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Bond Enthalpy

Bond Enthalpy

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Bond enthalpy is a measure of how much energy is needed to break apart the bonds in a covalent substance. It's also called bond dissociation energy or bond energy. To be more specific, it's the amount of energy required to break one mole of a particular covalent bond in a gas. When you're calculating bond enthalpy, you can only work with substances in the gas phase.

When you're writing about bond enthalpy, you'll need to include the specific covalent bond that you're talking about. For example, the bond enthalpy of one mole of diatomic hydrogen is written as E(H-H). A diatomic molecule means that there are two atoms in the molecule.

Throughout this article, we'll explain what bond enthalpy is and how it's calculated. We'll also show you how to use bond enthalpy to figure out the energy change in a reaction, and how to use it in other calculations. Finally, we'll explore the relationship between bond enthalpy and trends in the enthalpies of combustion of a homologous series.

What is meant by bond enthalpy?

When a molecule has more than one bond to break, the strength of each individual bond may differ due to the effect of other atoms in the molecule. Therefore, we use the mean bond enthalpy, which is the average amount of energy needed to break a specific type of bond across various molecules. These values are always positive, as breaking bonds always requires energy.

Using mean bond enthalpies, we can calculate the enthalpy change of a reaction when experimental data is not available. We can apply Hess' Law to do this. However, it's important to note that bond enthalpies are only approximate values and calculations using them may not be as accurate as using enthalpy of formation/combustion data.

To illustrate, let's consider an example. Suppose we want to calculate the enthalpy change for the reaction between carbon monoxide and steam in the manufacture of hydrogen. The bond enthalpies are given below:

C-O: 1072 kJ/mol
O-H: 463 kJ/mol
C-H: 413 kJ/mol

We can use a Hess cycle to calculate the enthalpy change. First, we draw a diagram of the reactants and products, and then we draw arrows representing the energy changes involved in the reaction. We can then use the bond enthalpy values to calculate the overall energy change.

In summary, bond enthalpy is a useful tool for calculating the energy required to break specific types of bonds. While it may not be as accurate as other methods, it allows us to estimate enthalpy changes when experimental data is not available.

Bond enthalpy calculation
Bond enthalpy calculation

Now let’s break the covalent bonds in each molecule into single atoms using their given bond enthalpies. Remember: There are two O-H bonds in ,One C-O bond in CO,Two C-O bonds in ,And one H-H bond in .

Bond enthalpy calculation
Bond enthalpy calculation

You can now use Hess’ Law to find an equation for the two routes.

∆H = [ 2(464) +1077 ] - [ 2(805) + 436 ]

∆H = -41 kJ

In the next example, we won't use a Hess cycle - you simply count the number of bond enthalpies broken in the reactants and the number of bond enthalpies formed in the products. Let’s have a look!

Some exams might specifically ask you to calculate ∆H using the following method.

Calculate the enthalpy of combustion for ethylene shown below, using the given bond enthalpies.

Enthalpy of combustion is the change in enthalpy when one mole of a substance reacts in excess oxygen to make water and carbon dioxide.

You must begin by rewriting the equation so that we have one mole of ethylene. 

Count the number of bonds being broken and the number of bonds being formed:

Fill the values in the equation below

= 2912 - 4142

= -1230

That's it! You have calculated the enthalpy change of reaction! You can see why this method might be easier than using a Hess cycle.

Perhaps you are curious about how you would calculate ∆H of a reaction if some of the reactants are in the liquid phase. You will need to change the liquid to a gas using what we call the enthalpy change of vaporisation.

Enthalpy of vaporisation () is simply the enthalpy change when one mole of a liquid turns to gas at its boiling point.

To see how this works, let's do a calculation where one of the products is a liquid.

The combustion of methane is shown below.

Calculate the enthalpy of combustion using the bond dissociation energies in the table.

One of the products, , is a liquid. We have to change it to a gas before we can use bond enthalpies to calculate ∆H. The enthalpy of vaporisation of water is +41 kJ.

Use the equation:

= ∑bond enthalpies broken in reactants - ∑bond enthalpies formed in products

∆H = 2648 - 3548

∆H = -900 kJ

Before we round up this lesson, here’s one last interesting thing related to bond enthalpy. We can observe a trend in the enthalpies of combustion in a 'homologous series'.A homologous series is a family of organic compounds. Members of a homologous series share similar chemical properties and a general formula. For example, alcohols contain an -OH group in their molecules and the suffix ‘-ol’. Have a look at the table below. It shows the number of carbon atoms, the number of hydrogen atoms and enthalpy of combustion of members of the alcohol homologous series. Can you see a pattern?

Trends in the combustion enthalpies of a homologous series
Trends in the combustion enthalpies of a homologous series

When it comes to the enthalpy of combustion, we notice a steady increase as the number of carbon and hydrogen atoms in the molecule increases. This is because the combustion process involves breaking C and H bonds, and each successive alcohol in a homologous series has an extra bond. As a result, the enthalpy of combustion for this series increases by approximately 650 kJ/mol for each extra bond. This pattern is really helpful for predicting the enthalpies of combustion for a homologous series using a graph. The calculated values from the graph are often more accurate than the experimental values obtained from calorimetry, which may be affected by factors such as heat loss and incomplete combustion. By comparing the calculated and experimental values of enthalpies of combustion for a homologous series, we can better understand how the combustion process works and make more accurate predictions about the energy released during combustion.

Bond Enthalpy - Key takeaways

In summary, bond enthalpy is the amount of energy required to break one mole of a specific covalent bond in the gas phase. The strength of a covalent bond is affected by its environment, meaning that the same type of bond can have different bond energies in different molecules. To account for this, we use the mean bond energy, which is an average over different molecules. We can use mean bond energy to calculate the enthalpy change (ΔH) of a reaction using the formula: ΔH = Σ bond energies broken - Σ bond energies made. However, this method is only valid when all substances are in the gas phase. In a homologous series, the enthalpies of combustion show a steady increase due to the number of C and H bonds being broken in the combustion process. This trend can be graphed to predict the enthalpies of combustion for a homologous series. The calculated values from the graph are often more accurate than experimental values obtained from calorimetry, which may be affected by factors such as heat loss and incomplete combustion.

Bond Enthalpy

What is bond enthalpy?

Bond enthalpy (E) is the amount of energy required to break one mole of a specific covalent bond in the gas phase. We show the specific covalent bond being broken by putting it in brackets after the symbol E. For example, you write the bond enthalpy of one mole of diatomic hydrogen (H2) as E (H-H).

How do you calculate average bond enthalpy?

Chemists find bond enthalpies by measuring the energy required to break one mole of a specific covalent molecule into single gaseous atoms. Bond enthalpy is calculated as an average over different molecules known as mean bond enthalpy. This is because the same type of bond can have different bond enthalpies in different environments.

Why do bond enthalpies have positive values?

Average bond enthalpies are always positive (endothermic), as breaking bonds always requires energy from the environment.

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