All the chemical reactions proceed with absorption or evolution of energy. The chemistry dealing with the energy changes during the chemical reactions is called chemical energetic.
The various units of energy are Joules(J), Ergs, Calories(Cal), etc.
1 J = 107 ergs
1 Cal = 4.2 J
- Some thermo – chemical terms
- Different types of system
- State functions and path functions
- Extensive and Intensive properties
- Different types of thermodynamic process
- Internal Energy (E)
- Enthalpy (H)
- Different types of Heat of reaction or Enthalpy of reaction
- Exothermic and Endothermic reaction and their energy profile diagram
- First law of thermodynamics (Law of Conservation of Energy)
- Hess’s law (of constant heat summation)
- Numerical problems
Some thermo – chemical terms
1. System : A system is the specific portion of universe in which the energy change is to be studied. Water taken in a beaker is an example of a system.
2. Surroundings : The remaining portion of universe which is not the part of system is called surrounding. System interacts with the surroundings.
3. Boundary : The part that separates system from the surrounding is called boundary. Boundary is that portion of universe through which system and surrounding can interact with each other. Boundary may be real or imaginary.
Different types of system
1. Open system : A system which can exchange matter as well as energy with the surrounding is called an open system. Eg. Coffee held in a cup.
2. Closed system : A system which can exchange energy but not matter with its surrounding is called a closed system. Eg. Coffee held in a closed metallic (steel) flask.
3. Isolated system: A system which can neither exchange energy nor any matter with the surrounding is called an isolated system.
State functions and path functions
A physical quantity is said to be state function if its value depends upon the initial and final state of the system only and does not depend upon the path by which this state has been attained . For example: temperature , pressure , volume , enthalpy, entropy, free energy, etc.
On the other hand, a physical quantity is said to be path function if its value depends upon the path by which this state has been attained . For example: work , heat, molar heat capacity, etc.
For example, a person standing on the fifth floor of a building has a fixed value of potential energy, irrespective of the fact whether he reached there by stairs or by a lift. Thus, the potential energy of the person is a state function. On the other hand, the work done by the legs of the person to reach the same height is not same in the two cases. Hence, work done is a path function.
Extensive and Intensive properties
Extensive properties: The property whose value depends upon the amount of the substance present in the system is called extensive property. For example: mass, volume, energy, etc.
Intensive properties: The property whose value does not depend upon the amount of the substance present in the system, but depends upon the nature of the substance is called intensive property. For example: temperature, pressure, density, concentration, surface tension, viscosity, etc.
Different types of thermodynamic process
i) Isothermal process: A thermodynamic process in which the temperature of the system remains constant is called an isothermal process. For an isothermal process, ∆T=0.
ii) Adiabatic process: A thermodynamic process in which heat of the system remains constant is called an adiabatic process. For an adiabatic process, ∆q=0.
iii) Isobaric process : A thermodynamic process in which pressure of the system remains constant is called an isobaric process. For an isobaric process, ∆p=0.
iv) Isochoric process : A thermodynamic process in which volume of the system remains constant is called an isochoric process. For an isobaric process, ∆v=0.
v) Cyclic process : When a system in a given state goes through a number of different processes and finally returns to its initial state, then the overall process is called cyclic process. For a cyclic process, ∆E=0 and ∆H=0.
Internal Energy (E)
The total amount of energy stored in a system (substance) under a given set of conditions is called internal energy. The internal energy is the sum of the different types of energies associated with atoms or molecules, such as electronic energy (Ee), nuclear energy (En), chemical bond energy (Ec), potential energy (Ep), and kinetic energy (Ek).
E = Ee + En + Ec + Ep + Ek
If the internal energy of a system in the initial state is E1 and in the final state is E2, then the change in internal energy (∆E) is :
∆E = E2 – E1
Enthalpy is the total heat content of a system. It is equivalent to the sum of the internal energy and the product of the pressure and volume of the system.
H = E + PV
H = Enthalpy, E = Internal Energy, P = Pressure and V = Volume of the system.
