6. THERMODYNAMICS
It is a branch of science that deals with the relationship between heat and work. Chemical thermodynamics is a
branch of chemistry that deals with the heat changes associated with chemical reactions.
Some definitions:
System and Surroundings
System is the part of the universe which is under observation or investigation. The part of the universe except
system is called surroundings. The system and surroundings are separated by a boundary which may be real or imaginary.
System + surroundings → Universe
Types of systems
Depending on the ability to exchange energy and matter with the surroundings, systems are classified into three
types:
1. Open system: It is a system that can exchange both energy and matter with the surroundings.
E.g. Hot water taken in an open vessel.
2. Closed system: It is a system that can exchange only energy and not matter with the surroundings.
E.g. Hot water taken in a closed vessel.
3. Isolated system: It is a system that cannot exchange both energy and matter with the surroundings.
E.g. Hot water taken in a thermo flask.
Depending on the number of particles, systems are of two types:
1. Microscopic system: It is a system containing a few number of particles.
2. Macroscopic system: it is a system contains a large number of particles.
The properties of a macroscopic system are called macroscopic properties. The important macroscopic properties
are: Temperature (T), Pressure (P), Volume (V), length (l), breadth (b), height (h), internal energy (U), enthalpy (H), entropy
(S), Gibb’s energy (G) etc.
Macroscopic properties can be divided into two:
1. Extensive properties: These are properties which depend on the amount of matter present in the system. Or, these are
the properties which change when a system is divided.
E.g.: Volume (v), length (l), breadth (b), height (h), internal energy (U), enthalpy (H), entropy (S), Gibb’s energy (G), heat
capacity etc.
2. Intensive properties: These are properties which are independent of the amount of matter present in the system. Or,
these are the properties which do not change when a system is divided.
E.g. : Temperature (T), Pressure (P), Volume (V), density, refractive index, molar heat capacity, viscosity, surface tension etc.
State and Path functions
A function or a property that depends only on the initial and final state of a system and not on the path followed is called a
state function.
E.g. for state functions: T, P, V, U, H, S, G etc.
Path functions: These are properties which depend on the path followed also.
E.g. heat (q) and work (w)
Thermodynamic process
A process is the method (path) by which a state change occurs in a system. The different types of thermodynamic
process are:
1. Isothermal process: It is a process that occurs at constant temperature. For this process ΔT = 0 but Δq ≠ 0
2. Isobaric process: It is a process that occurs at constant Pressure. For such a process ΔP = 0
3. Isochoric process: It is a process that occurs at constant volume. For such a process ΔV = 0
4. Adiabatic process: It is a process that occurs at constant heat energy. Here no heat enters into or leaves from the
system. For such a process Δq= 0 but ΔT ≠ 0
5. Cyclic process: It is a process that takes place in a cyclic manner. Here the system undergoes a series of changes and
finally returns to its initial state. For such a process ΔU=0 and ΔH = 0
6. Reversible process: Every process is associated with two types of forces – driving force and opposing force. Driving
force favours the process while opposing force opposes it. If the driving and opposing forces are differed by an
infinitesimally small quantity the process takes place in both directions. Such a process is called reversible process.
ANIL KUMAR K L,HSST CHEMISTRY,GHSS ASHTAMUDI, ASHTAMUDI (P.O),KOLLAM [HSSLiVE.IN]
, 7. Irreversible process: If the driving and opposing forces are differed by a large quantity, then the process takes place
in only one direction. Such a process is called irreversible process.
Heat and Work
Heat: It is a form of energy. Heat flows from a hot body to a cold body, when there is a thermal contact between the two.
When a body absorbs heat, its energy increases and when it evolves heat, its energy decreases. Thus by international
convention, when heat is absorbed by a system, q becomes +ve and when heat is evolved by a system q becomes –ve.
Work: In thermodynamics, there are two types of work – expansion work and non-expansion work.
1. Expansion work (Wexp): It is related to gaseous systems. It is the product of pressure and change in volume. i.e.,
expansion work (Wexp) = -PΔV
For irreversible process, (Wexp) = -PΔV and
For reversible process, (Wexp) = -2.303nRT log(V2/V1)
2. Non-expansion work (Wnon-exp): It is related to electrochemical cells. It is the product of potential difference (E) and
charge (Q). i.e., Wnon-exp = E x Q
By international convention, w becomes +ve, when work is done on the system and w becomes –ve, when work is done by
the system.
Free expansion: Expansion of a gas into vacuum (external pressure = 0) is called free expansion. No work is done during free
expansion of an ideal gas whether the process is reversible or irreversible.
