Chemical Energetics (3) Enthalpies of Combustion of Alcohols and Bond Energies

Chemical Energetics (3) The enthalpy of combustion of alcohols and bond energies.

Let’s examine in this blog the implications that follow from measuring the enthalpy of combustion of the alcohols

Here is the data we might have collected:

Name of Alcohol	Formula of alcohol	Enthalpy of combustion ΔHoc     kJ/mol
Methanol	CH3OH	–726
Ethanol	CH3CH2OH	–1367.3
Propan-1-ol	CH3CH2CH2OH	–2021.0
Butan-1-ol	CH3CH2CH2CH2OH	–2675.6
Pentan-1-ol	CH3(CH2)4OH	–3328.7
Hexan-1-ol	CH3(CH2)5OH	–3983.8

Let’s plot this data in the form of standard enthalpy of combustion vs number of carbon atoms per alcohol.

As you can see the plot is linear.

Why?

It seems that adding a —CH₂— group to each alcohol increases the enthalpy of combustion by the same amount.

This feature suggests that the additional bonds contribute a specific amount of energy to the overall combustion of an alcohol. 

In fact, we could probably predict the enthalpy of combustion of heptan-1-ol from this data.

The implication is that bonds require a specific energy value to break a mole of them. 

So what do we mean by bond energy?

Sometimes bond energy is termed bond enthalpy or bond strength or even bond dissociation enthalpy!! (I’m making a collection of these terms !!!)

All four terms though refer to the same thing. 

Bond enthalpy is not merely the energy required to break one mole of covalent bonds.

You need to define the states of the particles and the type of species involved

Here is a definition from a well-known book of data:

“E(X—Y) Bond Energy defined

a)   for X₂ molecules as the molar enthalpy change for the process X₂(g)  =  2X(g)

b)   for XY_n molecules as the molar enthalpy change for the process

1/n XY_n(g)   =   1/n X(g)   +   Y(g)

Both processes are at 298K with individual species pressures of 1 atm. 

A looser definition is here:

The bond dissociation enthalpy is the energy needed to break one mole of the bond to give separated atoms - everything being in the gas state.

Or this from Avogadro.com:

The Bond Enthalpy is the energy required to break a chemical bond. It is usually expressed in units of kJ mol^-1, measured at 298 K. The exact bond enthalpy of a particular chemical bond depends upon the molecular environment in which the bond exists.

Because bonds only exist in elements or compounds the context or chemical environment affects the bond energy.

The energy of a mole of C—H bonds in methane is slightly different then from that of a mole of the same bonds in alcohol.

So tables of bond energies usually caution you to remember that these are AVERAGE values taken over several molecular environments. 

Breaking bonds requires energy to pull the atoms apart so all bond enthalpies are endothermic values.

Now the point of bond enthalpies is to use them to calculate enthalpy changes in different reactions especially in what are called Hess Cycles.

I will blog about Hess Cycles here later. 

Chemical Energetics (2) Measuring the enthalpy change of combustion

Chemical Energetics (2) Measuring the enthalpy change of combustion. 

Fuels like alcohols are one of the easiest compounds to use to measure the enthalpy of combustion.

Most school and college courses will include this experiment.

At a basic level, the experiment involves using a simple copper calorimeter filled with water,  a spirit burner and a thermometer to measure the temperature rise in the water. 

You might be asked to describe how you might compare two or three alcohols and so provide a fair test of the measurements. 

And at a more sophisticated level, you would first calibrate the apparatus with an alcohol of known enthalpy of combustion.  (see below)

Let’s take the basic approach first:

Here’s a simple diagram of the apparatus you might use or have used.

And yes here’s another case of watch the internet.  As you can see there is no thermometer in the water in the calorimeter and the calorimeter is made of glass not a better conducting material like copper. 

But we have a spirit burner and the calorimeter placed at a constant height above the burner. 

There are also two draught shields either side of the apparatus to prevent convection of heat away from the calorimeter. 