Enthalpy change(∆H) :
∆H = ∆E + P∆V
Different types of Heat of reaction or Enthalpy of reaction
1. Heat of combustion: The heat of combustion of a substance is defined as the heat change (usually heat evolved) when one mole of a substance is completely burnt or oxidized in oxygen. Eg.
Completely oxidized means:
For example, carbon may be oxidized to CO and CO2 . Completely oxidation means oxidation to CO2 and not to CO.
2. Heat of formation : The heat of formation of a substance is defined as the heat change i.e. heat evolved or absorbed when one mole of substance is formed from its constituent elements under a given conditions of temperature and pressure. It is usually represented by ∆Hf.
Standard heat of formation: The heat change i.e. heat evolved or absorbed when one mole of substance is formed from its constituent elements in standard states of temperature and pressure ( i.e. 298K and 1 atm) is called standard heat of formation. It is usually represented by ∆Hf0. Eg.
When one mole of CO2 is formed from its elements – C and O2 in standard state, 393.5KJ heat is produced. Hence, the standard heat of formation of CO2 is 393.5KJ.
3. Heat (enthalpy) of neutralization: The heat change i.e enthalpy change when one gram equivalent of an acid is neutralized by base in dilute aqueous solution is called heat of neutralization.
For example: when one gram equivalent of HCl is neutralized by one gram equivalent of NaOH in dilute aqueous solution, 57.1KJ of heat is produced. Thus, heat of neutralization of HCl with NaOH is 57.1KJ.
- Heat of vaporization– Enthalpy change when one mole of a substance changes from liquid state into vapour state.
- Heat of solution– Enthalpy change when one mole of a solute is dissolved in an excess of solvent at a given temperature so that further dilution involves no heat change.
- Heat of sublimation– Enthalpy change when one mole of solid converts into vapour at a given temperature and pressure.
- Heat of fusion– Enthalpy change when one mole of solid is converted into liquid state at its melting point at one atmospheric pressure.
Exothermic and Endothermic reaction and their energy profile diagram
The reaction in which heat is liberated /released from the system is called exothermic reaction.
For example:- Combustion of methane in oxygen is an exothermic process.
Endothermic reaction :
The reaction in which heat is absorbed from the surrounding is called endothermic reaction. For example:- Combination of H2 and I2 to give HI is an endothermic process.
First law of thermodynamics (Law of Conservation of Energy)
This law state that, “the total energy of the universe i.e. system and surrounding always remains constant, however it may changes from one form to another.” i.e. Energy can neither be created nor destroyed.
Let us consider a system having initial internal energy (E1). If heat (q) is supplied to the system and work (w) is done on the system then the internal energy in the final stage i.e. E2 is given as,
E2 = E1 + q + w
E2-E1 = q + w
Therefore, ∆E = q + w ………..(i)
This equation (i) is the mathematical statement of first law of thermodynamics which shows the relationship among internal energy, work and heat.
In thermodynamics, w=P∆V, then,
∆E = q + P∆V
Where, ∆V is the change in volume and P is the external pressure.
Hess’s law (of constant heat summation)
“The total amount of heat evolved or absorbed in a reaction (i.e. enthalpy change) in a reaction is always same whether the reaction takes place in one step or in a number of steps. In other words, the total amount of heat change in a reaction depends only upon the nature of the initial reactants and the final products and is independent of the path or the manner by which this change is brought about.”
Consider a general reaction,
Suppose the heat evolved in this reaction is Q1 Joules.
Now suppose the same reaction takes place in three steps as follows:
Suppose the heat evolved in these three steps are q1, q2 and q3 Joules respectively.
Thus the total heat evolved (suppose Q2) = q1 + q2 + q3 Joules.
Then, according to Hess’s law, we must have Q1 = Q2
To illustrate the Hess’s law of constant heat summation, let us take the example in which carbon is burnt to CO2.
It is also possible to carry out this reaction in two steps as,
From this example, it is clear that :