Chap-06 -Chemical Thermodynamics

https://docs.google.com/document/d/1T0XC3RR4Xgvd_P16HG3uyRnhbhZaebvo/edit?usp=sharing&ouid=109474854956598892099&rtpof=true&sd=true intensive properties, state functions, types of processes. First law of thermodynamics - Concept of work, heat internal energy and enthalpy, heat ca pacity, mola r heat capacity; Hess’s law of consta nt heat summation; Enthalpies of bond dissociation, combustion, formation, atomization, sublimation, phase transition, hydration, ionization and solution. Second law of thermodynamics- Spontaneity of processes; ΔS of the universe and ΔG of the system as criteria for spontaneity, ΔG º (Standard Gibbs energy change) and equilibrium constant. TYPES OF SYSTEM System and surrounding interact by (i) Matter and (ii) Energy. Depending on the interactions between system and its surroundings, three kinds of system are distinguishable. Isolated system : A system which neither exchanges energy nor matter with its surroundings is called an isolated system. For example a liquid in contact with its vapour in an insulated closed vessel is an isolated system. Closed system : A system which may exchange energy but not matter with its surroundings is called a closed system. If in the above example the vessel containing liquid in contact with its vapour is closed one but is not insulated. It is a closed system. Open system : A system which may exchange both energy and matter with its surrounding is called an open system. The evaporation of water in an open vessel is an example of open system. Water absorbs heat from the surroundings and escapes into surroundings as water vapour. STATE VARIABLES A thermodynamic system has to be macroscopic (i.e. of sufficiently large size); this facilitates measurement of its properties such as pressure, volume, temperature, composition and density, such properties are therefore called macroscopic or bulk properties. These are also called state or thermodynamic variables. CHAPTER INCLUDES Types of System State variables Extensive and Intensive variables Internal Energy First Law of thermodynamics Enthalpy Hess's Law of constant heat summation Entropy & Entropy Change Gibb's free energy & spontaneity of reaction For example PV = nRT (an ideal gas equation) P = nRT V Here P is the dependent variable while n, T and V are independent variables. The state of a system is defined when the state variables have definite values EXTENSIVE AND INTENSIVE VARIABLES These are the observable properties of a system. Extensive property of a system are those whose magnitude depends on the amount of matter present in the system. It is an additive property. For example volume, internal energy, enthalpy, entropy, mass etc. The property whose magnitude does not depend upon quantity of matter present in a system is known as intensive property of the system. It is not an additive property and don’t alter with time. Example are; temperature, pressure, density etc. INTERNAL ENERGY (U) It is the energy associated with a system by virtue of its molecular constitution and motion of its molecule. Such motion may be translational, rotational, vibrational etc. U = E total = Et + Er + Ev + Eb + Ee + Eo + VT + ENuc. where Et = Translational kinetic energy K.E. Er = Rotational K.E. Ev = Vibrational K.E. Eb = Bond energy Ee = Electronic energy E = Zero point energy i.e., energy at absolute zero (–273oC) VT = Interamolecular attraction energy ENuc = Nuclear energy Internal energy is a state function and an extensive property. Sign of heat (q) (1) +q indicates that heat is absorbed by the system. (2) –q indicates that heat flows out of the system. Sign of work (w) When work is done on the system, w = +ve When work is done by the system, w = –ve Work of Expansion of Gas W = –PΔV → for irreversible isothermal process. V2 W = –2.303nRT log V1 → for reversible isothermal process. FIRST LAW OF THERMODYNAMICS Statements Energy can be transformed from one form to another but energy can neither be created nor destroyed. The total mass and energy of an isolated system remains constant There is an exact equivalence between work and heat Whenever a quantity of one kind of energy is produced, an exactly equivalent amount of another kind of energy must disappear Mathematical formulation q = ΔE + w Another form for isochoric process. q = ΔE, ΔW = 0 for adiabatic process ΔE = –w, q = 0 for cyclic process q = w, ΔE = 0 ENTHALPY For the process or a chemical reaction carried out at constant volume, the heat absorbed/evolved is equal to the corresponding change in internal energy (ΔU) or (ΔE). Many of the reactions are carried out at constant pressure. To measure heat changes at constant pressure it is useful to define new state function called enthalpy (H). i.e., internal energy and PV energy of any system under a particular set of conditions is known as enthalpy (H) i.e., H = U + PV and for reaction i.e. process Δng = number of moles of gaseous product – number of moles of gaseous reactant. Standard Enthalpy Changes ΔH° S. No. Name of Enthalpy Changed Defination Example 1. Enthalpy of combustion ΔH° C It is the energy which is released when 1 mole of substance is burnt completely with required amount of oxygen C + O2 → Co2 CH4 + 2O2 → CO2 + 2H2O 2. Enthalpy of formation ΔH° f It is the energy which is released when 1 mole of substance/ compound i s formed from its elements in their most stable state. H2 + S + 2O2 → H2SO4 S + O2 → SO3 3. Enthalpy of atomisation ΔH° a It is the energy which is absorbed when 1 mole of substance broken up into its isolated atoms in the gas phase (i) Na(s)→ Na(g) 4. Enthalpy of neutralization ΔH° N It is the energy which is released (fix for strong acid - Strong base neutralization i.e. -57.1 kJ or –13.7 kcal) when one mole of water is formed by the neutralization of an acid by a base. CH3COOH + NaOH → CH3COONa + H2O H2SO4 + 2NaOH → Na2SO4 + 2H2O 5. Enthalpy of fusion ΔH° f It is the energy which is absorbed when 1 mole of liquid is formed from the solid without a change in temperature. (i) H2O( l )→ H2(g) HESS’S LAW OF CONSTANT HEAT SUMMATION If a reaction is the sum of two or more constituent reactions, then ΔH for the overall process must be sum of the ΔH of the constituent reactions. The enthalpy change for a reaction is the same whether it occurs in one step or in a series of steps. A ΔH1 B ΔH2 ΔH3 C For A → D, ΔH = ΔH1 + ΔH2 + ΔH3. ENTROPY AND ENTROPY CHANGE Entropy is measure of randomness. Higher the disorder higher is the entropy. Entropy for solid < liquid < gas ; where qrev = Heat absorbed when the process is carried out reversibly and isothermally. Entropy change is given in unit of J/K or JK–1. GIBB'S FREE ENERGY AND SPONTANEITY OF REACTION Free energy of a system is defined as the maximum amount of energy available to a system that can be converted into useful work during a process. In other words, it is a measure of capacity of a system to do useful work. It is denoted by symbol G and is mathematically given by G = H – TS where H is the enthalpy of the system, S is its entropy and T is the absolute temperature. If the process is carried out at constant temperature and pressure, the terms Δ(PV) and Δ(TS) become Δ(PV) = PΔV and Δ(TS) = TΔS ΔG = ΔE + PΔV – TΔS or Spontaneity of the reaction ΔH ΔS ΔG = ΔH – TΔS Reaction Spontaneity – + – Spontaneous at all temperature – – – or + Spontaneous at low temperatures where ΔH outweighs TΔS (TΔS < ΔH) Nonspontaneous at high temperatures where TΔS outweighs ΔH + – + Nonspontaneous at all temperatures + + – or + Spontaneous at high temperatures where TΔS outweighs ΔH (TΔS > ΔH) Nonspontaneous at low temperatures where ΔH outweighs TΔS ❑ ❑ ❑