Lesson Explainer: Bond Energy Chemistry

In this explainer, we will learn how to describe bond energies and use net changes in bond energy during reactions to calculate the bond energies of reactants and products.

Let us consider a molecule of hydrogen gas.

The molecule consists of two covalently bonded hydrogen atoms. The covalent bond can be broken to separate the hydrogen atoms: H()H()+H()2ggg

This process requires energy. The average amount of energy required to break one hydrogen–hydrogen single bond is 7.17×10 joules. This value is incredibly small and not particularly useful as reactions involve many millions of molecules. It is much more useful to know the energy required to break one mole of hydrogen–hydrogen bonds by using Avogadro’s number: 7.17×101×6.022×101=432000/.JbondbondsmolJmol

This value can be modified even further by converting joules into kilojoules with the conversion factor: 1000=1,432000/×11000=432/.JkJJmolkJJkJmol

Thus, breaking one mole of hydrogen–hydrogen single bonds will require 432 kJ of energy. The average amount of energy required to break a particular bond in one mole of gaseous molecules is called the bond energy (BE) or bond enthalpy and commonly has the unit kJ/mol.

Definition: Bond Energy (BE)

Bond energy is the average amount of energy required to break a particular bond in one mole of gaseous particles. This is also referred to as bond enthalpy.

Breaking bonds is an endothermic process. This means that energy needs to be supplied by the surroundings in order for the bonds to be broken.

Definition: Endothermic Process

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

As bond breaking is an endothermic process, bond energies will always have a positive value indicating that the energy must be absorbed by the system from the surroundings. Average bond energies are often listed in tables. We can use these values as an approximation of the amount of energy required to break one mole of bonds, but we should recognize that the actual value can vary slightly depending on other atoms in the molecule.

BondAverage Bond Energy (kJ/mol)BondAverage Bond Energy (kJ/mol)
HH432CN887
HN386CF485
HO459CCl327
HF656CBr285
HCl428NN167
HBr362NN418
CH411NN942
CC346NO466
CC602NO626
CC835OO142
CO358OO494
CO799FF155
CO1‎ ‎072ClCl240
CN305II148
CN615

Example 1: Definition of Hydrogen Molecules Bond Energy

Hydrogen molecules (H2) have a single HH bond. This bond has an energy of 432 kJ/mol. What does this mean?

  1. 432 kJ of energy is required to break one mole of HH bonds.
  2. 432 kJ of energy is required to break each HH bond.
  3. 432 kJ of energy is released when one mole of HH bonds is broken.
  4. 432 kJ of energy is released when one HH bond is broken.
  5. 432 kJ of energy is required to form one mole of HH bonds.

Answer

Bond energy or bond enthalpy is the amount of energy required to break a particular type of chemical bond. However, the amount of energy necessary to break one bond is very small. Therefore, bond energy is reported in kilojoules per mole to indicate the amount of energy required to break one mole of a particular type of bond. In this example, 432 kJ is necessary to break one mole of hydrogen–hydrogen single bonds. The correct answer is choice A.

The strength of a particular bond is quantified by the bond energy. In general, triple bonds are stronger than double bonds which in turn are stronger than single bonds when comparing like-bonded atoms. Thus, the bond energy required to break a carbon–carbon triple bond is greater than the bond energy required to break a carbon–carbon double bond.

Notice that the bond energy required to break a carbon–carbon double bond is not twice as much as the bond energy required to break a carbon–carbon single bond. Bond energy does not directly scale with the number of bonds.

In addition, when comparing bonds of the same type that involve elements from the same group of the periodic table, the bond strength will tend to decrease as we move down a group.

It is important to recognize that bond energies indicate the strength of a particular type of intramolecular force (covalent or ionic bonds) in a molecule or formula unit. Bond energy does not represent the strength of the intermolecular forces between two molecules. As such, bond energies cannot be used to explain properties like melting point and boiling point.

Example 2: Recognizing Bond Enthalpy Trends

Which of the following hydrogen halides has the smallest bond enthalpy?

  1. HI
  2. HF
  3. HBr
  4. HCl

Answer

Bond enthalpy or bond energy is the amount of energy required to break one mole of a particular type of bond. Each of these hydrogen halides consists of one hydrogen atom single bonded to a halogen atom. When an atom participates in the same type of bond with atoms from the same group, the bond enthalpy will tend to decrease as the size of the atom increases (moving down a group).

Looking at the periodic table, we can see that fluorine is at the top of the halogen group. This means that an atom of fluorine is the smallest of the halogen atoms, and the bond enthalpy of the hydrogen–fluorine bond should be the largest of the hydrogen halides. Moving down the group, we see chlorine, bromine, and iodine. Iodine is near the bottom of the halogen group. This means that an atom of iodine is larger than an atom of fluorine, chlorine, or bromine. This also means that the bond enthalpy of the hydrogen–iodine bond should be smaller than that of HF, HCl, or HBr. The hydrogen halide with the smallest bond enthalpy is choice A.

Formation of a bond releases the same amount of energy as was required for the bond to break. For example, if 432 kJ is required to break one mole of hydrogen–hydrogen bonds, this means that 432 kJ will be released when one mole of hydrogen–hydrogen bonds is formed: H()H()+H()kJmolH()+H()H()kJmol22gggggg+432/432/

Bond formation is an exothermic process. This means that energy is released to the surroundings.

Definition: Exothermic Process

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

During the course of a chemical reaction, bonds in the reactants may be broken and new bonds may be formed to create the products. The total amount of energy needed to break the bonds and the total amount of energy released when new bonds are formed will likely not be the same. The difference between the total energy absorbed and the total energy released is the change in enthalpy of a reaction (Δ𝐻). This is represented by the following equation: Δ𝐻=,BEBE()()bondsbrokenbondsformed where Δ𝐻 is the change in enthalpy and BE is the sum of the bond energies.

The change in enthalpy of a reaction may be represented by Δ𝐻, Δ𝐻, Δ𝐻rxn, or Δ𝐻rxn. When the sign of Δ𝐻 is positive (+), the overall reaction is endothermic. The reaction will absorb energy from the surroundings and the reaction vessel may begin to feel cold. When the sign of Δ𝐻 is negative (), the overall reaction is exothermic. The reaction will release energy to the surroundings and the reaction vessel may feel hot.

Let us consider the following reaction: 2H()+O()2HO()222ggg

For the reaction to proceed, two hydrogen–hydrogen single bonds and one oxygen–oxygen double bond must be broken, and four oxygen–hydrogen single bonds must be formed.

We can calculate the overall change in enthalpy of the reaction if we know the bond energy for each type of bond.

BondBond Energy (kJ/mol)
HH432
OO494
OH459

We know that two hydrogen–hydrogen single bonds and one oxygen–oxygen double bond were broken. We can substitute the bond energies into the change in enthalpy equation: Δ𝐻=Δ𝐻=2×+Δ𝐻=((2×432/)+494/).BEBEBEBEBEkJmolkJmolBE()()()()bondsbrokenbondsformedHHOObondsformedbondsformed

We also know that four oxygen–hydrogen single bonds are formed. We can substitute the bond energy into the equation Δ𝐻=((2×432/)+494/)4×Δ𝐻=((2×432/)+494/)(4×459/)kJmolkJmolBEkJmolkJmolkJmolOH and solve for the change in enthalpy of the reaction: Δ𝐻=478/.kJmol

The negative sign indicates that energy is released to the surroundings and that the reaction is exothermic.

Example 3: Bond Energy Changes in the Water–Gas Shift Reaction

The water–gas shift reaction is a major source of hydrogen gas for industrial processes. In this reaction, carbon monoxide is reacted with steam to produce hydrogen and carbon dioxide: CO+HOCO+H222

The energies of selected bonds are listed in the table. Calculate the total change in bond energy for this reaction, per mole of hydrogen gas produced.

BondHHCHOHCOCOCO
Bond Energy (kJ/mol)4324114593587991‎ ‎072

Answer

First, we should ensure that the chemical equation is balanced. After verifying that there is one carbon atom, two oxygen atoms, and two hydrogen atoms on both sides of the reaction, we should then determine the chemical structure of each species in the reaction. Carbon monoxide (CO) contains one carbon–oxygen triple bond. Water (HO2) has two hydrogen–oxygen single bonds. Carbon dioxide (CO2) contains two carbon–oxygen double bonds. Hydrogen gas (H2) has one hydrogen–hydrogen single bond.

To calculate the total change in bond energy or the change in the enthalpy of the reaction (Δ𝐻), we can use the following equation: Δ𝐻=,BEBE()()bondsbrokenbondsformed where Δ𝐻 is the change in the enthalpy of the reaction and BE is the sum of the bond energies. The reactants contain bonds that will be broken and the products contain bonds that will be formed. In this reaction, one CO bond and two OH bonds will be broken: Δ𝐻=+2×.BEBEBECOOHbondsformed()

Two CO bonds and one HH bond will be formed: Δ𝐻=+2×2×+.BEBEBEBECOOHCOHH

We can substitute the appropriate bond energies into the equation Δ𝐻=(1072/+(2×459/))((2×799/)+432/)kJmolkJmolkJmolkJmol and solve for the change in the enthalpy of the reaction: Δ𝐻=40/.kJmol

The negative sign indicates that the overall reaction is exothermic. The question asks for the total change in bond energy per mole of hydrogen gas. Looking at the balanced chemical equation, we can see that only one mole of hydrogen gas is produced via this reaction. Thus, the 40 kJ/mol is the total change in bond energy for this reaction, per mole of hydrogen gas.

