thermodynamics

energy

the capacity to do work or transfer heat

heat

the form of energy that flows between two objects because of their difference in temperature

work is done

as a result of motion against an opposing force

system

specific part of universe under study

surroundings

everything in the universe outside the system

law of conservation of energy

• The total energy is unchanged in a chemical reaction

• The total energy of the universe is constant

∆E

∆E system + ∆E surroundings = 0

exothermic

releases heat

endothermic

takes in heat

diathermic

when container walls allow the passage of heat energy

adiabatic

when walls do not allow the passage of heat energy

quantity of heat transferred to or from an object depends on

• The size of temperature change

• The nature of the material gaining or losing heat

• The amount of material grams or moles

heat capacity

is the heat energy required to change the temperature of a substance by 1K

specific heat capacity

the amount of heat (q) required to change the temperature of a given amount (g) of material by 1K, (units J/g/K)

molar heat capacity

the amount of heat required to change the temperature of one mole of material by 1 Kelvin (units = J/mol/K)

enthalpy

the heat change transferred at constant pressure by a chemical reaction or process

state function

Change in value depends only on the initial and final state.

sublimation

changing directly from a solid to a gas

latent heat

the absorption or release of heat energy without a change in temperature

endothermic - more energy required to break bonds than

gained in bond formation

standard enthalpy of formation for an element in its standard state

is zero

standard enthalpy of a solution

the enthalpy change when 1 mole of a substance dissolves in a large excess of pure solvent at 1 bar pressure

internal energy

the sum of the potential energy (PE) and kinetic energy (KE).

Internal energy (U) = PE + KE

internal energy depends on

1. Amount of particles (molecules or atoms)

2. Types of particles

3. Temperature

external KE

the energy associated with the overall motion of an object

internal KE

the energy associated with the random motion of particles within an object

transfer of energy to or from a system

1. Heat (or thermal energy) absorbed or lost (q), energy transferred as a result of T difference only

2. Work performed (w), energy transferred when an object is moved by a force. Work relates to the movement of an object against a force.

forms of work

• Electrical (battery, electrolysis)

• Volume change (in chemistry usually expansion)

first law of thermodynamics

For any system the change in internal energy (∆U) is equal to the energy transferred in the form of heat (q) and work (w)

isothermal irreversible gas expansion

Each step has to be calculated and then added up. The more steps between the final and initial states the more work is done - the more the gas cools

isothermal reversible gas expansion

This means that the heat transfer takes place over an infinite number of steps so that gas pressure equals external pressure throughout

path of maximum work

reversible

spontaneous change

once started, proceeds on its own without continuous external influence.

At 0 degrees Celsius ice spontaneously absorbs heat energy from the surroundings to form liquid water

entropy

a measure of molecular and energy dispersion - it’s the distribution between different states, it is denoted by the symbol S

entropy increases

increase in dispersion

entropy decrease

decrease in dispersion

second law of thermodynamics

we cannot consider the system as an isolated entity. All spontaneous processes result in an increase in entropy of the universe

third law of thermodynamics

At 0K a perfect crystal has S=0, there is no variation in a perfect crystal at 0K

rules for predicting whether entropy will increase in a reaction

1. Reactions in which larger molecules are broken down into smaller components - de-polymerisation.

2. Reactions in which there is an increase in the number of moles of gas

3. Processes in which solids change into liquids or liquids into gases