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Thermodynamics; Introduction

Thermodynamics is used to describe heat and energy changes and chemical equilibria. It allows the prediction of equilibria positions, but NOT the rates of reaction or rates of change.


V The temperature scale and absolute zero

V The Zeroth Law of Thermodynamics

V The First Law of Thermodynamics
V The Second Law of Thermodynamics
V The Third Law of Thermodynamics
V Internal energy


'classical thermodynamics ...... is the only physical theory of universal content which I am convinced will never be overthrown'

Albert Einstein 


This is a very basic introduction to the terms and ideas of thermodynamics. Of great importance is that the laws of thermodynamics are absolute laws of the universe and cannot be circumvented. Remember this and you will not be misled by perveyors of 'too good to be true' ideas such as running cars on just water.


Heat is the amount of energy flowing from one body of matter to another spontaneously due to their temperature difference or by any means other than through work or the transfer of matter


Energy is difficult to define precisely being based on the Laws of thermodynamics. In different circumstances, it can be the capacity to do work, or the capacity to provide heat or radiation. There can be converted between different forms of energy, but energy cannot be created or destroyed.


Work is the energy associated with the action of a force.


Although advanced thermodynamics can appear daunting when first encountered, there are just three primary concepts: energy, entropy, and absolute temperature

The temperature scale and absolute zero

Absolute zero is the lowest limit for the thermodynamic temperature scale; nothing can be colder. It is defined as zero kelvin and cannot be reached. The temperature scale defines the triple point of water as +273.16 K exactly and this sets the kelvin temperature scale and the size of the kelvin. The closest approach to 0 K achieved so far is about 0.000 000 001 K.

The Zeroth Law of Thermodynamics

The Zeroth Law of Thermodynamics states that two bodies in contact will come to the same temperature. It follows that If a body A is in thermal equilibrium with two other bodies, B and C, then B and C are in thermal equilibrium with one another.

The First Law of Thermodynamics

The First Law of Thermodynamics is the law of conservation of energy:

                                            'Energy can neither be created nor destroyed'


It can be expressed in everyday terms:


                                           You can't win, you can only break even

                                           You do not get anything for nothing
                                           There is no such thing as a free lunch

                                           The energy of the universe is constant


It states that the energy in an isolated system is conserved, where energy is the capacity to do work. Heat energy can do work by (for example) changing a temperature or pressure. The isolated system may be a chemical reaction, a natural process, a cell, the earth, etc. If these systems are isolated, neither energy nor matter can enter or leave.

Thermodynamics introduces a term H that is the ENTHALPY; a measure of the heat content of the system.

    ΔH is the CHANGE IN ENTHALPY; the heat lost or gained


By convention:

    ΔH is negative when heat is released by the system; such as in exothermic processes

    ΔH is positive when heat is absorbed by the system; such as in endothermic processes


In a sequence of reactions the overall change in enthalpy is the sum of the enthalpies involved:

ΔHoverall = Σ ΔH

             A = B             ΔH1                     e.g.         C + ½O2       -->   CO            ΔH1 
             B = C             ΔH2                     e.g.         CO + ½O2    -->   CO2           ΔH2                 
sum      A+B = B+C    ΔH1 + ΔH2                            C + O2        -->   CO2        total ΔH = ΔH1 + ΔH2    


But ΔH does not tell us if or how fast the process will go: e.g.

            desk burning;       wood + O2     -->     CO2 + H2O        ΔH is negative, heat is given out

We know that a desk will not spontaneously burn as the reaction is incredibly slow. It would burn if we created a fire'


            melting of ice;                   ice      -->    water                 ΔH is positive, heat is absorbed

We know that ice will melt if the temperature is above 0 °C.


Therefore an enthalpy change, by itself, cannot predict the direction of a process.

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The Second Law of Thermodynamics

The Second Law of Thermodynamics states that the total disorder in an isolated system can only increase over time.


It can be expressed in everyday terms:


                                           You can't even break even (except at absolute zero)
                                           The house always wins
                                           The universe is becoming more chaotic

                                           Disorder within the universe always increases with time
                                           A perpetual motion machine cannot be built

                                           No process for converting heat into energy is 100 % efficient

                                           Heat only spontaneously flows from a hot object to a cold one, not from cold to hot




Entropy explanation



It states that the total order in an isolated system cannot increase over time. If part of such a system becomes more ordered, other parts must become even more disordered. In ideal cases, the amount of order may remain constant. There is only one way in which the entropy of a supposedly closed system can be decreased, and that is to transfer heat from the system (that is not closed).


