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Thermodynamics Thermodynamics (Greek word thermo means heat and dynamics means motion) is the branch of science which deals with the study of different forms of energy and the quantitative relationships between them. The complete study of thermodynamics is based upon three generalizations called first, second and third law of thermodynamics. These laws have been arrived purely on the basis of human experience and there is no theoretical proof for any of these laws.
Basic concepts
(1) System, surroundings and Boundary : A specified part of the universe which is under observation is called the system and the remaining portion of the universe which is not a part of the system is called the surroundings. The system and the surroundings are separated by real or imaginary boundaries. The boundary also defines the limits of the system. The system and the surroundings can interact across the boundary. (2) Types of systems (i) Isolated system : This type of system has no interaction with its surroundings. The boundary is sealed and insulated. Neither matter nor energy can be exchanged with surrounding. A substance contained in an ideal thermos flask is an example of an isolated system. (ii) Closed system : This type of system can exchange energy in the form of heat, work or radiations but not matter with its surroundings. The boundary between system and surroundings is sealed but not insulated. For example, liquid in contact with vapour in a sealed tube and pressure cooker.
(iii) Open system : This type of system can exchange matter as well as energy with its surroundings. The boundary is neither sealed nor insulated. Sodium reacting with water in an open beaker is an example of open system. (iv) Homogeneous system : A system is said to be homogeneous when it is completely uniform throughout. A homogeneous system is made of one phase only. Examples: a pure single solid, liquid or gas, mixture of gases and a true solution. (v) Heterogeneous system : A system is said to be heterogeneous when it is not uniform throughout, i.e. , it consist of two or more phases. Examples : ice in contact with water, two or more immiscible liquids, insoluble solid in contact with a liquid, a liquid in contact with vapour, etc. (vi) Macroscopic system : A macroscopic system is one in which there are a large number of particles (may be molecules, atoms, ions etc. ) (3) Macroscopic properties of the system Thermodynamics deals with matter in terms of bulk (large number of chemical species) behaviour. The properties of the system which arise from the bulk behaviour of matter are called macroscopic properties. The common examples of macroscopic properties are pressure, volume, temperature, surface tension, viscosity, density, refractive index, etc. The macroscopic properties can be subdivided into two types, (i) Intensive properties : The properties which do not depend upon the quantity of matter present in the system or size of the system are called intensive properties. Its examples are pressure, temperature, density, specific heat, surface tension, refractive index,
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Thermodynamics and Thermochemistry
Chapter
viscosity, melting point, boiling point, volume per mole, concentration etc. (ii) Extensive properties : The properties whose magnitude depends upon the quantity of matter present in the system are called extensive properties. Its examples are total mass, volume, internal energy, enthalpy, entropy etc. These properties are additive in nature. Any extensive property if expressed as per mole or per gram becomes an intensive property. (4) State of a system and State Variables Macroscopic properties which determine the state of a system are referred to as state variables or state functions or thermodynamic parameters. The change in the state properties depends only upon the initial and final states of the system , but it is independent of the manner in which the change has been brought about. In other words, the state properties do not depend upon a path followed. (5) Thermodynamic equilibrium : “A system is said to have attained a state of thermodynamic equilibrium when it shows no further tendency to change its property with time”. The criterion for thermodynamic equilibrium requires that the following three types of equilibrium exist simultaneously in a system, (i) Chemical Equilibrium : A system in which the composition of the system remains fixed and definite. (ii) Mechanical Equilibrium : No chemical work is done between different parts of the system or between the system and surrounding. It can be achieved by keeping pressure constant. (iii) Thermal Equilibrium : Temperature remains constant i.e. no flow of heat between system and surrounding. (6) Thermodynamic process : When the thermodynamic system changes from one state to another, the operation is called a process. The various types of the processes are (i) Isothermal process : In this process operation is done at constant temperature. dT = 0 thus E 0. (ii) Adiabatic process : In this a process there is no exchange of heat takes place between the system and surroundings. The system is thermally isolated, i.e., dQ = 0 and its boundaries are insulated. (iii) Isobaric process : In this process the pressure remains constant throughout the change i.e., dP = 0. (iv) Isochoric process : In this process volume remains constant throughout the change, i.e., dV = 0. (v) Cyclic process : When a system undergoes a number of different processes and finally return to its
