enthalpy, reaction rates and equilibria

Question 1

enthalpy definition

Answer 1

the total internal energy inside a chemical system
- this includes thermal and chemical energy

Question 2

chemical system definition

Answer 2

all the atoms, ions and molecules that make up all the chemicals in a space

Question 3

why does the enthalpy of a chemical system change during reactions

Answer 3

when a reaction takes place, energy is transferred between the system and its environment, causing enthalpy changes

Question 4

exothermic definition

Answer 4

when products have less energy than reactants, therefore energy is given off in the form of heat (this can be measured with a thermometer)
- overall enthalpy decrease

Question 5

endothermic definition

Answer 5

when products have more energy than reactants so energy is taken into system in the form of heat
- overall enthalpy increase

Question 6

what are standard conditions
-temp
-pressure
-concentration
-state

Answer 6

temp - 298K / 25 C
pressure - 100 kPa
concentration - 1 moldm-3
state - the physical state of a substance under standard conditions

Question 7

enthalpy change of formation definition

Answer 7

the enthalpy change that occurs when 1 mole of a substance is formed from its raw elements under standard conditions and states

remember the enthalpy change of formation of a raw element is always 0

Question 8

enthalpy change of combustion definition

Answer 8

the enthalpy change that occurs when 1 mole of a substance burns completely in excess oxygen under standard conditions

Question 9

enthalpy change of neutralisation definition

Answer 9

the enthalpy change that occurs when 1 mole of water is formed in a reaction between an acid and a base under standard conditions

Question 10

what are the units for enthalpy change

Answer 10

kJmol-1

Question 11

activation energy definition

Answer 11

the minimum amount of energy needed for a reaction to take place

Question 12

enthalpy change formula

Answer 12

= H(products) - H(reactants)

Question 13

bond enthalpy definition
(also known as bond energy/bond dissociation energy)

Answer 13

the amount of energy needed to break and separate 1 mole of a specific bond in gaseous molecules so that the resulting gaseous particles exert no force upon each other, under standard conditions

Question 14

what is the difference between bond enthalpy and average bond enthalpy

Answer 14

bond energies are affected by other atoms in the molecule

average bond enthalpy is an average for many bonds taken from a wide range of compounds containing the bond

e.g. an O-H bond in water will have a different bond enthalpy to an O-H bond in methanol

Question 15

why might listed average bond enthalpy for a specific bond differ from the average measured bond enthalpy of a specific bond in a specific molecule

Answer 15

the listed figure is an average taken from many compounds with that specific bond

the measured value is specific to the bond within a specific molecule, although this is also an average of breaking many of this type of bond

Question 16

how do you determine bond enthalpies

Answer 16

bond enthalpies cannot be determined directly so enthalpy cycles are used to calculate the average

bond breaking = exothermic (-) as energy is needed to break bonds
bond forming = endothermic (+) as energy is released when they are made so

enthalpy change of a reaction = (+ enthalpy of bonds broken) + (- enthalpy of bonds formed)

Question 17

why might calculations for the same bond enthalpy from 2 different reactions produce different answers

Answer 17

the other bond enthalpies are averages so they will differ from the actual enthalpies involved in the reactions

Question 18

why does enthalpy change of combustion of alkanes become more exothermic as chain length increases

Answer 18

as chain length increases there are more atoms and bonds burned so more energy is released, so more exothermic

Question 19

what is hess’ law

Answer 19

the idea that products can be formed directly from raw elements or indirectly from raw elements

Question 20

outline collision theory

Answer 20

collision theory states that for a chemical reaction to take place, the particles need to collide with each other in the correct orientation AND with sufficient energy
- ineffective collisions occur when particles collide in the wrong orientation or without enough energy, bouncing off each other without causing a chemical reaction

Question 21

collision frequency definition

Answer 21

the number of collisions per unit of time

Question 22

what is the relationship between collision frequency and reaction rate

Answer 22

as collision frequency increases, the number of particles with energy greater than the Ea increases, so reaction rate increases

Question 23

3 factors that affect the rate of reaction and how

Answer 23

concentration
high conc = large no. of particles = high collision freq, etc

temperature
high temps = lots of movement = high collision freq, etc

pressure
high pressure = less space between molecules = high collision freq, etc

Question 24

catalyst definition

Answer 24

a substance which increases the rate of a reaction by facilitating an alternative mechanism with a lower activation energy, without being used up in the process of the reaction

Question 25

difference between homogenous and heterogenous catalysts

Answer 25

homogenous catalysts are in the same phase as reactants

heterogenous catalysts are in different phase to reactants

Question 26

4 benefits of using catalysts

Answer 26

• allows for less extreme conditions which saves money
• also saves energy, resulting in fewer CO2 emissions from burning fossil fuels
• enables different mechanisms to be used with better atom economy + reduced waste + less use of hazardous reactants
• many can be enzymes, which operate effectively close to room temperature and pressure

Question 27

example of a reaction with a homogenous catalyst

Answer 27

CH3OH(aq) + CH3COOH(aq)&raquo_space; CH3COOCH3(aq) + H2O(l)
catalyst = conc H2SO4(aq)

Question 28

example of a reaction with a heterogenous catalyst

Answer 28

haber process
H2(g) + 3H2(g)&raquo_space; 2NH3(g)
catalyst = iron (s)

Question 29

dynamic equilibrium definition

Answer 29

this occurs when a reversible reaction takes place in a closed system (or reactions in solution), when the forwards and backwards reactions are occurring at the same rate, so product and reactant concentrations remain constant
products and reactants constantly reacting together, so dynamic

Question 30

equilibrium position definition

Answer 30

refers to the relative amount of product and reactant in a reaction mixture

Question 31

what is le chatilier’s principle

Answer 31

if a change is made to a system in dynamic equilibrium, the position of equilibrium will move to counteract this change

Question 32

what 3 factors influence the position of equilibrium

Answer 32

• pressure
• concentration
• temperature

Question 33

how does concentration affect equilibrium position

Answer 33

if the concentration of reactants/products increases, the position of equilibrium will move to favour the other side

if the concentration of reactants/products decreases, the position of equilibrium will move to favour that side

Question 34

how does pressure affect equilibrium position

Answer 34

when pressure increases, position of equilibrium shifts to the side with less gaseous moles

when pressure decreases, position of equilibrium moves to favour the side with more gaseous moles

Question 35

how does temperature affect equilibrium position

Answer 35

when temperature increases, position of equilibrium moves to favour the endothermic reaction

when temperature decreases, position of equilibrium moves to favour exothermic reaction

Question 36

what is the effect of catalysts on dynamic equilibrium

Answer 36

catalytsts increase the rates of both the forwards and backwards reactions equally, so they have no effect on equilibrium position, although they can cause equilibrium to be reached faster

Question 37

describe the process by which equilibrium can be formed for a reaction in a closed system

Answer 37

• at the start, reactant concentration is high so the position of equilibrium moves to the right, favouring forward reaction
• as the reaction progresses, products concentration increases and products start reacting to reform reactants
• the rate of reaction of reactants decreases as reactant conc decreases, and the rate of reaction of the products increases as concentration increases
• eventually the rates of the forwards and backwards reaction will be equal, and the concentrations of reactants and products will be constant

Question 38

what conditions are used in the haber process and why

Answer 38

pressure = 200atm
high pressures move equilibrium position to side with less moles, increasing yield of ammonia, very high pressures are expensive/energy intensive to produce- compromise

temperature = 400C
forward reaction is exothermic, favoured by equilibrium at low temps, but this decreases reaction rate, so equilibrium wouldn’t be reached- compromise

catalyst = iron catalyst speeds up reaction rate

Question 39

what 2 other features are used to improve yields, costs and sustainability in the haber process

Answer 39

• a heat exchanger warms incoming gas mixture to maintain an optimal temp so reaction rate is high, can also remove excess heat to help keep NH3 yield high
• excess/reproduced H2 and N2 can be re-reacted to make more NH3, so high yield of ammonia

Question 40

substances in what states should be included in the kc expression

Answer 40

(g) + (l) + (aq) are included

solids (s) should be ignored

Question 41

what does it mean if kc»1

Answer 41

equilibrium lies to the RHS so reaction mixture is mostly products

Question 42

what does it mean if kc«1

Answer 42

equilibrium lies to the LHS so reaction mixture is mostly reactants

Question 43

what does it mean if kc is close to 1

Answer 43

there is a similar concentration of both reactants and products in the reaction mixture, equilibrium is somewhere in the middle