Thermodynamics introduces a term S that is the ENTROPY; a measure of disorder and chaos of the system. In a simple equiprobable system it may be defined as

S0 = kB ˣ Ln(N)

where kB is the Boltzmann constant and N is the number of configurations.

    ΔS is the CHANGE IN ENTROPY; the change in order or disorder

ΔS = Σ{SProduct } − Σ{SReagent}

ΔSoverall = Σ ΔS

By convention:  

    ΔS is positive when disorder increases; the system is more chaotic and disorganized; e.g. a liquid turning to a gas.

    ΔS is negative when order increases; the system is more ordered and organized; e.g. a liquid turning to a crystalline solid.


If there is no change in enthalpy but a process proceeds, there must be an increase in entropy; e.g. gases mixing.

If two systems are combined, the final entropy is greater than the sum of the parts.


Entropy change, by itself, cannot predict the direction of a process


                             2H2 + O2    -->      2H2O       clearly goes with negative entropy change as
     3 molecules of a mixture     -->     2 molecule of the same product


Thus, this is a process that proceeds to give a more ordered product. However, a large amount of heat is produced (negative enthalpy change) that increases the kinetic energy and disorder in the products and the surroundings.


Therefore an entropy change, by itself, cannot predict the direction of a process.

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The Third Law of Thermodynamics

The Third Law of Thermodynamics addresses the problem concerning the direction of a process and It follows from a combination of the First and Second Laws. The entropy of a perfect crystal at absolute zero is exactly equal to zero.


It can be expressed in everyday terms:


                                          You cant reach absolute zero

                                          You can't stay out of the game
                                          The First and Second Laws cannot be got around



 Thermodynamics introduces a term G is the GIBBS FREE ENERGY; the ability to do work of the system at constant temperature and pressure. G is sometimes called, just, the 'free energy' or 'Gibbs energy'.

    ΔG is the change in the Gibbs free energy; Gibbs free energy can do work at constant temperature and pressure


In living systems (constant temperature and pressure);

ΔG = ΔH - T ΔS


    ΔG is the maximum work obtainable from a process
    ΔG is negative when the system is able to proceed; the process is exergonic and

        there is a positive flow of energy from the system to the surroundings
    ΔG is positive when the system is unable to proceed; the process is endergonic and
       it takes more energy to start the reaction than what you get out of it.

    ΔG is zero when the system is at equilibrium


Every reaction has a characteristic ΔG under defined conditions

Under standard conditions (usually 1 M reactants and products, 298.15 K (25 °C), 100 kPa), this is called the Standard Free Energy change and given the symbol, ΔG°.

Where the pH = 7, rather than 1 M this is given the symbol, ΔG°'

For the reaction  A + B = C + D


where R = gas constant (8.31 J ˣ mol-1 ˣ K-1), T is the temperature (in kelvin) and Ln is the natural logarithm;


Given ΔG is zero when the system is in equilibrium, therefore

ΔG°'= - RTLn (Keq°')


In a sequence of reactions:

ΔGoverall = Σ ΔG


                        A + B = C + D           ΔG1
                       C + E = F                   ΔG2
sum          A + B + E = D + F         ΔG = ΔG1 + ΔG2


So long as the overall ΔG is negative the reaction will go from left to right; A + B + E --> D + F

As an example;

                                                                                                                    ΔG°', kJ ˣ mol-1
                          Glucose + phosphate = Glucose-6-phosphate + H2O             +13.8
                                         ATP + H2O = ADP + phosphate                              - 30.5
                                    Glucose + ATP = ADP + Glucose-6-phosphate             - 16.7


Under standard conditions (at pH 7), the process directions are determined by ΔG°';

Glucose-6-phosphate + H2O --> Glucose + phosphate

ATP + H2O --> ADP + phosphate  

Glucose + ATP --> ADP + Glucose-6-phosphate  


The ATP hydrolysis pulls the phosphorylation of the glucose.


ΔG depends on the concentration of the reactants and products as well as the temperature and ΔG°'.


Process direction depends on ΔG

Internal energy

Thermodynamics introduces a term U that is the INTERNAL ENERGY; the energy contained within the system.

H = U +( P ˣ V)

G = U + (P ˣ V) - (T ˣ S)

where P is the pressure and T is the temperature.


ΔU is the change in the Internal energy; ΔU equals the heat added to a system less the work done by the system.


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