initial state, it is termed cyclic process. For a cyclic process dE = 0 and dH = 0. (vi) Reversible process : A process which occurs infinitesimally slowly, i.e. opposing force is infinitesimally smaller than driving force and when infinitesimal increase in the opposing force can reverse the process, it is said to be reversible process. (vii) Irreversible process : When the process occurs from initial to final state in single step in finite time and cannot be reversed, it is termed an irreversible process. Amount of entropy increases in irreversible process. Irreversible processes are spontaneous in nature. All natural processes are irreversible in nature
Internal energy, heat and Work
(1) Internal energy ( E ) : “ Every system having some quantity of matter is associated with a definite amount of energy. This energy is known as internal energy .” E E translational E rotational E vibrational E bonding E electronic...... (i) Characteristics of internal energy (a) Internal energy of a system is an extensive property. (b) Internal energy is a state property. (c) The change in the internal energy does not depend on the path by which the final state is reached. (d) There is no change in internal energy in a cyclic process. (e) The internal energy of an ideal gas is a function of temperature only. (f) Internal energy of a system depends upon the quantity of substance, its chemical nature, temperature, pressure and volume. (g) The unit of E is ergs in CGS or joules in SI 1 Joule = 107 ergs. (ii) Change in internal energy ( E ) : It is neither possible nor necessary to calculate the absolute value of internal energy of a system then, E Ef Ein ; E is positive if E (^) f Ein and negative if Ef Ein. (2) Heat (q) and work (w) : The energy of a system may increase or decrease in several ways but two common ways are heat and work. Heat is a form of energy. It flows from one system to another because of the difference in temperature between them. Heat flows from higher temperature to lower temperature. Therefore, it is regarded as energy on the move.
(3) Since gases on heating show considerable tendency towards expansion if heated under constant pressure conditions, an additional energy has to be supplied for raising its temperature by 1 oC relative to that required under constant volume conditions, i.e. , Cp Cv or C (^) p Cv Work doneinexpansion, P V ( R ) where, Cp molar heat capacity at constant pressure Cv molar heat capacity at constant volume. (4) Some useful relations of Cp and Cv (i) C (^) p Cv R 2 calories 8. 314 J (ii) Cv 23 R (for monoatomic gas) and Cv 23 x
(for di and polyatomic gas), where x varies from gas to gas.
(iii) v
p C C (Ratio of molar capacities) (iv) For monoatomic gas, Cv 3 calories whereas, C (^) p Cv R 5 calories
(v) For monoatomic gas, 1. 66 2
3
2
5 ( ) R
R C
C v
p
(vi) For diatomic gas 1. 40 2
5
2
7 ( ) R
R C
C v p
(vii) For triatomic gas(^ )^ CC 68 RR ^1.^33 v
p
Expansion of an ideal gas
(1) Isothermal Expansion : For an isothermal expansion, T 0 ; E 0. According to first law of thermodynamics, E q w q w This shows that in isothermal expansion, the work is done by the system at the expense of heat absorbed. Since for isothermal process, E and T are zero respectively, hence, H 0 (i) Work done in reversible isothermal expansion : Consider an ideal gas enclosed in a cylinder fitted with a weightless and frictionless piston. The cylinder is not insulated. The external pressure, Pext is equal to pressure of the gas, Pgas. Pext Pgas P
If the external pressure is decreased by an infinitesimal amount dP, the gas will expand by an infinitesimal volume, dV. As a result of expansion, the pressure of the gas within the cylinder falls to Pgas dP , i.e., it becomes again equal to the external pressure and, thus, the piston comes to rest. Such a process is repeated for a number of times, i.e., in each step the gas expands by a volume dV.
Since the system is in thermal equilibrium with the surroundings, the infinitesimally small cooling produced due to expansion is balanced by the absorption of heat from the surroundings and the temperature remains constant throughout the expansion. The work done by the gas in each step of expansion can be given as, d (^) w ( P (^) ext dP ) dV Pext. dV dP. dV dP. dV , the product of two infinitesimal quantities, is negligible. The total amount of work done by the isothermal reversible expansion of the ideal gas from volume V 1 to volume V 2 is, given as, 1
log 2 V w nRT V e or
1
- 303 log 10 2 V w nRT V At constant temperature, according to Boyle’s law, P 1 (^) V 1 P 2 V 2 or 12 P 21
P
V
V^ So, 2
- 303 log 10 1 P w nRT P Isothermal compression work of an ideal gas may be derived similarly and it has exactly the same value with positive sign.
1
2 2
- 303 log^1 2. 303 log P nRT P V w nRT V compressio n (ii) Work done in irreversible isothermal expansion : Two types of irreversible isothermal expansions are observed, i.e., (a) Free expansion and (b) Intermediate expansion. In free expansion, the external pressure is zero, i.e., work done is zero when gas expands in vacuum. In intermediate expansion, the external pressure is less than gas pressure. So, the work done when volume changes from V 1 to V 2 is given by (^2) ( 2 1 ) 1 w P dV PextV V V V ext
P gas
P ext
P gas
P ext – dP
dV
Fig. 10.
Since (^) Pext is less than the pressure of the gas, the work done during intermediate expansion is numerically less than the work done during reversible isothermal expansion in which Pext is almost equal to Pgas. (2) Adiabatic Expansion : In adiabatic expansion, no heat is allowed to enter or leave the system, hence, q 0. According to first law of thermodynamics, E q w E w work is done by the gas during expansion at the expense of internal energy. In expansion, E decreases while in compression E increases. The molar specific heat capacity at constant volume of an ideal gas is given by
v (^) dT v C dE or dE C dT v. and for finite change E Cv T So, w E Cv T The value of T depends upon the process whether it is reversible or irreversible. (i) Reversible adiabatic expansion : The following relationships are followed by an ideal gas under reversible adiabatic expansion. PV^ constant where, P = External pressure, V = Volume
v
p C
C
where, Cp molar specific heat capacity at constant pressure, (^) Cv molar specific heat capacity at constant volume.
1
2 1 2
1 2
1 P
P
P
P
T
T
knowing (^) , P 1 , P 2 and initial temperature T 1 , the final temperature T 2 can be evaluated. (ii) Irreversible adiabatic expansion : In free expansion, the external pressure is zero, i.e, work done is zero. Accordingly, E which is equal to w is also zero. If E is zero, T should be zero. Thus, in free expansion (adiabatically), T 0 , E 0 , w 0 and H 0. In intermediate expansion, the volume changes from V 1 to V 2 against external pressure, Pext.
w Pext ( V 2 V 1 )
^
1
1 2
2 P
RT
P
P RT
ext
Pext TPPPTP R
^ 12
21 12
or
^ 1 2
( 2 1 )^2112 PP w C T T RP TP TP v ext
Spontaneous and Non-spontaneous processes
A process which can take place by itself under the given set of conditions once it has been initiated if necessary, is said to be a spontaneous process. In other words, a spontaneous process is a process that can occur without work being done on it. The spontaneous processes are also called feasible or probable processes. On the other hand, the processes which are forbidden and are made to take place only by supplying energy continuously from outside the system are called non-spontaneous processes. In other words, non spontaneous processes can be brought about by doing work. Examples of Spontaneous and Non-spontaneous processes (1) The diffusion of the solute from a concentrated solution to a dilute solution occurs when these are brought into contact is spontaneous process. (2) Mixing of different gases is spontaneous process. (3) Heat flows from a hot reservoir to a cold reservoir is spontaneous process. (4) Electricity flows from high potential to low potential is spontaneous process. (5) Expansion of an ideal gas into vacuum through a pinhole is spontaneous process. All the above spontaneous processes becomes non-spontaneous when we reverse them by doing work. Spontaneous process and Enthalpy change : A spontaneous process is accompanied by decrease in internal energy or enthalpy, i.e., work can be obtained by the spontaneous process. It indicates that only exothermic reactions are spontaneous. But the melting of ice and evaporation of water are endothermic processes which also proceeds spontaneously. It means, there is some other factor in addition to enthalpy change ( H ) which explains the spontaneous nature of the system. This factor is entropy.
