Basic Terms and Concept – I
* The branch of science that deals with the study of different forms of energy and the quantitative relationship between them is known as thermodynamics.
* A system in thermodynamics refers to that part of the universe in which observations are made.
* The part of the universe that is not a part of the system but can interact with the system is called surroundings.
* The part that separates the system from the surroundings is known as boundary.
* A system is said to be open if it can exchange both energy and matter with its surroundings.
* A system is said to be closed if it can exchange energy but not matter with its surroundings.
* A system is said to be isolated if it can exchange neither energy nor matter with its surroundings.
Basic Terms and Concept – II
* The state of a system means the condition of the system in terms of observable properties such as temperature, pressure and volume.
* Extensive properties are properties that depend upon the quantity of the matter contained in a system.
* Intensive properties are properties that are independent of the amount of the substance present in a system.
* The energy stored within a substance or a system is known as its internal energy.
* A change in internal energy is brought out by doing work or by transfer of heat. w is negative if work is done by the system, and positive if work is done on the system. q is negative if heat transfer takes place from the system to the surrounding, and positive if the heat transfer is from the surrounding to the system.
* In an adiabatic process, no heat exchange takes place.
Thermodynamic Quantity: Work
* Thermodynamics is concerned with mechanical work, also known as pressure – volume work.
* The expression of pressure – volume work is written as w = -Pex ∆V.
* A reversible process is one that takes place infinitesimally slowly and the direction of which at any point can be reversed by an infinitesimal change in the state of the system. In a reversible process the system is in equilibrium in the initial, final and all intermediate stages.
Wrev = – 2.303 nRT log \frac{vf}{vi}
If ∆V = 0, then –Pex ∆V = 0
* No work is done during free expansion and change in internal energy equals the heat supplied.
Thermodynamic Quantity: Enthalpy
* Enthalpy of a system is defined as the sum of the internal energy and the product of pressure and volume.
* Enthalpy is a state function. Enthalpy of a substance cannot be measured, but change in enthalpy can be measured.
* If a reaction is carried out at constant temperature and pressure, then the heat exchanged between the system and the surrounding is equal to the change in enthalpy.
* ∆H is negative for exothermic reactions and positive for endothermic reactions.
* The capacity to absorb heat energy and store it is known as the heat capacity of a system. Heat capacity is defined as the amount of heat required to raise the temperature of the system by one degree at a specified temperature.
* The Sl unit of molar heat capacity are joule per degree Kelvin per mole.
* Molar heat capacity of a substance is the quantity of heat needed to raise the temperature of one mole substance through one degree Celsius or one Kelvin.
Cp – Cv = R
Measuring Change in Internal Energy and Change in Enthalpy
* Calorimetry is an experimental technique that involves the measurement of heat changes associated with physical or chemical processes.
* Heat evolved is the heat of combustion at constant volume and is a measure of internal energy
change ∆U.
* When heat is absorbed or evolved by the system at constant pressure, we are actually measuring the changes in the enthalpy. Therefore, ∆H = qp
* In an exothermic reaction, heat is evolved and the system loses heat to the surroundings. Therefore, qp as well as enthalpy change ∆H is negative.
* In an endothermic reaction, heat is absorbed by the system, from the surroundings and therefore, qp as well as the enthalpy change ∆H will be positive.
Reaction Enthalpy
* Enthalpy of a reaction is defined as the amount of heat evolved or absorbed in a chemical reaction, when the numbers of moles of the reactants as expressed in the chemical equation have completely reacted. Enthalpy of a reaction is represented as ∆r H.
* Enthalpy of reaction is expressed as ∆r H = ∑i aᵢ Hproducts – ∑i bᵢ Hreactants
* The standard enthalpy of a reaction is the enthalpy change accompanying the reaction when all the reactants and products are taken in their standard states.
* A substance is said to be in a standard state when it is present in its most stable state under a pressure of one bar at a specified temperature.
* Standard Enthalpy of fusion is the enthalpy change accompanying the melting of one mole of a solid substance in standard state into its liquid state. It is also called as molar enthalpy of fusion and is represented as ∆fus Ho.
* The amount of heat required to convert one mole of a liquid into its vapor state at constant temperature and under standard pressure is called its standard enthalpy of vaporisation or molar enthalpy of vaporisation.
* Enthalpy of sublimation is defined as the enthalpy change accompanying the conversion of one mole of a solid directly into vapor phase at constant temperature and under standard pressure.
Enthalpy of Formation, Thermochemical Equations and Hess’s Law
* The enthalpy change associated with the formation of one mole of a compound from its constituent elements, all substances being in their standard states, that is at 298k and 1bar pressure is called the standard molar enthalpy of formation, it is represented as ∆f H0.
* By convention, the standard enthalpy of formation of all elements is assumed to be zero.
* A balanced chemical equation which indicates the amount of heat evolved or absorbed in the chemical reaction, is known as a thermochemical equation.
