How Heat of Combustion & Formation Affects Energy

Instructor: Laura Foist

Laura has a Masters of Science in Food Science and Human Nutrition and has taught college Science.

The heat of combustion and heat of formation tells us how much energy is in a compound. In this lesson, we will learn how they are determined and how they can be used to calculate energy of a reaction.

The First Law of Thermodynamics

The first law of thermodynamics states that energy cannot be created or destroyed. Let's examine this energy diagram:


Energy graph


In this diagram, the ending energy is lower than the starting energy. So, if energy cannot be destroyed, then what happened to this energy? There is another assumption within the first law of thermodynamics: in a ''closed system,'' energy is never created or destroyed. So if the reaction represented in the above energy diagram were to occur in a closed system, then we would see an increase in heat in the surroundings, because energy was released as heat.

If we put this reaction into a closed system and then measure the change in temperature, we can determine just how much energy was lost (or gained) in the reaction. This is exactly what we do to determine the heat of combustion. We take the material of interest, explode it with excess oxygen in a closed system (called a bomb calorimeter), and measure the temperature change. Thus, the heat of combustion is the amount of energy released when a material is completely burned up.

The heat of combustion is always positive, because energy is released. On the other hand, the heat of formation is negative, because energy is absorbed or used in the reaction. The heat of formation is the amount of energy needed to form a material.

Bomb Calorimeter Calculations

Let's imagine that we have run an experiment, burning up ethanol. We put 2 g of ethanol into the bomb calorimeter and quickly burned it, until it was only carbon dioxide and water. The container holding the ethanol was surrounded by 100 g of water; the temperature of the water started out as 25° C and increased to 75° C. Let's determine the heat of combustion.

Heat of combustion is measured per mole, so let's determine how many moles of ethanol we have. The molecular weight of ethanol is 46 g/mol:


Calculate moles of ethanol


We have 0.04 moles of ethanol. Next, we need to determine how much energy was put into the water, using the specific heat of water (4.2 J/g° C). We know that it takes 4.2 J of energy to raise 1 gram of water 1 degree Celsius. We have raised 260 grams of water by 50 degrees Celsius:


Calculate energy added to water


So the water took in 54600 J (54.6 kJ) of energy. This was 54.6 kJ for 0.04 moles:

54.6 kJ / 0.04 mol = 1365 kJ/mol

So, the heat of combustion for ethanol is 1365 kJ/mol.

To unlock this lesson you must be a Study.com Member.
Create your account

Register to view this lesson

Are you a student or a teacher?

Unlock Your Education

See for yourself why 30 million people use Study.com

Become a Study.com member and start learning now.
Become a Member  Back
What teachers are saying about Study.com
Try it risk-free for 30 days

Earning College Credit

Did you know… We have over 200 college courses that prepare you to earn credit by exam that is accepted by over 1,500 colleges and universities. You can test out of the first two years of college and save thousands off your degree. Anyone can earn credit-by-exam regardless of age or education level.

To learn more, visit our Earning Credit Page

Transferring credit to the school of your choice

Not sure what college you want to attend yet? Study.com has thousands of articles about every imaginable degree, area of study and career path that can help you find the school that's right for you.

Create an account to start this course today
Try it risk-free for 30 days!
Create an account
Support