Synopsis
Energy is one of the great subject matters of our time, but what is energy actually? In this simulation you will learn the fundamental thermodynamic concepts of enthalpy, entropy and Gibbs free energy. You will also determine the internal energy of a chemical compound by using bomb calorimetry, and you’ll even travel inside the calorimeter to see it in action!
The challenge of storing energy
What is the best way to store energy? There is no simple answer, but in this simulation, you will be encouraged to give it some thought while learning about the very nature of energy and how it connects to chemical reactions. You will assess whether the gasoline component octane is a suitable chemical for energy storage by using a bomb calorimeter.
The fundamental concepts of thermodynamics
The first and second laws of thermodynamics and the core concepts of enthalpy, entropy and Gibbs free energy are introduced in this simulation. You can play around with the energy levels of reactants and products on a virtual energy surface to learn about endothermic and exothermic reactions. The concept of reaction spontaneity is linked to the concept of Gibbs free energy and its temperature dependence is explored in an interactive game. You will have access to a state-of-the-art bomb calorimeter and can travel inside to see it in action, in order to really be able to understand how it works. From here the concept of chemical bond energy is linked to the thermodynamic calculations of enthalpy on the calorimeter output.
Combine theory and practice
Throughout the simulation you will combine the fundamental teachings of thermodynamics with the experimental results from the calorimeter.
Will you be able to suggest a solution for the energy storage challenge?
Learning Objectives
At the end of this simulation, you will be able to…
Define the core thermodynamics concepts of entropy, enthalpy, and free Gibbs energy, and their units
Explain the first and second laws of thermodynamics
Understand and apply the concept of reaction spontaneity
Explain the differences between the enthalpy of combustion, and enthalpy of formation
Understand the relationship between internal energy and enthalpy
Present Hess’s law in connection with performing enthalpy change calculations
Present the concepts of exothermic and endothermic reactions
Techniques in Lab
Calorimetry
Theory
The first law of thermodynamics
The first law of thermodynamics is a consequence of the law of conservation of energy, which states that energy can be transformed from one type of energy to another, but it can be neither created nor destroyed - in other words, the energy of the universe is constant.
Figure 1: The first law of thermodynamics states that the energy of the universe is constant.The change in energy of chemical reactions
When looking at a chemical reaction, this means that if the system (the molecules taking part in the reaction) experiences a change in internal energy, ΔU, then the surroundings (the rest of the universe) will experience a similar change in energy, but with the opposite sign.
The change in internal energy, ΔU, of a system is expressed as
ΔU = q + w
Where q is the amount of heat being transferred and w is the work being performed. The sign of q and w reflects the direction of the flow of heat and work, seen from the systems point of view: If heat flows into the system (which is the case for an endothermic reaction) then q is positive, whereas the opposite is true for an exothermic reaction. If work is being performed on the system (like a piston being pushed to compress the volume) then w is positive, whereas work is being performed by the system (a gas expanding to move a piston) then w is negative.
Calorimetry
Calorimetry is the science of using a calorimeter.
A calorimeter is a device used to measure the change in heat associated with a chemical reaction. The actual output from a calorimeter is a change in temperature of the reaction medium or a water bath. The change in heat of the reaction monitored, ΔH, can be calculated by multiplying the change in temperature, ΔT, with the total heat capacity, C:
ΔH = C⦁ΔT
The change in heat (also called the change in enthalpy), ΔH, of a reaction is closely related to the change in energy, ΔU, associated with that reaction.
Two types of calorimeters exist: Constant pressure calorimeters are used to measure the heat of a reaction in a liquid medium, whereas constant volume calorimeters are used to measure the change in heat of a combustion reaction.
Enthalpy
Enthalpy is a property of a system also referred to as the heat of that system. Heat and temperature are not the same, but they are related by a simple formula. This can be utilized in calorimetry, where the change in temperature is used to calculate the change in enthalpy associated with a chemical reaction. Enthalpy is denoted with the symbol H, and it is measured in joules, J, which is a measure of energy.
Enthalpy, H, is defined as
H = U + pV
Where U is the internal energy of the system, p is the pressure and V is the volume. pV is the same as the work, w. You may recall from the first law of thermodynamics that the change in energy of a system, U, is given by the sum of the heat and the work applied to that system.
Standard state
The standard state of an element or a compound is its state at 25 oC and atmospheric pressure.
The enthalpies of formation of pure elements in their standard state is 0.
Hess's law
Hess's law states that in going from one particular set of reactants to one particular set of products it doesn't matter whether the reaction takes place in one or multiple steps: The change in enthalpy will be the same.
Figure 2: Hess's law is illustrated by this reaction, going from the reactants to the products over the constituent elements in their standard states.
It follows from Hess' Law, that the change in enthalpy of a reaction, ΔHr, can be calculated by subtracting the sum of the enthalpies of formation of the reactants from the sum of the enthalpies of formation of the products, according to following equation:
ΔHr = ΣΔHf(products) - ΣΔHf(reactants)
The second law of thermodynamics
The second law of thermodynamics states that every spontaneous process causes an increase in the entropy of the universe.
A popular example is the fact that your room seems to get messed up all by itself (spontaneously), whereas tidying up requires energy. Why is that? Well, when you decide that the orderly place for your bag to be is on the hook on the door, then every other place increases the disorder of your room. Since there are many more disorderly places for your bag (and every other item in your room) than orderly, the chances of disorder are simply higher and the increase in entropy is a matter of statistics.
Figure 3: Increasing the entropy of your room is a spontaneous process, whereas tidying up requires energy.
Entropy
Entropy is often interpreted as a measure of the disorder or randomness of a system. It is probably more correct to say that it is a representation of the number of microstates available to a system. The second law of thermodynamics states that every spontaneous process moves towards greater entropy, so towards more available microstates.
Gibbs free energy
In thermodynamics, the Gibbs free energy, G, is the energy associated with a chemical reaction that can be used to do work at a constant temperature and pressure. The unit of Gibbs free energy, G, is joules, J.
The Gibbs free energy is given as the sum of the enthalpy, H, and the temperature, T, times the entropy, S:
G = H - TS
The change in Gibbs free energy, ΔG, for a reaction at a certain temperature determines the
spontaneity of the reaction at that temperature and is given by the equation
ΔG = ΔH - TΔS
If ΔG is:
Negative: the process is spontaneous and may proceed as written. A reaction with a negative ΔG value is called an exergonic reaction.
Positive: the process is non-spontaneous as written, but it may move spontaneously in the reverse direction. A reaction with a positive ΔG value is called an endergonic reaction.
Zero: the process is at equilibrium.
Remember that ΔG will have different values, or even different signs, at different temperatures.