Thermodynamics

• The “state” of a system is characterised by a set of variables such as P, V, T, n, etc., which are state variables (or state functions) and can be classified as either intensive or extensive.
• If temperatures are different, energy (heat) flows from high temperature to low temperature.
• All chemical reactions are accompanied by an energy change.
• The law of conservation of energy is often represented as ∆ U = q + w, where ∆ U is the change in internal energy, q is system heat, w is work, and U is internal energy.
• Temperature is a measure of a particle’s average kinetic energy, an indication of heat content, and objects at the same temperature do not exchange heat energy.
• The difference between the change in internal energy and the enthalpy is the amount of work that needs to be done to make room for the products of the reaction.
• Hess' Law states that ∆ H ° for an overall process is the sum of the ∆ H ° 's for the individual steps of the process.
• Bond Dissociation Enthalpy values are for GASES ONLY and all reactants & products must be in the gaseous state to use BDE values alone.
• Ionic bonds result from a complete transfer of an electron.
• Expansion work is negative since energy leaves the system as the system does work on the surroundings.
• Compression: work is done on the system by the surroundings.
• Pressure-Volume (or P-V) work is the most common type of work encountered in chemical systems.
• Work is the energy exchange that results when a force moves an object through a distance.
• Molar Heat Capacity (C m or C) is the amount of heat needed to raise 1 mole through 1C ° or 1 K unit, or the heat capacity per mole, units: J∙K –1.
• Specific Heat (C S or c) is the amount of energy to raise 1 gram through 1C ° or 1 K unit, or the heat capacity per gram of substance, units: J∙g –1.
• If heat is given off, the rxn is Exothermic (∆ H is negative).
• If heat is absorbed, the rxn is Endothermic (∆ H is positive).
• Hess's Law states that thermochemical equations can be added (or subtracted) to yield other thermochemical equations.
• The name for the reverse process of Enthalpy of Vaporization is Deposition.
• Lattice enthalpies are larger for short ion separations and/or larger ionic charges.
• The lattice energy of ionic compounds is determined by the electrostatic interactions between ions.
• Reactions at Constant Pressure, i.e., open to the atmosphere, allow the volume to change, hence, ∆ V can be ≠ 0.
• A state function for a system is independent of how that value was achieved and only depends on the difference between the final and initial state, not on the path it takes to get there.
• Reactions at Constant Volume, i.e., a sealed container, result in zero work and zero change in internal energy, hence, ∆ U = 0.
• When an ideal gas expands at constant temperature, ∆ U is zero.
• Both enthalpy (H) and internal energy (U) are state functions.
• If work done, such as pushing back the atmosphere, is considered, ∆ U = -P ext ∆ V.
• Entropy is a thermodynamic function that indicates whether or not a reaction will occur spontaneously.
• Spontaneous reactions always proceed so as to move towards a state of equilibrium, i.e., Q → K.
• The direction of spontaneity is determined by the sign of ∆G.
• Positional distribution of species in space is referred to as mixing of two gases or the expansion of a gas into a vacuum.
• For any substance, entropy rises with temperature due to more atomic/molecular motions and therefore more ways the system can contain energy.
• Thermal distribution of energy among species, or distribution of species over energy levels, is referred to as heat flow from hot to cold objects.
• A spontaneous process is one that occurs without any external influence.
• An instantaneous process is one that occurs almost instantly, that is very rapid.
• A spontaneous process may be fast or slow, so it is not generally true that a spontaneous process is also instantaneous.
• Expansion of a gas into a vacuum is not only spontaneous but also instantaneous.
• The entropies of the four gases are likely to be different because their molecular structures are different.
• If the reaction is spontaneous, ∆ G is negative.
• The number of moles of gases is the same on both sides of the equation, however, so the entropy change is likely to be small if the temperature is constant.