Second law of thermodynamics
All the limitations of the first law of thermodynamics can be remove by the second law of thermodynamics. This law is generalisation of certain experiences about heat engines and refrigerators. It has been stated in a number of ways, but all the statements are logically equivalent to one another. (1) Statements of the law (i) Kelvin statement : “It is impossible to derive a continuous supply of work by cooling a body to a
where, k is Boltzmann's constant (4) Entropy changes in system & surroundings and total entropy change for Exothermic and Endothermic reactions : Heat increases the thermal motion of the atoms or molecules and increases their disorder and hence their entropy. In case of an exothermic process, the heat escapes into the surroundings and therefore, entropy of the surroundings increases on the other hand in case of endothermic process, the heat enters the system from the surroundings and therefore. The entropy of the surroundings decreases. In general, there will be an overall increase of the total entropy (or disorder) whenever the disorder of the surroundings is greater than the decrease in disorder of the system. The process will be spontaneous only when the total entropy increases. (5) Entropy change during phase transition : The change of matter from one state (solid, liquid or gas) to another is called phase transition. Such changes occur at definite temperature such as melting point (solid to liquid). boiling point (liquid to vapours) etc, and are accompanied by absorption or evolution of heat. When a solid changes into a liquid at its fusion temperature, there is absorption of heat (latent heat). Let Hf be the molar heat of fusion. The entropy change will be
f f f T
S H
Similarly, if the latent heat of vaporisation and sublimation are denoted by Hvap and Hsub ,^ respectively,^ the^ entropy^ of^ vaporisation^ and sublimation are given by
b vap vap T S H and s sub sub T
S H
Since H (^) f , Hvap and HSub are all positive, these processes are accompanied by increase of entropy and the reverse processes are accompanied by decrease in entropy. (6) Entropy change for an ideal gas : In going from initial to final state, the entropy change, S for an ideal gas is given by the following relations, (i) When T and V are two variables,
1
2 1
ln 2 ln V nR V T S nCv T . Assuming (^) Cv is
constant (ii) When T and p are two variables,
1
2 1
ln 2 ln p nR p T S nC T P . Assuming Cp ,is constant (a) Thus, for an isothermal process (T constant),
1
2 1
ln 2 ln p or nR p V S nR V (b) For isobaric process ( p constant),
1
ln 2 T S nC T p (c) For isochoric process ( V constant),
1
ln 2 T S nC T v (d) Entropy change during adiabatic expansion : In such process q=0 at all stages. Hence S 0. Thus, reversible adiabatic processes are called isoentropic process.
Free energy and Free energy change
Gibb's free energy ( G ) is a state function and is a measure of maximum work done or useful work done from a reversible reaction at constant temperature and pressure. (1) Characteristics of free energy (i) The free energy of a system is the enthalpy of the system minus the product of absolute temperature and entropy i.e.,G H TS (ii) Like other state functions E, H and S , it is also expressed as G. Also G H T Ssystem where S is entropy change for system only. This is Gibb's Helmholtz equation. (iii) At equilibrium G 0 (iv) For a spontaneous process decrease in free energy is noticed i.e. , G ve. (v) At absolute zero, T S is zero. Therefore if G is – ve , H should be – ve or only exothermic reactions proceed spontaneously at absolute zero. (vi) Gsystem T Suniverse , where H 0 (vii) The standard free energy change, Go^ 2. 303 RT log 10 K , where K is equilibrium constant. (a) Thus if K 1 ,then Go^ ve thus reactions with equilibrium constant K >1 are thermodynamically spontaneous. (b) If K <1, then Go^ ve and thus reactions with equilibrium constant K< 1 are thermodynamically spontaneous in reverse direction. (2) Criteria for spontaneity of reaction : For a spontaneous change G ve and therefore use of
G H T S , provides the following conditions for a (^) change to be spontaneous. Table :10.1 Criteria for spontaneity of reaction ΔH ΔS ΔG Reaction characteristics Example
- Always negative Reaction is spontaneous at all temperatures
2 O 3 (^) ( g ) 3 O 2 ( g )
- – Always positive Reaction is non spontaneous at all temperatures
3 O 2 (^) ( g ) 2 O 3 ( g )
- – Negative at low temperature but positive at high temperature
Reaction is spontaneous at low temperature but becomes non spontaneous at high temperature
CaO (^) ( s ) CO 2 ( g ) CaCO 3 ( s )
- Positive at low temperature but negative at high temperature
Reaction is non spontaneous at low temperature but becomes spontaneous at high temperature
CaCO (^) 3 ( s ) CaO ( s ) CO 2 ( g )
Third law of thermodynamics
This law was first formulated by German chemist Walther Nernst in 1906. According to this law, “The entropy of all perfectly crystalline solids is zero at the absolute zero temperature. Since entropy is a measure of disorder, it can be interpretated that at absolute zero, a perfectly crystalline solid has a perfect order of its constituent particles.” The most important application of the third law of thermodynamics is that it helps in the calculation of absolute entropies of the substance at any temperature T. S 2. 303 Cp log T Where CP is the heat capacity of the substance at constant pressure and is supposed to remain constant in the range of 0 to T. Limitations of the law (1) Glassy solids even at 0o K has entropy greater than zero. (2) Solids having mixtures of isotopes do not have zero entropy at 0 o K. For example, entropy of solid chlorine is not zero at 0o K. (3) Crystals of CO , N 2 O , NO , H 2 O , etc. do not have perfect order even at 0 o K thus their entropy is not equal to zero. Thermochemistry “Thermochemistry is a branch of physical chemistry which is concerned with energy changes accompanying chemical transformation. It is also termed as chemical energetics. It is based on the first law of thermodynamics.”