Internal Energy (U)
Every body is associated with a definite amount of energy. This energy is called internal or intrinsic energy. It is the
energy possessed by a body. It is the sum of different types of molecular energies like translational kinetic energy, rotational
kinetic energy, vibrational kinetic energy, electronic energy, nuclear energy etc.
Internal energy of a body is an extensive property and state function. We cannot calculate the exact value of
internal energy, but we can calculate the change in internal energy (ΔU) during a process by using an apparatus called Bomb
Calorimeter.
The change in internal energy (ΔU) = U2-U1, where U1 and U2 are the initial and final internal energies respectively.
The unit of internal energy is kJ/mol.
The internal energy of a system can change in mainly by two ways:
1. By the transfer of heat
2. By doing work
First law of thermodynamics: It is same as law of conservation of energy. It states that energy can neither be
created nor be destroyed. Or, the total energy in the universe is always a constant. Or, the total energy of an isolated system
is always a constant.
Mathematically ΔU = q + w ………………. (1)
Where q is the amount of heat absorbed by the system and w is the amount of work done on the system.
If there is only expansion work, the above equation becomes ΔU = q – PΔV ………………(2)
(since, Wexp = -PΔV)
Special cases:
For a system containing only solids or liquids, ΔV = 0. So ΔU = q
For an isothermal reversible process, ΔU = 0. So q = -w = 2.303nRT log V2/V1
For an adiabatic process, q =0, so ΔU = w.
Significance of ΔU
We know that ΔU = q – PΔV
For a process taking place at constant volume, ΔV = 0. So ΔU = qv
i.e., ΔU gives the amount of heat absorbed or evolved by a system at constant volume.
Enthalpy (H)
It is the total heat content of a system. It is the sum of internal energy and pressure-volume energy of a system.
i.e. H = U + PV
It is a state function and an extensive property. i.e, whenever there is a change in enthalpy, it depends only on the
initial and final values and not on the path followed. If H1 is the enthalpy of a system in the initial state and H2 is that in the
final state, then the change in enthalpy, ΔH = H2 – H1.
The unit of enthalpy is kJ/mol. ΔH is determined by using an apparatus called calorimeter.
For a chemical reaction, if HP is the enthalpy of products and HR is that of reactants, then
ANIL KUMAR K L,HSST CHEMISTRY,GHSS ASHTAMUDI, ASHTAMUDI (P.O),KOLLAM [HSSLiVE.IN]
It is a branch of science that deals with the relationship between heat and work. Chemical thermodynamics is a
branch of chemistry that deals with the heat changes associated with chemical reactions.
Some definitions:
System and Surroundings
System is the part of the universe which is under observation or investigation. The part of the universe except
system is called surroundings. The system and surroundings are separated by a boundary which may be real or imaginary.
System + surroundings → Universe
Types of systems
Depending on the ability to exchange energy and matter with the surroundings, systems are classified into three
types:
1. Open system: It is a system that can exchange both energy and matter with the surroundings.
E.g. Hot water taken in an open vessel.
2. Closed system: It is a system that can exchange only energy and not matter with the surroundings.
E.g. Hot water taken in a closed vessel.
3. Isolated system: It is a system that cannot exchange both energy and matter with the surroundings.
E.g. Hot water taken in a thermo flask.
Depending on the number of particles, systems are of two types:
1. Microscopic system: It is a system containing a few number of particles.
2. Macroscopic system: it is a system contains a large number of particles.
The properties of a macroscopic system are called macroscopic properties. The important macroscopic properties
are: Temperature (T), Pressure (P), Volume (V), length (l), breadth (b), height (h), internal energy (U), enthalpy (H), entropy
(S), Gibb’s energy (G) etc.
Macroscopic properties can be divided into two:
1. Extensive properties: These are properties which depend on the amount of matter present in the system. Or, these are
the properties which change when a system is divided.
E.g.: Volume (v), length (l), breadth (b), height (h), internal energy (U), enthalpy (H), entropy (S), Gibb’s energy (G), heat
capacity etc.
2. Intensive properties: These are properties which are independent of the amount of matter present in the system. Or,
these are the properties which do not change when a system is divided.
E.g. : Temperature (T), Pressure (P), Volume (V), density, refractive index, molar heat capacity, viscosity, surface tension etc.
State and Path functions
A function or a property that depends only on the initial and final state of a system and not on the path followed is called a
state function.
E.g. for state functions: T, P, V, U, H, S, G etc.
Path functions: These are properties which depend on the path followed also.