* For exothermic reactions, ∆H is negative and for endothermic reactions it is positive.
* If a chemical equation is reversed, the magnitude of ∆H remains the same but the sign of ∆H is reversed.
* Hess’s law states that, if a chemical change can be made to take place in two or more different ways whether in one step or more than one step, the amount of total heat change is the same no matter by which method the change is brought about.
Enthalpy of Combustion and Atomisation
* The enthalpy of combustion of a substance is defined as the heat energy evolved when one mole of substance is completely burnt or oxidised in oxygen. It is represented as ∆cHo.
* Standard enthalpy of combustion is defined as “the enthalpy change per mole of a substance, when it undergoes combustion, with all the reactants and products being in their standard states at the specified temperature.” It is represented by ∆cHo.
* Calorific value is the amount of heat in calories or joules produced from complete combustion of one gram of fuel or food. Calorific value is usually expressed in kJ/gm or kcal/gm.
* When one mole of a given substance dissociates into gaseous atoms, the enthalpy change accompanying the process is called enthalpy of atomisation. It is represented by ∆aHo.
* Bond dissociation enthalpy is the change in enthalpy when one mole of covalent bonds of a gaseous covalent compound is broken to form products in the gas phase.
* For all diatomic molecules, bond dissociation enthalpy is the same as atomisation enthalpy.
Bond Enthalpy
* Enthalpy of a system is defined as the sum of the internal energy and the product of pressure and volume.
* Enthalpy is a state function. Enthalpy of a substance cannot be measured, but change in enthalpy can be measured.
* If a reaction is carried out at constant temperature and pressure, then the heat exchanged between the system and the surrounding is equal to the change in enthalpy.
* ∆H is negative for exothermic reactions and positive for endothermic reactions.
* The capacity to absorb heat energy and store it is known as the heat capacity of a system. Heat capacity is defined as the amount of heat required to raise the temperature of the system by one degree at a specified temperature.
* The Sl unit of molar heat capacity are joule per degree Kelvin per mole.
* Molar heat capacity of a substance is the quantity of heat needed to raise the temperature of one mole of substance through one degree Celsius or one Kelvin.
Cp – Cv = R
Enthalpy of Solution
* The amount of heat energy released or absorbed when one mole of a substance dissolves in a specified amount of solvent is known as enthalpy of solution.
* When water is used as a solvent, the dissolving process is called as hydration. If any other solvent is used, then the process is known as salvation of ions.
* Enthalpy of hydration, of an ion is the amount of energy released when a mole of the ion dissolves in a large amount of water forming an infinite dilute solution in the process.
* 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.
* The lattice enthalpy of a compound depends on the size of the compound, and the charge on the ions that make up the compound.
* A decrease in the size of an ion increases the lattice enthalpy, that is the enthalpy becomes negative and the reaction becomes exothermic.
* An increase in charge also increases the lattice enthalpy. This is because the attraction forces between ions increases as the charge increases.
* It is not possible to determine lattice enthalpy through an experiment. But, the same can be calculated by constructing an enthalpy diagram known as Born-Haber cycle.
Criterion for Spontaneity of a Reaction
* A process, which under some given conditions may take place itself or by initiation is a spontaneous process.
* Processes that take place only by supplying energy continuously from outside the system are called non-spontaneous processes.
* The tendency of a process to be spontaneous depends on two factors:
• Tendency for decrease in enthalpy
• Tendency for maximum randomness
• Randomness or disorder of a system is expressed by a thermodynamic state function known as entropy, represented by the letter S.
* For a given substance, the solid state will have the lowest entropy, while the gaseous state will have the highest entropy. The entropy of the liquid state will be in between the entropy values of solid state and gaseous state.
* Entropy change is inversely proportional to temperature. ∆S = \frac{ q_{rev} }{T}
* Spontaneous processes in an isolated system have a positive entropy.
* Entropy change for non-isolated systems, will be the sum of entropy of the system and the entropy of the surrounding.
* Entropy of a system at equilibrium is maximum; and there is no further change of entropy.
∆Stotal = ∆Ssystem + ∆Ssurr > 0
Gibbs Energy Change and Equilibrium
* Spontaneity of a reaction involves two thermodynamic properties – enthalpy and entropy.
* In 1870, J.Willard Gibbs an American mathematical physicist developed the concept of “free energy” to predict the spontaneity of a process.
* The following rules apply for predicting Gibbs energy for irreversible reactions:
• If at constant temperature and pressure, change in Gibbs energy, is negative, then the process is spontaneous.
• If change in Gibbs energy is positive, then the process is non-spontaneous.
• If change in Gibbs energy is equal to zero, then the process is at equilibrium.
* The following rules apply for predicting Gibbs energy for reversible reactions:
• If change in Gibbs energy at standard state, is negative, then the process is spontaneous.
• If change in Gibbs energy at standard state is positive, then the process is non-spontaneous.
• If change in Gibbs energy at standard state is equal to zero, the process is at equilibrium.