Exothermic and Endothermic reactions
(1) Exothermic reactions : The chemical reactions which proceed with the evolution of heat energy are called exothermic reactions. The heat energy produced during the reactions is indicated by writing + q or more precisely by giving the actual numerical value on the products side. In general
exothermic reactions may be represented as, A B C D q (heat energy) In the exothermic reactions the enthalpy of the products will be less than the enthalpy of the reactants , so that the enthalpy change is negative as shown below H Hp Hr ; H (^) p Hr ; H ve Examples : (i) C ( s ) O 2 ( g ) CO 2 ( g ) 393. 5 kJ (at constant temperature and pressure) or C ( s ) O 2 ( g ) CO 2 ( g ); H 393. 5 kJ (ii) H (^) 2 ( g ) 21 O 2 ( g ) H 2 O ( l ); H 285. 8 kJ (iii) Fermentation is also an example of exothermic reaction. (2) Endothermic reactions : The chemical reactions which proceed with the absorption of heat energy are called endothermic reactions. Since the heat is added to the reactants in these reactions, the heat absorbed is indicated by either putting (–) or by writing the actual numerical value of heat on the reactant side A B C D q (heat energy) The heat absorbed at constant temperature and constant pressure measures enthalpy change. Because of the absorption of heat, the enthalpy of products will be more than the enthalpy of the reactants. Consequently, H will be positive ( ve ) for the endothermic reactions. H Hp Hr ; H (^) p Hr ; H ve Example : (i) N (^) 2 ( g ) O 2 ( g ) 2 NO ( g ); H 180. 5 kJ (ii) C ( s ) 2 S ( s ) CS 2 ( l ) H 92. 0 kJ (iii) Preparation of ozone by passing silent electric discharged through oxygen is the example of endothermic reaction. (iv) Evaporation of water is also the example of endothermic reaction. For exothermic reaction : (^) H or E ve For endothermic reaction : H or E ve
It may be calculated by H^0 [ H^0 f (^) (Products)- H^0 f (Reactants)] The enthalpy or heat of combustion have a number of applications. Some of these are described below, (a) Calorific value of foods and fuels : Energy is needed for the working of all machines. Even human body is no exception. Coal, petroleum, natural gas etc. serve as the principal sources of energy for man-made machines, the food which we eat serves as a source of energy to our body. The energy released by the combustion of foods or fuels is usually compared in terms of their combustion energies per gram. It is known as calorific value. The amount of heat produced in calories or Joules when one gram of a substance (food or fuel) is completely burnt or oxidised. When methane burns, 890.3 kJ mol –^1 of energy is released. ( ) 2 2 () 2 () 2 2 (); 1 mole^4 (16g) CH g O g CO g HOl HCH (^) 4 890. 3 kJ
So, the calorific value of methane = ^89016.^3 55. 6 kJ / g (b) Enthalpies of formation : Enthalpies of formation of various compounds, which are not directly obtained, can be calculated from the data of enthalpies of combustions easily by the application of Hess's law.
- Heatofcombustionof products.
