decomposition of magnesium carbonate is as follows: MgCO()kJMgO()+CO()32ssg+117

Based on this chemical equation, which of the following statements is true?

  1. 117 kJ of energy is released when a single molecule of MgCO3 decomposes.
  2. 117 kJ of energy is needed to make 1 mole of MgCO3 decompose.
  3. 117 kJ of energy is needed to make 42 g of MgCO3 decompose.
  4. 117 kJ of energy is needed to make a single molecule of MgCO3 decompose.
  5. 117 kJ of energy is released when 1 mole of MgCO3 decomposes.

Answer

The 117 kJ shown in the equation represents the net change in energy over the course of the reaction. As the energy is shown as a reactant, we know that 117 kJ must be supplied to the reaction. We can therefore eliminate answers A and E.

The amount of energy shown is particular to the number of moles of reactants and products in the chemical equation. We could read this chemical equation as follows: one mole of magnesium carbonate decomposes into one mole of magnesium oxide and one mole carbon dioxide when 117 kJ of energy is added to the reaction. Therefore, the correct answer is B.

Example 5: Calculating the Amount of Energy Released by the Combustion of Glucose

The chemical equation for the combustion of glucose is shown: CHO()+6O()6CO()+6HO()kJ6126222sggl+2808

How much energy is released when 2.5 moles of glucose are burned?

Answer

The chemical formula of glucose is CHO6126. We can see from the equation that one mole of glucose reacts with six moles of oxygen to produce six moles of carbon dioxide and six moles of water. This process releases 2‎ ‎808 kJ of energy. Thus, we can say 2‎ ‎808 kJ of energy is produced per mole of glucose. To determine the amount of energy released by 2.5 moles of glucose, we can perform the following calculation: 28081×2.5=7020.kJmolCHOmolCHOkJ61266126

The amount of energy released when 2.5 moles of glucose are burned is 7‎ ‎020 kJ.

We can often find tables in textbooks that list the change in enthalpy for various substances. For example, we might find the following entries:

Compound(state)Δ𝐻 (kJ/mol)
H()2g0
HO()2g242
O()2g0

We know that Δ𝐻 indicates the change in enthalpy over the course of a reaction. The plimsoll symbol, , indicates that the reaction was performed under standard conditions. Standard pressure is 1 bar, sometimes given as 1 atm, and standard concentration of solutes is 1 mol/dm3. In addition, temperature is generally stated to be 25C (298.15 K).

It is important to indicate the conditions under which the reaction is performed because pressure, concentration, temperature, and state of matter can all affect the change in enthalpy. The subscript f stands for formation. This means that the change in enthalpy is for the formation of the substance from its constituent atoms.

It is also important to note that the enthalpy is given in kilojoules per mole. This means that one mole of water produced by the reaction of hydrogen and oxygen, will release 242 kJ of energy. This corresponds to the reaction equation H()+O()HO()222ggg12 where only one mole of water is produced. However, we typically do not write chemical equations with fractions and are much more likely to see the following: 2H()+O()2HO()222ggg

As two moles of water appear in this equation, we can expect that the energy released will be twice as great: 2×242=484.kJkJ

We also see in the table that hydrogen and oxygen have a standard enthalpy of formation of zero. This is because at 1 bar and 25C, hydrogen already exists as H2 and oxygen already exists as O2. These molecules have enthalpy, but as they already exist in the desired form, no change in enthalpy will occur. The natural state of an element at 1 bar and 25C is known as the standard state. For elements like carbon that have several allotropes at 1 bar and 25C, one allotrope is chosen as the standard state.

Key Points

  • Energy cannot be created nor destroyed only transferred between different categories.
  • Energy lost by a system must be gained by the surroundings and vice versa.
  • Breaking bonds requires an increase in energy while forming bonds results in a release of energy.
  • The difference in the energy of the reactants and products in a reaction is the change in enthalpy.
  • Standard conditions denoted by are 1 bar, 1 M, and usually 25C.
  • Standard conditions also imply that elements are in their natural state or standard state at 1 bar and 25C.

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