ThermodynamicsMind Map

Thermodynamic terms define the language of the chapter. A system is the part of the universe chosen for study, while surroundings include everything outside it. Systems may be open, closed or isolated depending on whether matter and energy can be exchanged. Properties such as pressure, volume, temperature and internal energy describe the state of a system. State functions depend only on initial and final states, whereas path functions such as heat and work depend on how the change occurs. Internal energy is the total microscopic energy stored in a system. Heat is energy transfer due to temperature difference, while work is energy transfer by force-displacement or pressure-volume change. Processes may be isothermal, adiabatic, isochoric or isobaric.

The First Law of Thermodynamics is the law of conservation of energy applied to thermodynamic systems. It states that energy can neither be created nor destroyed; it can only be converted from one form to another. For a chemical system, the change in internal energy equals heat supplied to the system plus work done on the system: ΔU = q + w. If heat is absorbed, q is positive; if heat is released, q is negative. If work is done on the system, w is positive; if the system expands and does work, w is negative. Enthalpy is defined as H = U + pV and becomes especially useful for constant-pressure reactions. For ideal gases, ΔH and ΔU differ by ΔngRT.

Calorimetry is the experimental measurement of heat exchanged during physical or chemical changes. The basic equation q = mcΔT connects heat with mass, specific heat capacity and temperature change. Heat capacity is the heat required to raise the temperature of a body by 1 K, while specific heat capacity is for unit mass. In chemistry, enthalpy change is the heat exchanged at constant pressure. Standard enthalpy changes are measured under standard conditions, usually 1 bar and specified temperature, commonly 298 K. Standard enthalpy of formation is the enthalpy change when one mole of a compound forms from elements in their standard states. Enthalpies of combustion, neutralization and solution are important reaction enthalpies frequently tested in NEET numericals.

Hess Law states that the enthalpy change of a reaction is the same whether the reaction occurs in one step or many steps, provided the initial and final states are the same. This works because enthalpy is a state function. Therefore, difficult reaction enthalpies can be calculated by adding, reversing or multiplying known thermochemical equations. Reaction enthalpy can also be calculated using standard enthalpies of formation or average bond enthalpies. Enthalpy of bond dissociation measures energy needed to break one mole of bonds in gaseous molecules. Enthalpy of atomization forms gaseous atoms from an element. Lattice enthalpy measures energy change during ionic crystal formation or separation. Born-Haber cycle applies Hess Law to ionic solids and connects sublimation, ionization, dissociation, electron gain and lattice enthalpy.

Spontaneity tells whether a process can occur on its own under given conditions, but it does not tell how fast it occurs. A spontaneous process may be fast, like acid-base neutralization, or slow, like rusting. Entropy measures randomness or dispersal of energy; it generally increases during expansion, mixing, melting, vaporization and formation of more gaseous particles. Total entropy of the universe increases for a spontaneous process. For chemical reactions at constant temperature and pressure, Gibbs free energy is the most useful criterion: ΔG = ΔH - TΔS. If ΔG is negative, the process is spontaneous; if positive, non-spontaneous; if zero, equilibrium. Gibbs energy is also related to equilibrium by ΔG° = -RT ln K, linking thermodynamics with chemical equilibrium.