Plus One Chemistry Notes Chapter 6 Thermodynamics is part of Plus One Chemistry Notes. Here we have given Kerala Plus One Chemistry Notes Chapter 6 Thermodynamics.
|Text Book||NCERT Based|
|Category||Plus One Kerala|
Kerala Plus One Chemistry Notes Chapter 6 Thermodynamics
The study of energy transformations forms the subject matter or thermodynamics.
A system in thermodynamics refers to the part of universe which is under investigation.
Surroundings include everything other than the system or remaining part of the universe.
Type of Systems
Open System. In an open system, there is ex-change of energy and matter between system and surroundings.
Closed System. In a closed system, there is no exchange of matter, but exchange of energy is possible between system and the surroundings.
Isolated System. In an isolated system, there is no exchange of energy or matter between the system grid the surroundings.
State function or State variables and Path function
A state function is a property of the system whose values depends upon the initial and final states of the system and independent of the path by which the state is reached.
Path function is a property of the system whose values depend on the path followed,
eg. Heat (q), Work (w).
Intensive and Extensive Properties
1. Intensive properties. The properties whose value depends on the nature of the substance and is independent of its amount in the system are called intensive properties
eg., density, temperature, pressure, etc.
2. Extensive properties. The properties whose value depends on the quantity of matter but are independent on the nature are called extensive properties,
eg., mass, volume, internal energy, enthalpy heat capacity, etc.
- Isothermal process. A process is said to be isothermal if the temperature of the system remains constant during each stage of the process. ΔT = 0
- Adiabatic process. In adiabatic process, there is no exchange of heat between the system and the surroundings. Δq = 0.
- Isobaric process. Pressure on 9 system remains constant. Δp = 0
- Isochoric process. Volume on a system remains constant. Δv = 0
- Cyclic process. In cyclic process, a system undergoes a series of operation or change and returns back to its initial state.
- Reversible process. The process is said to be reversible if the process is carried out so slowly that the system and the surroundings are always in equilibrium.
- Irreversible process. The process which is carried out so rapidly that the system does not get a chance to attain equilibrium.
Internal energy (U):
The total energy contained in a system is called its internal energy or intrinsic energy. The internal energy of a system is made up of a number of components such as translational kinetic energy of molecules, intermolecular interaction energy of the constituent particles of the system, bond energy, sum of the energies of the occupied orbitals or electronic energy etc.
The Internal Energy as a State Function
1. Work (W). Exchange of energy between system and surrounding can occur in the form of work, which may be mechanical, or electrical work, eg., let us take a system containing some quantity of water in a thermos flask or in an insulated beaker. This would not allow exchange of heat between the system and surrounding through its boundary, such systems are called adiabatic systems.
2. Heat (q). The internal energy of a system by transfer of heat from the surrounding to the system or vice-versa without expenditure of work.
This exchange of energy, which is a temperature difference is called heat noted by q.
First Law of Thermodynamics (Law of Conservation of Energy)
The first law of thermodynamics states that energy can neither be created nor destroyed. It can be transformed from one form to another. The total energy of the universe (system + surroundings) remains constant.
The first law of thermodynamics states that the total mass and energy of an isolated system remains unchanged or constant.
Mathematical statement of the first law of Thermodynamics
Let us consider a system in its initial state having internal energy U1 If heat equal to q is supplied to the system and work done (w) on the system, then internal energy of system in the final state
U2 is given by,
U2 = U1 + q + w
U2 – U1 = q + w
AU = q + w
This relationship between change in internal energy, heat given to the system and work done on the system is the mathematical statement of first law of thermodynamics.
Applications-Work (pressure-volume change)
Work done in reversible isothermal expansion of a gas.
Consider a cylinder which contains one mole of an ideal gas fitted with a frictionless piston. The external pressure, (Pex) on the piston is equal to the pressure of the gas inside, (Pin) the : cylinder.
i. e., Pex = Pin for reversible process. If the external pressure is lowered by an infinitesimal amount Δp, i.e., it falls from Pin to Pin Δp. Also the gas will expand by an infinitesimal volume,
Δv, i.e., volume changes from V to V +ΔV.
In the case of expansion of a gas, the work is done by the system on the surroundings and W has a negative sign.
Enthalpy is the sum total of internal energy and pressure volume or work
The heat required to rise the temperature of the system in case of heat absorbed by the system.
The increase of temperature is proportional to the heat transferred. The magnitude of the co-efficient depends on the size, composition and nature of the system. We can also write it as q = CΔT
The coefficient, C is called the heat capacity. C is directly proportional to amount of substance. The molar heat capacity of a substance, (Cm) is the heat capacity for one mole of the substance and is the quantity of heat needed to raise the temperature of one mole by one degree Celsius (or one kelvin).
Specific heat capacity is the quantity of heat required to raise the temperature of one unit mass of a substance by one degree Celsius (or one kelvin).
Exothermic and Endothermic Reaction
In an exothermic reaction, heat is evolved and the system loses heat to the surroundings. Here qp and ΔrH will be negative. In an endothermic reaction, heat is absorbed and the system gains heat from the surroundings.
Enthalpy changes during phase transformations
This include transition from solid to liquid (fusion), liquid to vapour (vapourisation) and solid to vapour (sublimation).
Molar enthalpy of fusion (ΔfusH) or Standard enthalpy of fusion (ΔfusH°).
The enthalpy change that accompanies melting of one mole of a solid substance into its liquid state at its melting point in the standard state is called standard enthalpy of fusion or molar enthalpy of fusion.
Molar enthalpy of vapourisation (ΔvapH) or Standard enthalpy of vapourisation (ΔfusH°). The enthalpy change that accompanies the vapourisation of one mole of a liquid substance into its gaseous state at its boiling point is called standard enthalpy of vapourisation or molar enthalpy of vapourisation.
Molar enthalpy of sublimation (ΔsubH) or Standard enthalpy of sublimation (ΔsubH°).
Standard enthalpy of sublimation as the change in enthalpy when one mole of 4 solid substance changes into its gaseous state (sublimes) at a temperature below its melting point in the standard state.
Molar enthalpy of formation (ΔfH) or Standard enthalpy of formation (ΔfH°).
The standard enthalpy change for the formation of one mole of a compound from its elements in their standard or reference states is called standard molar enthalpy of formation.
Hess’s law of Constant heat Summation
It states that “the enthalpy change of a chemical reaction is the same whether the change take place in one step or several steps”.
Enthalpies for Different types of Reactions
Standard enthalpy of combustion (ΔcH°).
It is defined as the enthalpy change per mole of a substance when it undergoes combustion and all the reactants and products are in their standard states.
Standard enthalpy of atomization (ΔaH°).
It is the enthalpy change on breaking one mole of bonds completely to obtain atoms in the gas phase in their standard states.
Bond dissociation enthalpy or bond enthalpy (ΔbondH°).
It is the enthalpy change when 1 mole of covalent bonds of a gaseous covalent compound is broken to form products in the gas phase.
Standard enthalpy of solution (ΔsolH°).
It is the enthalpy change when one mole of a substance is dissolved in a specified amount of solvent at a particular temperature in their standard states.
Lattice enthalpy (Δlattice H°).
The lattice enthalpy of an ionic compound is the enthalpy change which occurs when one mole of an ionic compound dissociates into its ions in gaseous state.
A process which has an urge or a natural tendency to occur under a given set of conditions is known as a spontaneous process.
The direction of a spontaneous process for I which the energy is constant is always the one that increases the molecular disorder or randomness. The thermodynamic property that measures the amount of molecular disorder is called the entropy, denoted by S.
For an isothermal process the change in entropy,
ΔS, is defined as
where, qrev is the heat absorbed during the reversible path.
Gibbs Energy and Spontaneity
Gibbs energy, G = H – TS ; where H, T and S are the enthalpy of the system, temperature in Kelvin and entropy respectively.
ΔG is Gibbs energy change accompanying a process.
For a spontaneous process, ΔGsystem = -ve
For a non-spontaneous process, ΔGsystem, = +ve
At equilibrium ΔGsystem = 0
Consider the equation ΔG = ΔH – TΔS
If a reaction has a +ve enthalpy change ; (ΔH) and positive entropy change (ΔS), it can be ; spontaneous when TΔS is large enough to out-weigh ΔH. This can happen in two ways.
- +ve entropy change of the system can be small in which T must be large.
- The positive entropy change of the system
can be large in which case T may be small The former is one of the reasons why reactions are often carried out at high temperatures.
Conditions for Spontaneity
For a process to be spontaneous ΔG should be -ve. But ΔG = ΔH – TΔS
- When ΔH = -ve and ΔS = +ve, ΔG is always -ve and process is spontaneous.
- When ΔH = +ve and ΔS = -ve, ΔG = +ve, and the process is always non spontaneous.
- When ΔH = +ve and ΔS = +ve, ΔG becomes -ve only when ΔH > TΔS, such reactions are spontaneous only at low temperature.
At equilibrium, ΔG = 0.
Enthalpy of solution
Enthalpy of solution is the enthalpy change as associated with the addition of a specified amount of solute to the specified amount of solvent at a constant temperature and pressure.
Enthalpy of dilution
It is the heat withdrawn from the surroundings when additional solvent is added to the solution.
|– ve||+ve||– ve||Reaction will be spontaneous at all temperatures.|
|+ ve||– ve||+ ve||Non-spontaneous at all temperature but reverse reaction will be spontaneous.|
|– ve||– ve||– ve (at low T) + ve(athighT)||Reaction will be spontaneous at low temperature. Reaction will be non-spontaneous at high temperature.|
|+ ve||+ ve||+ ve (at low T) – ve (at high T)||Reaction will be non-spontaneous at low temperature. Reaction will be spontaneous at high temperature.|
Entropy and second law of Thermodynamics
The second law of thermodynamics introduces the concept of entropy and its relation with spontaneous process.
The total entropy change (ΔStotal ) will be equal to the sum of the changes in entropy of the system
(ΔSsystem) and change in entropy of I surroundings (ΔSsurrounding).
ΔStotal = ΔSsystem +ΔSsurrounding
Predict in which of the following, entropy I increases/decreases. Give reason:
i. Temperature of crystalline solid is raised from 0 K to 115 K.
ii. H2(g) →2H(g)
i. Entropy will increase on increasing the temperature since the particles of solid move with greater speed at higher temperature.At 0 K, there is perfect order of the constituent particles,
Entropy will increase because the number of particles of product are double than that of reactant.
Third law of Thermodynamics
The entropy of any pure crystalline substance approaches zero as the temperature approaches absolute zero. This is called third law of thermodynamics.
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