Thermodynamics Flashcards Preview

A2 Chemistry > Thermodynamics > Flashcards

Flashcards in Thermodynamics Deck (36):
1

Define enthalpy change.

The heat energy transferred in a reaction at constant pressure.

2

What are standard conditions?

100kPa

298K

3

What is the enthalpy change for an exothermic reaction?

Negative.

4

What is the enthalpy change for an endothermic reaction?

Positive.

5

Lattice enthalpy is a measure of what?

Ionic bond strength.

6

Define lattice formation enthalpy, 

The total enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions.

7

Define lattice dissociation enthalpy.

The total enthalpy chaneg when 1 mole of a solid ionic substance is completely dissociated into its gaseous ions under standard conditions.

8

Define enthalpy change of formation.

The total enthalpy change when 1 mole of a substance is formed from its constiuent elements which are in their standard states under standard conditions.

9

Define enthalpy change of atomisation of an element.

The total enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state.

10

Define first ionisation enthalpy.

The total enthalpy change when 1 mole of gaseous 1+ ions is formed from 1 mole of gaseous atoms.

11

Define first electron affinity.

The total enthalpy change when 1 mole of gasesou 1- ions is formed from 1 mole of gaseous atoms and 1 mole of electrons.

12

Define enthalpy of hydration.

The total enthalpy change when 1 mole of aqueous ions is forme from 1 mole of gaseous ions.

13

Define bond dissociation enthalpy.

The total enthalpy change when all the bonds of the same type in 1 mole of molecules are broken.

14

Define the enthalpy of atomisation of a compound.

The total enthalpy change when 1 mole of a compound in its standard state is converted to gaseous atoms.

15

Define second ionisation enthalpy.

The enthalpy change when 1 mole of gaseous 2+ ions and 1 mole of electrons is formed from 1 mole of gaseous 1+ ions.

16

Define second electron affinity.

The total enthalpy change when 1 mole of gaseous 2- ions is formed from 1 mole of gaseous 1- ions and 1 mole of electrons.

17

Define enthalpy change of solution.

The total enthalpy change when 1 mole of solute is dissolved in sufficient solvent that no further enthalpy change occurs on further dilution.

18

What is Hess's law?

The total enthalpy change of a reaction is independent of the route of reaction taken.

19

To find the lattice formation enthalpy for NaCl, what is the Born-Haber cycle that must be constructed.

First the -(enthalpy of formation for NaCl)

Then the atomisation enthalpies for both Na(s) and 1/2Cl2(g)

Then the first ionisation enthalpy for Na(g)

Then the first electron affinity of Cl(g)

20

Why is it that theoretical lattice enthalpies differ from experimental values?

Because the theoretical values assumet that the lattice is purely ionic - experimental values suggest that ionic compounds have a degree of covalent character.

21

Why is it that ionic lattices have a degree of covalency?

Because the positive ions polarise their negative neighbourds, withdrawing electron density from them - the more polarisation that occurs, the greater the withdrawn electron density, the greater the covalency.

22

How can the covalency of an ionic bond be established?

By comparing the experimental and theoretical values for the bond - the greater the difference, the more covalent the bond is.

23

Which are stronger, covalent or ionic bonds?

Covalent bonds - their formation is more exothermic than ionic bonds'.

24

What are the process that occur when an ionic lattice is dissolved?

The bonds between the ions break - an endothermic process (lattice enthalpy of dissociation)

New bonds between the ions and water form - an exothermic process (enthalpy of hydration)

25

Enthalpy change of solution = ?

Lattice enthalpy of dissociation + Enthalpy of hydration

26

Why is it that the data book value for a bond enthalpy can differ drom the actual enthalpy of a particular bond?

Because data book values are mean values across a range of compounds - actual bond strength varies between compounds and even within a compound. For an accurate bond enthalpy, acquire experimental data for the relevant compound.

27

Define entropy.

A measure of disorder in a system that defines how much work that system is able to do.

28

What variables affect the entropy of a substance?

The physical state of the substance (Solid = ordered, liquid = less so, gas = chaotic)

Whether the substance has been dissolved.

The number of particles.

29

What is a spontaneous reaction?

A reaction which occurs by itself.

30

What is the total entropy change of a reaction?

The entropy of the products minus the entropy of the reactants.

ΔSsystem = Sproducts - Sreactants

31

What is the entropy change of surroundings defined as?

ΔSsurroundings = -ΔH / T

Where: ΔH = Enthalpy change

T = temperature of surroundings

32

What is Gibbs' free energy used for?

Predicting the feasibility of a reaction.

33

When is a reaction spontaneous?

When the Gibbs' free energy is less than or equal to zero.

34

What is the equation for Gibbs free energy?

ΔG = ΔH - TΔS

G = Gibbs free energy

ΔH = Enthalpy change

ΔS = Entropy change

T = Temperature

35

How is it possible to calculate the temperature at which a reaction becomes feasible?

By setting the Gibbs free energy equal to zero - this occurs just as the reaction becomes feasible.

This gives:

ΔH/ΔSsystem = T

36