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C 10K Chemical Thermodynamics and Electrochemistry Lecture 1 of 9 Lectures Jan, Feb, March 2009 Dr. W. Pinnock

What is Chemical Thermodynamics?   

Thermodynamics is the study of the transformation of energy into its various forms. Energy can be defined as the capacity to do work or to transfer heat. In Chemical Thermodynamics we are more interested in energy transformations in chemical reactions and in perfect (or ideal) gas processes. Energy Transformations (?)      

Energy is usually released when a chemical reaction occurs, but sometimes it is absorbed. (Exothermic and Endothermic reactions). Take for example the reaction representing the burning of gasoline. C8H18(l) + 12½O2(g) → 8CO2(g) + 9H2O(g) + heat Quite a lot of energy is released when gasoline is ignited and we sense this as heat. If you stand near to the sample of gasoline when it is ignited you feel a rush of hot gas. The reaction system does work in pushing back the surrounding atmosphere as the number of moles of gas increases in the reaction


What is Chemical Thermodynamics?  

Energy Transformations Where does energy come from in an exothermic reactions? Does the reaction create the energy we see released?  

C8H18(l) + 12½O2(g) → 8CO2(g) + 9H2O(g) + heat transferred and work done. Was it present in gasoline or in the oxygen, and is it simply being released in the process (?) or is energy being created in the reaction?

Where does the absorbed energy go in an endothermic reaction? Is it destroyed?    

Endothermic reactions can be exemplified by the mixture of solid hydrated barium hydroxide and ammonium nitrate. Ba(OH)2.8H2O(s) + 2NH4NO3(s) + Heat absorbed → Ba(NO3)2(s) +2NH3(g) + 10 H2O(l) + work done Heat is absorbed in this process and some work done. Is it destroyed or is it stored somewhere in the solution – in a form that is not obvious? wrp

What is Chemical Thermodynamics? 

Energy Transformations

The answer is, of course, that energy is conserved; it is not created nor is it destroyed; it is simply converted from one form to another. (Principle of Conservation of Energy) 

Chemical systems store energy internally as potential and kinetic energy, U and Ekin.  

The gasoline and oxygen molecules must have contained the energy that we see released when gasoline burns. It must have been stored in a form which we could not discern simply by looking at the litre of gasoline. The energy absorbed by the mixture of barium hydroxide and ammonium nitrate is stored in some way which we cannot see by looking at the solution formed.

Kinetic energy is energy of motion of molecules and Potential energy is energy stored in chemical bonds.

In summary then when we talk about “transformation of energy into its various forms” in chemical systems we mean its storage as potential and kinetic energy in substances, and its transfer between system and surroundings as heat or as work done. 



What We Will Do In This Thermodynamics Course? In the first part of the course we will define and use parameters which will help us keep an account of energy changes that occur in chemical reactions and in perfect (ideal) gas processes.


Internal Energy, U, Many American Textbooks use E for Internal Energy. We will use the IUPAC notations and definitions as they are represented in Atkins. Enthalpy, H. You have all come across Enthalpy in your “A” Level or CAPE Thermochemistry Course. In those courses the term Enthalpy is used very loosely to mean heat evolved or absorbed. In this C 10K Course we will define Enthalpy more carefully.

In the second part we will define and use parameters which will help us to predict the direction of spontaneous change, and predict the position of equilibrium in reversible reactions. These are:


Entropy, S and Gibbs Energy (or Free Energy) G.

Let’s start with Internal Energy


First Law of Thermodynamics 

Applying the Principle of Conservation of Energy to a chemical reaction that occurs in a system where mass is conserved (?) we can say more pointedly: “When a chemical reaction occurs in a closed system the change in energy of the system must appear as either heat or work.”  This is a statement of the First law of Thermodynamics.  Note that the Law applies to closed systems only. It is convenient in accounting for energy in chemical reactions, to call the energy which is present in the system, the INTERNAL ENERGY, U. The change in the energy content of the system can be designated as ∆U where  ∆U = Ufinal - Uinitial 

If energy is conserved and can appear only as heat or work, we can write ∆U = q + w for a closed system, such as that represented by a chemical equation. wrp


