The study of the relationships between heat, work, temperature, and energy is known as thermodynamics. The rules of thermodynamics define how energy evolves in a system and whether it can do beneficial work on its surroundings.
System: The region of the cosmos that is being seen is referred to as a system.
Surroundings: Everything in the cosmos other than the system is referred to as surroundings. The Universe is made up of the System and its surroundings.
Open System: A system in which energy and matter are exchanged with the environment.
An open system, for example, is the presence of reactants in an open beaker.
Closed System: When there is no interchange of matter yet there is an exchange of energy, a system is said to be closed.
The existence of reactants in a closed vessel consisting of conducting material, for example.
Isolated System: An isolated system is one in which there is no exchange of energy or matter with the environment.
An isolated system is, for example, the existence of reactants in a thermoflask or a material in an insulated closed vessel.
Homogenous System: When all of the elements of a system are in the same phase and uniform throughout, it is said to be homogeneous.
For example: A- mixture of two miscible liquids.
Heterogeneous system: When a combination consists of two or more phases and the composition is not uniform, it is said to be heterogeneous.
A combination of insoluble solids in water, for example.
A thermodynamic system's state refers to its macroscopic or bulk qualities that may be characterised using state variables:
Pressure (P), volume (V), temperature (T), and quantity (n), among other things.
Isothermal process: An operation is considered to be isothermal if it takes place at a constant temperature. dT = 0 for an isothermal process, where dT is the temperature change.
Adiabatic process: It is a process in which there is no heat transmission between the system and its surroundings.
Isobaric process: A process is considered to be isobaric if it is carried out under constant pressure. dP = 0 in this case.
Isochoric process: When a process is carried out at a constant volume, it is referred to as isochoric.
Cyclic process: A process in which a system goes through a sequence of modifications before returning to its original condition.
Reversible Process: When a change is introduced into a process in such a way that the process might be reversed at any time by an infinitesimal modification. Reversible refers to a change that can be reversed.
It's the total of all the many types of energy that a system can have.
It is symbolised by AM in thermodynamics, which can vary when
– Heat enters or exits the system.
— Work is done on the system or by it.
— Substance enters or exits the system.
Internal Energy Changes as a Result of Work
Let us endeavour to bring about a shift in internal energy.
Assume that the system's starting condition is A and that the temperature is TA. Internal energy = uA
The new state is termed state B, and the temperature is TB, after conducting some mechanical work. It is discovered to be
TB > TA
The internal energy after change is denoted by uB.
∴ Δu = uB – uA
Internal Energy Changes as a Result of Heat Transfer
The internal energy of a system may be adjusted without doing work by transferring heat from the environment to the system.
Δu = q
The heat absorbed by the system is denoted by q. It may be calculated as a difference in temperature.
When heat is transmitted from the environment to the system, q is positive. When heat is transmitted from the system to the environment, q is -ve.
When a change of state is accomplished by both effort and heat transmission.
Δu = q + w
Thermodynamics' first law
Energy cannot be generated or destroyed, according to this theory. An isolated system's energy remains constant.
Δu = q + w.
Change in enthalpy: At constant pressure, change in enthalpy is the heat absorbed or emitted by the system.
ΔH = qp
When an exothermic reaction occurs (System loses energy to Surroundings),
ΔH and qp both are -Ve.
In case of an endothermic reaction (System absorbs energy from the Surroundings).
ΔH and qp both are +Ve.
Capacity for heat
"If the temperature is kept constant, the total quantity of heat developed or absorbed in a reaction is the same whether the reaction takes place in a single step or in a series of stages."
Hess's Law in Practice:
H2(g) + 436 kJ mol-1 → 2 H(g)
Bond energy is the bond dissociation energy of diatomic compounds like H2.
Common salt dissolves in water
Water evaporation in an open vessel
The downward flow of water
The transfer of heat from one end to the other.
a candle is lit (initiation by ignition)
CaCO3 is heated to produce CaO and CO2.
The water is flowing uphill.
The transfer of heat from a cold to a warm body.
The Spontaneous Process's Driving Force: