Enthalpy, Rates and equilibrium, Equilibrium 2, Enthalpy and Entropy

Question 1

Enthalpy definition

Answer 1

A measure of the heat energy in a system.

Question 2

Law of conservation of energy

Answer 2

Energy cannot be created or destroyed, only transferred. The total energy of a chemical system remains constant.

Question 3

Activation energy

Answer 3

Minimum amount of energy required for a reaction to occur

Question 4

Standard conditions

Answer 4

Standard pressure- 100kPa
Standard temperature- 25°C (25+273= 298K)
Standard concentration- 1mol/dm³

Standard state- physical state of a substance under standard conditions.

Question 5

Standard enthalpy of reaction

Answer 5

The enthapy change that accompanies a reaction in the molar quantities shown in the equation in standard conditions, with all reactants and products in their standard states.

Question 6

Standard enthalpy of formation

Answer 6

The enthalpy change when *1 mole of a substace** is formed from its constituent elements under standard conditions, with all reactants and products in their standard states.

Question 7

Enthalpy change of combustion

Answer 7

Enthalpy change when 1 mole of a substance reacts completely with oxygen, under standard conditions, with all reactants and products in their standard states.

Question 8

Enthalpy change of neutralisation

Answer 8

The enthapy change when 1 mol of water is formed in a neutralisation reaction between an acid and a base, in standard conditions, with all reactants and products in their standard states.

Question 9

Why can calorimetry be inaccurate?

Answer 9

• heat loss to surroundings
• incomplete combustion of fuel
• evaporation of alchohol (fuel) from wick
• not in standard conditions

All apart from last cause results to be less exothermic.
Look at flashcards to see how to draw graph for enthalpy of neutralisation.

Question 10

Average bond enthalpy

Equation

Why only an average?

Answer 10

The enthalpy change to break 1 mole of a specified bond in a gaseous molecule.
Always endothermic.

Total(bond enthalpy in reactants)- total(bond enthalpy in products)

Bond enthalpy found from the mean enthalpy of the same bond in different molecules.

Question 11

Hess’s Law

Answer 11

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

Question 12

Rate of reaction definition

Answer 12

How fast a reaction is being used up/ change in concentration over time.

Rate greatest at start, as reactant at highest concentration.
Slows down, as reactants are being used up.
Reactants completely used up so rate is 0.

Question 13

Factors affecting rate of reaction

Answer 13

Concentration (or pressure when reactants are gas)

Temperature

Use of a catalyst

Surface area

Question 14

How does each factor affect the rate of reaction?

Answer 14

Concentration/pressure - more molecules per unit/volume, more successful collisions.
Temperature- greater kinetic energy, more successful collisions, more with activation energy
Catalysts- provides an alternate reaction pathway with lower activation energy so more particles have activation energy, more successful collisions.
Surface area- more area for particles to collide with.

Question 15

What makes a collision effective

Answer 15

Particles should collide with correct orientation
Particles must have enough energy to overcome activation energy barrier.

Question 16

Catalysts definition

Types- method

Advantages of catalysts.

Answer 16

Substance that increases the rate of reaction without being used up in reaction. Provided an alternate reaction pathway, with lower activation energy.

Homogenous- when all reactants are in same phase
Heterogenous- reactants are in different phases

Homogenous:
Reactant react with catalyst to form intermediate. Intermediate breaks down to form product and regenerates catalyst.

Heterogenous:
Reactant molecules are adsorbed (weakly bonded) onto surface of catalyst, where reaction occurs. Product leaves catalyst by desorption.

Reduces energy demand and lowers temperatures- reduces energy costs and increases sustainability by reducing combustion of fossil fuels, so reduction of CO2 emmissions.

Increases yield of desired products with fewer by-products.

Question 17

Examples of catalysts

Answer 17

Homogenous:
Ozone depletion, cl. radical 2O3<–>3O2

Heterogenous:
Hydrogenation of alkene, Ni catalyst
Haber process, Fe catalyst

Question 18

Boltzman distribution curve
- increasing temperature
- using a catalyst

Answer 18

Shows the spread of energies (Fraction of molecules with energy vs Energy)
Curve not symmetrical as particles with very high energies skew the results to the right

• more particles at higher energies, more particles with activation enery, more successful collisions
• More particles with energy more than the activation energy, so frequency of successful collisions increases

Question 19

Le Chatelier’s principle

Answer 19

If an external change occurs to a reaction at equilibrium, the position of equilibrium will move to oppose the change.

Question 20

Dynamic equilibrium is reached when

Answer 20

• In a closed system
• Forward and backward rate are equal
• The concentration of products and reactants are constant

Question 21

Increasing pressure

Answer 21

Moves to least moles of gas.

Question 22

Enthalpy of atomisation

Answer 22

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

Question 23

First ionisation energy

Answer 23

The energy required to remove 1 electron from each atom in 1 mole of gaseous atoms, to form 1 mole of gaseous 1+ ions.

Question 24

First electron affinity

Answer 24

The enthalpy change when 1 electron is added to each atom in 1 mole of gaseous atoms, to form 1 mole of gaseous -1 ions.

Question 25

Lattice enthalpy

Answer 25

Enthalpy change when 1 mole of an ionic crystal lattice is formed from its constituent ions in gaseous form.

Measure of ionic bond strength.

Question 26

Enthalpy of hydration

Answer 26

Enthalpy change when 1 mole of gaseous ions dissolve in H2O to become 1 mole of aqeous ions.

Question 27

Enthalpy of solution

Answer 27

Enthalpy change when 1 mole of a solute completely dissolves in water.

Question 28

The more exothermic lattice enthalpy is,

Answer 28

The more exo, the stronger the ionic bonds

Question 29

Factors affecting lattice enthalpy/ enthalpy of hydration

Answer 29

• ionic radius increases, distance between ions increases, attraction decreases, lattice enthalpy less exo
• ionic charge increases, attraction between ions increases, more exo

(in enthalpy of hydration, its attraction between ions and water)

Question 30

Entropy
Factors affecting entropy

Why does Ne(g) have greater entropy than Ne(l)

Units and equation for standard entropy

Answer 30

The dispersal of energy/ the level of disorder in a chemical system
Entropy increases when: increase in number of molecules, changes state (s–> l –> g)

In gas state atoms are more disordered, atoms have greater energy

J/K/mole
standard entropy= entropy of products- entropy of reactants

Question 31

Feasibility definition

Answer 31

Whether a reaction is likely to happen and is energetically feasible (spontaneous)
Free energy change must be -ve.

Question 32

Gibb’s Free energy change

Answer 32

ΔG = ΔH – TΔS

free energy change ΔG= kJ/mol
enthalpy change ΔH= kJ/mol
T= K
Standard entropy change ΔS= J/K/mole (convert to kJ/K/mol by dividing by 1000)

Question 33

Free energy change during change of state

Answer 33

ΔG = 0

y=mx+c
ΔG = (-ΔS)T + ΔH

Energy limitations: Some reactions have -ve ΔG, but still don’t occur:
- reaction rate may be very slow
- reaction may require large activation energy