THERMODYNAMICS - PART 1 || All Concepts, Tricks & PYQ || Ummeed NEET

Brief Summary

This video provides a comprehensive overview of thermodynamics and thermochemistry, covering fundamental concepts, definitions, and various types of thermodynamic processes. It explains key terms like ideal gas, system, surroundings, and different types of walls and systems. The lecture also discusses state functions, path functions, internal energy, heat, work, and the first law of thermodynamics. Furthermore, it covers enthalpy, heat capacity, Hess's law, and different types of enthalpies, including bond energy and enthalpy of atomization.

Thermodynamics and thermochemistry fundamentals
System types and properties
Thermodynamic processes and functions
Enthalpy and heat capacity
Hess's law and enthalpy calculations

Introduction

The lecture begins with a welcome and an introduction to the chapter on thermodynamics and thermochemistry, highlighting its importance in both chemistry and physics. The instructor expresses hope that students will be able to answer at least three questions in the NEET paper from this chapter, aiming for a perfect score.

Ideal Gas

The discussion starts with ideal gases, focusing on energy changes during expansion, compression, or chemical reactions. Thermodynamics helps determine whether a reaction will occur but does not provide information about the reaction's speed, which is covered in chemical kinetics.

System

A system is defined as the part of the universe where a reaction is taking place, such as a container with an ideal gas. Everything outside the system is considered the surroundings. The combination of the system and surroundings constitutes the universe.

Surrounding

The surrounding is everything outside the system. The universe is the combination of the system and the surrounding.

Types of Walls

There are two types of walls: diathermic and thermally insulated. Diathermic walls allow energy exchange between the system and surroundings, while thermally insulated walls do not. Energy exchange can occur through heat or work. Work is done when the boundary is movable, such as a piston moving up or down.

Types of System

Three types of systems are defined: open, closed, and isolated. An open system exchanges both energy and matter with the surroundings. A closed system exchanges energy but not matter. An isolated system exchanges neither energy nor matter. A thermos flask is given as an example of an isolated system.

State of a System

The state of a system is its condition, which can be described using four state variables: pressure (P), volume (V), temperature (T), and number of moles (n). Knowing these variables provides knowledge of the system's state. The ideal gas law, PV = nRT, can be used to find one variable if the other three are known.

Properties of a System

Systems have two types of properties: extensive and intensive. Extensive properties depend on the amount of matter in the system (e.g., mass, volume, energy), while intensive properties do not (e.g., temperature, refractive index, molar mass). Concentration terms like molarity, molality, and mole fraction are intensive properties.

Types of Processes

Several types of processes are defined: isothermal (constant temperature), adiabatic (no heat exchange), isobaric (constant pressure), isochoric (constant volume), and cyclic (initial and final states are the same). Reversible processes occur infinitely slowly, maintaining equilibrium, while irreversible processes are fast and real.

State Function

State functions depend only on the initial and final states of the system, not on the path taken (e.g., pressure, volume, temperature, internal energy, enthalpy, entropy, and Gibbs free energy). Small changes in state functions are denoted by "d," while large changes are denoted by "delta."

Path Function

Path functions depend on the path taken during a process (e.g., heat and work). Small changes in path functions are denoted by "del."

Internal Energy

Internal energy (U or E) is the sum of all energies within a system, including nuclear, rotational, vibrational, and electronic energies. It depends on temperature and, for non-ideal gases, also on volume. For ideal gases, it depends only on temperature. Energy absorbed by the system is positive, while energy released is negative.

Heat

Heat (q) is a mode of energy transfer between the system and surroundings. The amount of heat is calculated using the formula q = m * c * delta T, where m is mass, c is specific heat, and delta T is the change in temperature. Heat absorbed is positive, while heat released is negative.

Work

Work (w) is another mode of energy transfer between the system and surroundings. The formula for work done is w = -P_ext * delta V, where P_ext is external pressure and delta V is the change in volume. For reversible processes, w = -nRT * ln(V2/V1). Work done on the system is positive, while work done by the system is negative.

First Law of Thermodynamics

The first law of thermodynamics states that energy is conserved: delta U = q + w. For isothermal processes, delta U = 0, so q = -w. For adiabatic processes, q = 0, so delta U = w. For isochoric processes, w = 0, so delta U = q_v. For isobaric processes, q_p = delta H.

Enthalpy

Enthalpy (H) is defined as H = U + PV. The change in enthalpy (delta H) is equal to the heat change at constant pressure (q_p). Delta H = delta U + P * delta V. Delta H is positive for endothermic reactions and negative for exothermic reactions.

Heat Capacity

Heat capacity (C) is the amount of heat required to raise the temperature of a substance by 1 Kelvin. Specific heat capacity is for 1 gram of a substance, while molar heat capacity (Cm) is for 1 mole. Cp is heat capacity at constant pressure, and Cv is heat capacity at constant volume. Cp is always greater than Cv. For ideal gases, Cp - Cv = R.

Poison’s Ratio

Poisson's ratio (gamma) is the ratio of Cp to Cv (gamma = Cp/Cv). For monoatomic gases, gamma = 1.66; for diatomic gases, gamma = 1.4; for triatomic or polyatomic gases, gamma = 1.33. For mixtures of gases, the average Cp and Cv are used to calculate gamma.

Isothermal Reversible Expansion

In an isothermal reversible expansion, the work done is w = -nRT * ln(V2/V1). Since delta T = 0, delta U = 0 and delta H = 0.

Isothermal Irreversible Expansion

In an isothermal irreversible expansion, the work done is w = -nRT * (1 - P2/P1).

Free Expansion of an Ideal Gas

In the free expansion of an ideal gas, the external pressure is zero, so no work is done (w = 0). Therefore, delta U = 0 and delta H = 0.

Isochoric Vs Isobaric Process

In an isochoric process, the volume is constant, so no work is done (w = 0). Therefore, delta U = q_v. In an isobaric process, the pressure is constant, and q_p = delta H.

Isothermal Vs Adiabatic Expansion

In isothermal expansion, PV = constant. In adiabatic expansion, PV^gamma = constant. The slope of the adiabatic curve is steeper than that of the isothermal curve.

Adiabatic Expansion Continues

Formulas for adiabatic processes are provided, including the relationships between temperature, volume, and pressure.

Hess Law

Hess's law states that the enthalpy change for a reaction is the same whether it occurs in one step or multiple steps. This is because enthalpy is a state function.

Laws of Thermochemistry

If two reactions are added, their delta H values are also added. If a reaction is reversed, the sign of delta H is also reversed. If a reaction is multiplied by a stoichiometric coefficient, the value of delta H is also multiplied by the same coefficient.

Different types of Enthalpies

Different types of enthalpies are discussed, including enthalpy of formation, standard enthalpy of formation, enthalpy of combustion, enthalpy of hydrogenation, enthalpy of hydration, enthalpy of solution, enthalpy of ionization, enthalpy of allotropic transformation, enthalpy of fusion, enthalpy of vaporization, and enthalpy of sublimation.