Recall of the Argument Leading to a Designation of Internal Energy, U  

The energy which is present within a chemical system is what we call the INTERNAL ENERGY, U, of the system. The change in the energy of the system when a reaction occurs can be represented by ∆U, where : ∆U = Ufinal - Uinitial

The energy that leaves or enters the system – if no mass leaves the system - must manifest itself as heat or work, so: ∆U = q + w

This equation is perhaps the best “statement” of the First Law. In words the equation can be expressed thus:  “The change in internal energy of a closed system is the sum of the energy that passes through its boundaries as heat and work” Note carefully that a chemical reaction as written represents a wrp system that is closed because mass is conserved.

Work Done and Heat Transferred 

Work - is done when a body is moved some distance by a force.     

Work = Force x Distance moved in the direction of the force. w=Fxd usually written as vectors Work - is also done when electrical charge is moved against a potential as in the discharge of an electrical cell. Electrical work = charge moved x potential against which it is moved w=qxE charge in Coulombs and potential in volts

Heat - transferred shows up inevitably as a change in temperature, or as an amount of substance that has changed phase. 

Heat = mass x specific heat capacity x change in temperature  q = m c (T2 – T1) where c is specific heat capacity

Heat = mass x enthalpy per unit mass for phase change  q = m (∆ ∆H) where ∆H is the “enthalpy” for the phase change wrp


Units of Energy The SI unit of energy, in whatever form it exists or in whatever mode it is transferred is the Joule. 1 Joule = 1 Newton x meter = 1 N m


Newton = Force required to give a mass of 1 kg an acceleration of 1 ms-2 = 1 kg m s-2 Specific heat capacity is given in units of J K-1mol-1 usually, but is sometimes given in J K-1 g-1. C for water is 4.18 J K-1g-1 and 75.3 J K-1 mol-1


1 Joule = 1 Coulomb x Volt

Some texts use other units like calories, electron-volts, kilowatt-hours or BTU’s for energy, depending on the application, but we will use the SI unit in most cases.



Energy Stored in Chemical Substances 

We have defined a parameter called Internal Energy, U, as the sum of the kinetic energy of atoms/molecules that make up a chemical system, and the potential energy of the system due to interactions between atoms and molecules. 

Kinetic Energy – is energy possessed by virtue of motion  

In Chemistry this is energy of motion of atoms and molecules. It is usually quantized and made up of components: Etrans, Erot, Evib Recall that T, the absolute temperature, is proportional to Ekin according to the Kinetic Theory, so high kinetic energy of atoms and molecules is apparent by high temperature.

Potential Energy – is energy possessed by virtue of position in a “field”  

The most common form of potential energy we encounter is gravitational potential energy, given by Epot = mgh What this means in chemical system is something we will look at later on, but for the moment let us accept the somewhat intuitive idea that this is energy “stored in bonds”. High potential energy is not usually apparent from observation of the wrp substance.


A Sign Convention for the First Law? 

When the change in the energy of the system is designated as ∆U where ∆U = Ufinal - Uinitial we can write: ∆U = q + w for a closed system

It is not very obvious, but there is a sign convention which must be followed to be consistent with the way we have designated ∆U. To illustrate, let’s consider the case where w = 0, i.e no work is done.


∆U = q Suppose heat energy leaves the system when the reaction occurs. Ufinal is less than Uinitial So ∆U (= Ufinal – Uinitial) is a negative quantity and is equal to the heat that has left the system. So q is negative when heat leaves the system.

The sign convention which goes with this statement follows logically from the way we have defined ∆U:    

if heat is added to the system q is +ve if heat is given out by the system q is -ve if work is done on the system w is +ve if work is done by the system w is -ve


The Basis of the Sign Convention Energy Leaving System Uinitial


Energy Entering System Uinitial

Ufinal ∆U = Ufinal - Uinitial

The sign convention which goes with this statement follows logically from the way we have defined ∆U: if heat is added to the system q is +ve if heat is given out by the system q is -ve if work is done on the system w is +ve if work is done by the system w is -ve END LECTURE I



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