Heatofreaction Heatofcombustionof reactants
(iii) Heat of neutralisation : It is the amount of heat evolved (i.e., change in enthalpy) when one equivalent of an acid is neutralised by one equivalent of a base in fairly dilute solution, e.g., Neutralisation reactions are always exothermic reaction and the value of His ( ve ). HCl ( aq .) NaOH ( aq .) NaCl ( aq .) H 2 O H 13. 7 kcal The heat of neutralisation of a strong acid against a strong base is always constant ( 13. 7 kcalor 57 kJ mole ^1 ). It is because in dilute solutions all strong acids and bases ionise completely and thus the heat of neutralisation in such cases is actually the heat of formation of water from H and OH ions, i.e., H OH H 2 O ; H 13. 7 kcal
In case of neutralisation of a weak acid or a weak base against a strong base or acid respectively, since a part of the evolved heat is used up in ionising the weak acid or base, it is always less than 13. 7 kcalmole ^1 ( 57 kJmole ^1 ). For example, HCN ( aq .) NaOH ( aq .) NaCN ( aq .) H 2 O H 2. 9 kcal HCN ( aq .) ⇌ H +^ + CN – ; H = 10.8 Kcal
- 8 KCal of heat is absorbed for ionisation of HCN it is heat of dissociation or ionisation (iv) Heat of solution : It is the amount of heat evolved or absorbed (i.e., change in enthalpy) when one mole of the solute is dissolved completely in excess of the solvent (usually water). For example, NH (^) 4 Cl ( s ) H 2 O ( l ) NH 4 Cl ( aq .); H 3. 90 kcal BaCl (^) 2 ( s ) H 2 O ( l ) BaCl 2 ( aq .); H 2. 70 kcal (v) Heat of hydration : It is the amount of heat
evolved or absorbed (i.e change in enthalpy) when 1 mole
of an anhydrous or a partially hydrated salt combines with the required number of moles of water to form a specific hydrate. For example, CuSO 4 ( s ) 5 H 2 O ( l ) CuSO 4. 5 H 2 O ( s ); H 18. 69 (vi) Heat of vapourisation : When a liquid is allowed to evaporate, it absorbs heat from the surroundings and evaporation is accompanied by increase in enthalpy. For example: 10. 5 k c als is the increase in enthalpy when one mole of water is allowed to evaporate at 25 oC. When the vapours are allowed to condense to liquid state, the heat is evolved and condensation of vapour is accompanied by decrease in enthalpy. The evaporation and condensation can be represented as, H (^) 2 O ( l ) H 2 O ( g ); H 10. 5 kcals ( 43. 93 kJ ) H (^) 2 O ( g ) H 2 O ( l ); H 10. 5 kcals ( 43. 93 kJ ) Thus the change in enthalpy when a liquid changes into vapour state or when vapour changes into liquid state is called heat of vapourisation. (vii) Heat of fusion : When a solid is allowed to melt, it changes into liquid state with the absorption of heat (increase in enthalpy) and when a liquid is allowed to freeze, it changes into solid with the evolution of heat (decrease in enthalpy). The change in enthalpy of such type of transformations is called enthalpy of fusion. For example, H (^) 2 O ( ice ) H 2 O ( liquid ); H 1. 44 kcals ( 6. 02 kJ ) H (^) 2 O ( liquid ) H 2 O ( ice ); H 1. 44 kcals ( 6. 02 kJ )
(viii) Heat of precipitation : It is defined as the amount of heat liberated in the precipitation of one mole of a sparingly soluble substance when solutions of suitable electrolytes are mixed, for example Ba^2 SO 42 ( aq ) BaSO 4 ( s ): H 4. 66 kcal (ix) Heat of sublimation : Sublimation is a process in which a solid on heating changes directly into gaseous state below its melting point. Heat of sublimation of a substance is the amount of heat absorbed in the conversion of 1 mole of a solid directly into vapour phase at a given temperature below its melting point. I (^) 2 ( s ) I 2 ( g ) ; H 62. 39 kJ Most solids that sublime are molecular in nature e.g. iodine and naphthalene etc. H (^) sub . H fusion Hvaporisation (3) Experimental determination of the heat of reaction : The heat evolved or absorbed in a chemical reaction is measured by carrying out the reaction in an apparatus called calorimeter. The principle of measurement is that heat given out is equal to heat taken, i.e.,Q ( W m ) s ( T 2 T 1 ),
Where Q is the heat of the reaction (given out), W is the water equivalent of the calorimeter and m is the mass of liquid in the calorimeter and s its specific heat,
T 2 is the final temperature and T 1 the initial