E.g. heat (q) and work (w)
Thermodynamic process
A process is the method (path) by which a state change occurs in a system. The different types of thermodynamic
process are:
1. Isothermal process: It is a process that occurs at constant temperature. For this process ΔT = 0 but Δq ≠ 0
2. Isobaric process: It is a process that occurs at constant Pressure. For such a process ΔP = 0
3. Isochoric process: It is a process that occurs at constant volume. For such a process ΔV = 0
4. Adiabatic process: It is a process that occurs at constant heat energy. Here no heat enters into or leaves from the
system. For such a process Δq= 0 but ΔT ≠ 0
5. Cyclic process: It is a process that takes place in a cyclic manner. Here the system undergoes a series of changes and
finally returns to its initial state. For such a process ΔU=0 and ΔH = 0
6. Reversible process: Every process is associated with two types of forces – driving force and opposing force. Driving
force favours the process while opposing force opposes it. If the driving and opposing forces are differed by an
infinitesimally small quantity the process takes place in both directions. Such a process is called reversible process.
ANIL KUMAR K L,HSST CHEMISTRY,GHSS ASHTAMUDI, ASHTAMUDI (P.O),KOLLAM [HSSLiVE.IN]
, 7. Irreversible process: If the driving and opposing forces are differed by a large quantity, then the process takes place
in only one direction. Such a process is called irreversible process.
Heat and Work
Heat: It is a form of energy. Heat flows from a hot body to a cold body, when there is a thermal contact between the two.
When a body absorbs heat, its energy increases and when it evolves heat, its energy decreases. Thus by international
convention, when heat is absorbed by a system, q becomes +ve and when heat is evolved by a system q becomes –ve.
Work: In thermodynamics, there are two types of work – expansion work and non-expansion work.
1. Expansion work (Wexp): It is related to gaseous systems. It is the product of pressure and change in volume. i.e.,
expansion work (Wexp) = -PΔV
For irreversible process, (Wexp) = -PΔV and
For reversible process, (Wexp) = -2.303nRT log(V2/V1)
2. Non-expansion work (Wnon-exp): It is related to electrochemical cells. It is the product of potential difference (E) and
charge (Q). i.e., Wnon-exp = E x Q
By international convention, w becomes +ve, when work is done on the system and w becomes –ve, when work is done by
the system.
Free expansion: Expansion of a gas into vacuum (external pressure = 0) is called free expansion. No work is done during free
expansion of an ideal gas whether the process is reversible or irreversible.
Internal Energy (U)
Every body is associated with a definite amount of energy. This energy is called internal or intrinsic energy. It is the
energy possessed by a body. It is the sum of different types of molecular energies like translational kinetic energy, rotational
kinetic energy, vibrational kinetic energy, electronic energy, nuclear energy etc.
Internal energy of a body is an extensive property and state function. We cannot calculate the exact value of
internal energy, but we can calculate the change in internal energy (ΔU) during a process by using an apparatus called Bomb
Calorimeter.
The change in internal energy (ΔU) = U2-U1, where U1 and U2 are the initial and final internal energies respectively.
The unit of internal energy is kJ/mol.
The internal energy of a system can change in mainly by two ways:
1. By the transfer of heat
2. By doing work
First law of thermodynamics: It is same as law of conservation of energy. It states that energy can neither be
created nor be destroyed. Or, the total energy in the universe is always a constant. Or, the total energy of an isolated system
is always a constant.
Mathematically ΔU = q + w ………………. (1)
Where q is the amount of heat absorbed by the system and w is the amount of work done on the system.
If there is only expansion work, the above equation becomes ΔU = q – PΔV ………………(2)
(since, Wexp = -PΔV)
Special cases:
For a system containing only solids or liquids, ΔV = 0. So ΔU = q
For an isothermal reversible process, ΔU = 0. So q = -w = 2.303nRT log V2/V1
For an adiabatic process, q =0, so ΔU = w.
Significance of ΔU
We know that ΔU = q – PΔV
For a process taking place at constant volume, ΔV = 0. So ΔU = qv
i.e., ΔU gives the amount of heat absorbed or evolved by a system at constant volume.
Enthalpy (H)
It is the total heat content of a system. It is the sum of internal energy and pressure-volume energy of a system.
i.e. H = U + PV
It is a state function and an extensive property. i.e, whenever there is a change in enthalpy, it depends only on the
initial and final values and not on the path followed. If H1 is the enthalpy of a system in the initial state and H2 is that in the
final state, then the change in enthalpy, ΔH = H2 – H1.
The unit of enthalpy is kJ/mol. ΔH is determined by using an apparatus called calorimeter.
For a chemical reaction, if HP is the enthalpy of products and HR is that of reactants, then
ANIL KUMAR K L,HSST CHEMISTRY,GHSS ASHTAMUDI, ASHTAMUDI (P.O),KOLLAM [HSSLiVE.IN]