So how do we make this experiment a fair measure of heat of combustion between two or three alcohols?

First, keep the calorimeter at the same height above the burner.

Second, keep the flame the same size (very tricky to do).

Third, keep the wick the same length.

Fourth, use the same volume of water in each experiment, say 300ml.

Fifth, raise the temperature of the water by the same number of degrees in each experiment, say 10.0^oC.  The thermometer should read to 0.1^oC.

Better results come with a higher temperature change and a greater volume of water as the percentage error in each measure is reduced. 

You’ll also need snuffer for the spirit burner. 

Here is a better diagram to follow:

How to carry out this experiment

The procedure goes like this:

• Weigh the burner with snuffer (it prevents evaporation of the alcohol before and after burning so preventing the weight actually burned changing, if not the weight burned  would be higher than it actually is)

• Light the wick and place under calorimeter.

• Stir the water in the calorimeter until the temperature rises 10^oC

• Snuff out the burner and reweigh.

You can then calculate the energy transferred to the water in the calorimeter like this:

Energy transferred (in Joules) E  = m c ΔT

M = the mass of water in the calorimeter (not repeat not the mass of alcohol burned!!)

C = specific heat capacity of water (This is the number of Joules of energy that raise the temperature of 1 g of water by 1^oC i.e. 4.2 J/g/ ^oC)

ΔT  =  the temperature change in ^oC

What you can do with the data:

At this point, there are several things you can do with the data.

You could calculate the energy released per gram of fuel burned like this:

Energy per gram (J/g)      =        energy released (J)  

                                                             mass of fuel burnt (g)

But suppose you wanted to measure the enthalpy of combustion of your alcohol.

Let’s remind ourselves of the definition first:

The standard enthalpy change of combustion (denoted  ΔH^o_c ) is the enthalpy change that occurs when one mole of a compound is completely burned in oxygen.

Given this definition we are not going to get close to the value unless we calibrate the apparatus because the apparatus does not use pure oxygen, and it has many places where the heat is wasted because it does not heat the water in the calorimeter but the surrounding air. 

There are draughts despite the draught shields.

The flame is usually yellow so the alcohol is not burning completely.

And we do not burn one mole of the compound.

We need to calibrate the apparatus with a known alcohol whose enthalpy of combustion is known which will give us a conversion factor (heat capacity of the apparatus) based on using the same apparatus each time. 

Provided all else remains the same then burning a different alcohol and applying the conversion factor we can determine a heat of combustion for a different alcohol. 

Here’s how to use propan-1-ol to calibrate your apparatus:

Look up the molar mass of propan-1-ol and its enthalpy change of combustion

M_r is  60.1 g/mol and ΔH^o_c  is -2021kJ/mol. 

The calibration factor is calculated like this:

First, calculate the energy released using this equation:

Mass propan-1-ol burned/g * ΔH^o_c     =     energy produced Q (kJ) — 1

Molar mass propan-1-ol

Then the heat capacity or conversion factor becomes:

heat capacity of the apparatus c (kJ/K)    =    Energy produced (Q)  — 2

                                                                                 Temperature rise (K)

All you then need to do is use exactly the same set up as before only this time with a different alcohol. 

Then with the heat capacity of the apparatus known you can use the temperature rise to calculate the energy produced from a different alcohol.

This energy produced value will of course be the same as before because the same temperature rise has been measured.

Energy produced (kJ)  =  heat capacity (kJ/K) *  temperature rise (K)  — 3

But what’s different with the second alcohol is that the measured mass of alcohol burned is different.

If the alcohol has more carbon atoms the measured mass burnt should be smaller.

So the enthalpy of combustion can then be calculated like this:

Enthalpy of combustion ΔH^o_c (kJ/mol)   =     Energy produced (from — 3)

                                                                                    Number of moles burned

You should be able to see why this is not the standard enthalpy of combustion: the value is not measured under standard conditions e.g. the pressure is not likely to be 1 atm.