temperature of the system. Different types of calorimeters are used but two of the common types are, (i) Water calorimeter and (ii) Bomb calorimeter Bomb calorimeter : This is commonly used to find the heat of combustion of organic substances. Since the reaction in a bomb calorimeter proceeds at constant volume, the heat of combustion measured is E E W mwt t s Mkcal 1
Where M is the molecular mass of the substance,
w 1 is the weight of substance taken, W is the water
equivalent of calorimeter, m is the mass of liquid in the calorimeter and s is the specific heat of liquid. H can be calculated from the relation, H E nRT
Laws of thermochemistry
(1) Levoisier and Laplace law : According to this law enthalpy of decomposition of a compound is numerically equal to the enthalpy of formation of that compound with opposite sign, For example, C ( s ) O 2 CO 2 ( g ); H 94. 3 kcal CO (^) 2 ( g ) C ( s ) O 2 ( g ); H 94. 3 kcal (2) Hess's law (the law of constant heat summation) : This law was presented by Hess in 1840. According to this law “ If a chemical reaction can be made to take place in a number of ways in one or in several steps, the total enthalpy change (total heat change) is always the same, i.e. the total enthalpy change is independent of intermediate steps involved in the change .” The enthalpy change of a chemical reaction depends on the initial and final stages only. Let a substance A be changed in three steps to D with enthalpy change from A to B , H 1 calorie, from B to C , H 2 calorie and from C to D , H 3 calorie. Total enthalpy change from A to D will be equal to the sum of enthalpies involved in various steps, Total enthalpy change H (^) steps H 1 H 2 H 3 Now if D is directly converted into A, let the enthalpy change be H direct. According to Hess's law H steps H direct 0 , i.e. H steps must be equal to H direct numerically but with opposite sign. In case it is not so, say H steps(which is negative) is more that H direct (which is positive), then in one cycle, some energy will be created which is not possible on the basis of first law of thermodynamics. Thus, H steps must be equal to H directnumerically. (i) Experimental verification of Hess's law (a) Formation of carbon dioxide from carbon First method : carbon is directly converted into CO 2 ( g ). C ( s ) O 2 ( g ) CO 2 ( g ); H 94. 0 kcal Second method : Carbon is first converted into CO ( g ) and then CO ( g )into CO 2 ( g ), i.e. conversion has been carried in two steps, C ( s ) 21 O 2 CO ( g ) ; H 26. 0 kcal
CO ( g )^12 O 2 CO 2 ( g ); H 68. 0 kcal Total enthalpy change C ( s ) to CO 2 ( g ); H 94. 0 kcal
Joule thomson coeffient
p H
T
For cooling ve For Heating ve Neither cooling nor heating 0
The temperature at which a real gas shows neither
cooling nor heating effect on adiabatic expansion ( i.e ., 0 ) is called inversion temperature.
Hydrogen has highest calorific value.
13. 7 Kcal / mol 57 KJ / mol (be cause of 1 cal =4.
Joule)
Enthalpy of fusion of ice per mole is 6 KJ.
Order of bond energy in halogen
Cl (^) 2 Br 2 F 2 I 2.
Heat of vapourisation of water per mole is 10.
KCal.
The heat of reaction is independent of the time
consumed in the process.
Basic concepts
1. Internal energy of an ideal gas depends on (a) Volume (b) Temperature (c) Pressure (d) None of these 2. Any series of operations so carried out that at the end, the system is back to its initial state is called (a) Boyle's cycle (b) Reversible process (c) Adiabatic process (d) Cyclic process 3. One calorie is equal to [CPMT 1988] (a) 0.4184 Joule (b) 4.184 Joule (c) 41.84 Joule (d) 418.4 Joule 4. The total internal energy change for a reversible isothermal cycles is (a) Always 100 calories per degree (b) Always negative
(c) 0 (d) Always positive
5. A well stoppered thermos flask contains some ice cubes. This is an example of a [AIIMS 1992] (a) Closed system (b) Open system (c) Isolated system (d) Non-thermodynamic system 6. Identify the intensive quantity from the following [IIT JEE 1993] (a) Enthalpy and temperature (b) Volume and temperature (c) Enthalpy and volume (d) Temperature and refractive index 7. Which of the following units represents the largest amount of energy [CPMT 1989; MP PET 2000] (a) Electron volt (b) Erg (c) Joule (d) Calorie 8. Energy equivalent to one erg, one joule and one calorie is in the order [NCERT 1980; CPMT 1997] (a) (^1) erg 1 joule 1 calorie (b) 1 erg 1 calorie 1 joule (c) 1 calorie 1 joule 1 erg (d) 1 joule 1 calorie 1 erg