We can also use the bond energies of select bonds and the change in enthalpy of the reaction to determine an unknown bond energy.

Example 4: Bond Energies of Chloromethane

Chloromethane is produced by the reaction of methane (CH4) with chlorine gas in the presence of UV light. The equation for this reaction is shown:

The reaction of 1.00 mol of methane releases 104 kJ of energy. The energies of selected bonds in the reactants and products are given in the table.

BondHClClClCH
Energy (kJ/mol)428240411

Calculate, to the nearest kilojoule per mole (kJ/mol), the energy of the CCl bond.

Answer

First, we should ensure that the chemical equation is balanced. After verifying that there is one carbon atom, two chlorine atoms, and four hydrogen atoms on both sides of the reaction, we can continue with the problem.

To calculate the bond energy of a CCl bond, we can use the change in enthalpy equation: Δ𝐻=,BEBE()()bondsbrokenbondsformed where Δ𝐻 is the change in enthalpy of the reaction and BE is the sum of the bond energies.

We know that the reaction releases 104 kJ of energy for one mole of methane. We can rewrite this information as 104 kJ/mol. The negative sign indicates that the energy is released. This value is the change in enthalpy of the reaction (Δ𝐻) and is substituted into the equation: 104/=.kJmolBEBE()()bondsbrokenbondsformed

We can see from the reaction equation that one carbon–hydrogen single bond and one chlorine–chlorine single bond must be broken. One carbon–chlorine single bond and one hydrogen–chlorine single bond must be formed in this reaction.

Adding this information to the equation gives us 104/=++.kJmolBEBEBEBECHClClCClHCl

We can substitute the known bond energies into the equation 104/=(411/+240/)+428/kJmolkJmolkJmolBEkJmolCCl and rearrange to solve for the CCl bond energy: 104/=651/+428/755/=+428/755/=428/327/=327/=.kJmolkJmolBEkJmolkJmolBEkJmolkJmolBEkJmolkJmolBEkJmolBECClCClCClCClCCl

The bond energy of the CCl bond is 327 kJ/mol.

A reaction profile diagram can be used to represent or determine the change in enthalpy of a reaction. Below is a reaction profile diagram for the reaction of hydrogen and oxygen to produce water: 2H()+O()2HO()222ggg

In a reaction profile diagram, the reactants are shown on the left-hand side and the products are shown on the right-hand side. In between the reactants and products, we see the separated atoms at the peak of the graph. The peak represents an intermediate state that exists after bonds in the reactants have been broken but before bonds in the products have been formed. The difference in the energy of the reactants and intermediate is the total bond energy of the broken bonds. The difference in the energy of the intermediate and products is the total bond energy of the formed bonds. The difference in energy of the reactants and products is the change in enthalpy.

In an exothermic reaction, the energy of the products will be less than the energy of the reactants. In an endothermic reaction, represented by the reaction profile diagram below, the energy of the products will be greater than the energy of the reactants: 2NH()N()+3H()322ggg

Example 5: Interpreting a Reaction Profile Diagram

In which of these reaction profile diagrams is the energy required to break chemical bonds the greatest?

Answer

In a reaction profile diagram, the leftmost line represents the reactants and the rightmost line represents the products. The peak of the graph represents an intermediate state after bonds in the reactants have been broken but before bonds in the products have been formed.

The difference in the energy of the reactants and intermediate represents the energy required to break the bonds involved in the reaction. The difference in the energy of the intermediate and products represents the energy released when the new bonds are formed. The difference in the energy of the reactants and products is the change in enthalpy of the reaction (Δ𝐻).

We need to identify the graph in which the energy required to break chemical bonds is the greatest. This graph will have the greatest difference in reactant energy and intermediate energy. In all four graphs, the reactants have the same starting energy, but graph a has the highest intermediate peak.

This means that graph a will have the greatest difference in reactant and intermediate energies. The graph with the greatest energy required to break chemical bonds is graph a.

Key Points

  • The energy required to break one mole of a particular type of bond is known as the bond energy (BE) and has the unit kJ/mol.
  • Breaking bonds is an endothermic process, while forming bonds is an exothermic process.
  • In general, triple bonds have a larger bond energy than double bonds which have a larger bond energy than single bonds.
  • The change in enthalpy of a reaction (Δ𝐻) can be calculated using the following equation: Δ𝐻=.BEBE()()bondsbrokenbondsformed
  • When Δ𝐻 is positive, the overall reaction is endothermic, and when Δ𝐻 is negative, the overall reaction is exothermic.
  • A reaction profile diagram can be used to represent the energy required to break bonds, the energy released when forming bonds, and the change in enthalpy of a reaction.

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