Chemical Energetics (1) Endo and Exothermic Reactions

Chemical Energetics (1): Exothermic and Endothermic Reactions

I love reactions that give out heat energy or take in heat energy.

Prometheus gives humanity fire by Jose A Fadul.

It was the ancient Greeks who believed that their god Prometheus gave the world fire!

This abstract watercolor painting depicts Prometheus as the giver of fire for the use of humankind.

In Greek mythology, Zeus once assigned Prometheus the task of forming humans from water and earth, which he did, but in the process, became fonder of humans than Zeus had anticipated.

Zeus did not share Prometheus' feelings and wanted to prevent humans from having power, especially over fire.

Prometheus cared more for humans than for the wrath of the increasingly powerful and autocratic king of the gods, so he stole fire from Zeus' lightning, concealed it in a hollow stalk of fennel, and brought it to humanity.

And wasn’t it Zeus who had Prometheus chained up because of what he had set loose on the world?

But think of the benefits of fire: the warmth and light and wonder!!!

Well back to the less sentimental let’s first define exo and endo thermic reactions.

An exothermic reaction is one in which energy is transferred into the surroundings.

An endothermic reaction is one where energy is taken in from the surroundings

But what’s all that stuff mean: what are the surroundings?

As you can see the surroundings are anything but the reaction itself.

So in this example of a reaction in a beaker or polystyrene cup the water, the cup, the thermometer are all part of the surroundings as is the rest of the Universe!!

Sometimes you see the kind of picture that can leave you with the misleading impression that the surroundings are everything outside the container (flask, beaker etc.) and the system is what’s inside the container. 

Not so!!:

The system in this kind of set up is the reaction chemicals in the container.

The aqueous solution in which the reaction takes place is part of the surroundings absorbing or releasing heat. 

This point will become very important when we describe how to calculate the change of heat or enthalpy for a particular reaction in aqueous solution. 

We can recognise an endo or exo thermic reaction by the temperature change taking place in the surroundings.

In an exothermic reaction the temperature of the surroundings increases.

In an endothermic reaction the temperature of the surroundings decreases. 

Before we go any further we need to understand what the term Enthalpy Change means.

The symbol for Enthalpy Change is ΔH pronounced “delta H”.

Greek letter capital delta is the symbol for “change of” and H stands for “heat” or “enthalpy”.

Enthalpy Change refers to the heat change in a chemical reaction taking place under constant pressure. 

The enthalpy change may be negative

But these changes in temperature must not be confused with the change in heat content in these two reactions.

Let’s look at how the change in enthalpy is represented using a simple enthalpy level diagram.

Here they are:

You will see that in an exothermic reaction the temperature of the surroundings increases but the change in enthalpy is negative: - ΔH

And in an endothermic reaction though the temperature of the surroundings falls the change in enthalpy is positive: + ΔH

We can see how this works out using cold packs here .

Let’s finish this first blog on chemical energetics with some definitions of different enthalpy changes. 

The standard enthalpy of reaction (denoted ΔH^o_reaction) is the enthalpy change for the reaction in the equation that occurs under standard conditions i.e. at a pressure of 1 atmosphere, at a temperature of 298K, with the substances in their physical states normal to these standard conditions.  Solutions must have a concentration of 1mol dm^-3

The standard enthalpy change of formation (denoted  ΔH^o_f ) of a compound is the enthalpy change that takes place when one mole of that compound is formed from its elements in their standard states under standard conditions

The standard enthalpy change of combustion (denoted  ΔH^o_c ) is the enthalpy change that occurs when one mole of compound is completely burned in oxygen.

The standard enthalpy change of neutralisation (denoted  ΔH^o_neut ) is the enthalpy change that occurs when solutions of acid and alkali react together under standard conditions to produce one mole of water.

Standard enthalpy of atomization (denoted  ΔH^o_at ) is the enthalpy change when 1 mole of gaseous atoms is formed from its element in its defined physical state under standard conditions (298K, 1 atm).