Thermodynamics Flashcards
Standard Enthalpy Change
Change in heat energy under standard conditions, i.e. 298K and 100kPa
Standard Enthalpy of Formation
Enthalpy change when one mole of a compound is formed from its constituent elements all in their standard states under standard conditions
Standard Enthalpy of Combustion
Enthalpy change when one mole of a substance is completely burnt in oxygen under standard conditions
Enthalpy of Atomisation
Enthalpy change when one mole of gaseous atoms is formed from an element in its standard state
First Ionisation Energy
Enthalpy change when one mole of gaseous atoms each lose an electron to form one mole of gaseous ions with a single positive charge
First Electron Affinity
Enthalpy change when one mole of gaseous atoms each acquire an electron to form one mole of gaseous ions with a single negative charge
Lattice Enthalpy of Formation
Enthalpy change when one mole of solid ionic compound is formed from its gaseous ions infinitely far apart
Lattice Enthalpy of Dissociation
Enthalpy change to separate one mole of solid ionic compound into its gaseous ions infinitely far apart
Enthalpy of Hydration
Enthalpy change when one mole of gaseous ions form aqueous ions
Enthalpy of Solution
Enthalpy change when one mole of an ionic compound completely dissolves in sufficient water to form an infinitely dilute solution
Mean Bond Enthalpy
Enthalpy change when breaking one mole of a given bond in a molecule of a gaseous species (averaged over different molecules)
Entropy
A measure of disorder of a system
Feasible/Spontaneous
- reactions that occur on their own accord
- free energy is zero or negative
Gibbs Free Energy
Thermodynamic quantity that combines enthalpy and entropy under a constant temperature and pressure to determine the spontaneity of a reaction (direction it takes)
First Law of Thermodynamics
Energy can be neither created nor destroyed, only converted into different forms
Second Law of Thermodynamics
Entropy of an isolated system not at equilibrium will tend to increase over time
Third Law of Thermodynamics
Entropy of a perfect crystal at absolute zero is zero
Suggest why ionic model and actual lattice enthalpies are different
- Ionic model assumes ions are spherical and arranged in a lattice
- Calculated value is more negative suggesting stronger attraction
- Compound has covalent character
Suggest what can be deduced from comparisons between theoretical and actual lattice enthalpies
Similar = Purely Ionic Different = Covalent Character
Which types of ionic compounds have covalent character
- high charge
- small positive ion
- large negative ion
= negative ion polarised > more sharing > covalent character
Suggest why successive ionisation energies get larger
- negative electron is removed from a positive ion
- requires more energy than removing from an atom
- removing electrons from orbitals closer to the nucleus require much more energy
Suggest why first electron affinity is negative but second is not
- first electron affinity is exothermic since energy is released when an electron is added to an atom due to attraction
- second electron affinity is endothermic since energy is required to overcome repulsion from negative ion
How to determine Enthalpy of Solution
- difficult to measure directly
- calculated using lattice dissociation enthalpy and enthalpy of hydration
Suggest why Enthalpy of Hydration is always negative
- (hydrogen) bonds are always formed between ions and water
- bond making is exothermic
Suggest what factors result in the most negative Enthalpy of Hydration
- highly charged
- small ions
(high charge density)
Suggest why calcium fluoride is most likely to be found as CaF2 rather than CaF or CaF3
- Ca2+ has a higher charge density than Ca+ so formation of CaF2 is more exothermic so more favourable
- Third ionisation energy of Ca is much higher than first and second so formation of CaF3 is endothermic so less favourable than CaF2
Why is Gibbs free energy zero at equilibrium
Forward and backward reaction occur at the same rate at the same time
What is the value for Gibbs free energy when a reaction is feasible
zero or negative
When is a reaction first feasible
When Gibbs free energy is zero
When does entropy increase
- moles of gas increase
- complexity increases
- more product particles
Suggest why calculating the lattice enthalpy of sodium carbonate would be more difficult than sodium oxide
Carbonate ion is complex and calculating lattice enthalpy would require gaseous ions
Suggest the purpose of constructing a Born Haber cycle
Calculating enthalpy changes that cannot be measured directly, e.g. Lattice Enthalpy or Enthalpy of Solution
Born Haber Method for Lattice Enthalpy
Start with Enthalpy of Formation
Born Haber Method for Enthalpy of Solution
Start with Lattice Enthalpy
Explain why the entropy is zero at 0K
- absolute zero = minimum energy
- no molecular motion
- no entropy / randomness
Suggest why the gradient of a free energy-temperature graph changes below a certain temperature
Compound changes state (condenses/freezes)
Suggest why magnesium is stable in air at room temperature despite the formation of MgO having a negative free energy value
- protective layer of MgO prevents further attack on Mg
- slow reaction
- high activation energy
Partial Pressure
Pressure that would be exerted by one of the gases in a mixture if it occupied the same volume on its own
= mole fraction x total pressure
Mole Fraction
number of moles of gas/ total moles of gases
Suggest why free energy is zero during a change in state
- forward and backward reaction occur at same rate
- at equilibrium
Suggest the condition under which diamond would have an entropy of zero
absolute zero/0K
Suggest why a feasible reaction does not take place in the absence of a catalyst
- high activation energy
- catalyst lowers activation energy by providing an alternative route for reaction
Suggest what is meant by perfect ionic model
- ions are spherical and arranged in lattice
- only electrostatic attraction between ions/ no covalent character
Describe what occurs when free energy is positive for a reversible reaction
reaction proceeds in reverse direction
