topic 5/15- energetics/thermochemistry

Question 1

what is heat?

Answer 1

a form of energy, sometimes referred to as thermal energy

Question 2

what is temperature?

Answer 2

a measure of the average kinetic energy of the particles

Question 3

total energy is —- in chemical reactions. Why?

Answer 3

conserved; the first law of thermodynamics states that energy cannot be created/destroyed

Question 4

describe an endothermic reaction

Answer 4

• heat is taken into the system from the surroundings
• the temperature of the surroundings decreases
• the enthalpy change of the system is positive

Question 5

describe an exothermic reaction

Answer 5

• heat is given out from the system to the surroundings
• the temperature of the surroundings increases
• the enthalpy change of the system is negative.

Question 6

describe enthalpy

Answer 6

the heat stored by a substance

Question 7

what three things make up a system?

Answer 7

• the chemical reaction
• the reactants
• the products

Question 8

what two things make up the surroundings?

Answer 8

the air or solvent around the reactant and product molecules

Question 9

give the equation for heat change

Answer 9

q=mcΔT

Question 10

the density and specific heat capacities of aqueous solutions are assumed to be…

Answer 10

equal to those of water

Question 11

what is ΔH meaured in?

Answer 11

kJ/mol

Question 12

ΔH of a reaction =

Answer 12

ΣΔH(products)-ΣΔH(reactants)

Question 13

what is Hess’ Law?

Answer 13

the enthalpy change of a reaction is independent of the route between the initial and final states

Question 14

why may Hess’ law be useful?

Answer 14

In some cases, the heat of formation is impossible to measure directly.
Hess’s law helps us break a reaction or process into a series of small/ more easily measured steps.

Question 15

the enthalpy change for a reaction that is carried out over a series of steps is equal to

Answer 15

the sum of the enthalpy changes for the individual steps

Question 16

define standard enthalpy of formation

Answer 16

the enthalpy change when 1 mole of a substance is formed from its elements in their standard states, under standard conditions

Question 17

look at how to form standard enthalpy of formation/combustion equations

Answer 17

-

Question 18

define standard enthalpy of combustion

Answer 18

the enthalpy change when 1 mole of a substance combusts completely in oxygen, under standard conditions

Question 19

what does standard state refer to?

Answer 19

the normal, most pure stable state of a substance measured at 100kPa

Question 20

describe how temperature fits into the definition of standard state

Answer 20

temperature is not a part of the definition of standard state, but 298K is commonly given as the temperature of interest

Question 21

bond forming —— energy

Answer 21

releases

Question 22

bond breaking ——- energy

Answer 22

requires

Question 23

define average bond enthalpy

Answer 23

the energy needed to break one mol of a bond in a gaseous molecule averaged over similar compounds

Question 24

average bond enthalpies are only valid for ——

Answer 24

gases

Question 25

ΔH =

Answer 25

bonds broken - bonds made

Question 26

why may calculations involving bond enthalpies be inaccurate?

Answer 26

as they do not take into account intermolecular forces.

Question 27

draw and describe an energy profile diagram for an exothermic reaction

Answer 27

- reactants are higher in energy than the products - the reactants are therefore closer in energy to the transition state - this means that exothermic reactions have a lower activation energy compared to endothermic reactions

Question 28

in an exothermic reaction, are the reactants or products more stable?

Answer 28

the products are more stable than the reactants

Question 29

draw and describe an energy profile diagram for an endothermic reaction

Answer 29

- reactants are lower in energy than the products - the reactants are therefore further away in energy to the transition state - this means that endothermic reactions have a higher activation energy compared to exothermic reactions

Question 30

in an endothermic reaction, are the reactants or products more stable?

Answer 30

the products are less stable than the reactants

Question 31

define the transition state

Answer 31

a stage during the reaction at which chemical bonds are partially broken and formed

Question 32

describe the transition state

Answer 32

The transition state is very unstable – it cannot be isolated and is higher in energy than the reactants and products

Question 33

To compare the changes in enthalpy between reactions, all thermodynamic measurements are made under

Answer 33

standard conditions: - A pressure of 100 kPa - A concentration of 1 mol dm-3 for all solutions - Each substance involved in the reaction is in its standard state (solid, gas or liquid) - 298K

Question 34

discuss the bond strength in ozone relative to oxygen in its importance to the atmosphere

Answer 34

- UV radiation Fromm the sun can be very damaging to molecules like DNA (damage to genes->skin cancer) - both oxygen and ozone absorb UV radiation O2 (g) → O⋅ (g) + O⋅ (g) O⋅ (g) + O2 (g) → O3 (g) O3 (g) → O⋅ (g) + O2 (g) O3 (g) + O⋅ (g) → 2O2 (g) The temperature in the stratosphere is maintained by the balance of ozone formation and ozone depletion in a process known as the Chapman Cycle The double bond in oxygen is stronger than the delocalised π bonds in ozone We say the bond order of oxygen is 2 and the bond order of ozone is 1.5

Question 35

define standard enthalpy change of atomisation,ΔHat

Answer 35

the enthalpy change when one mole of gaseous atoms are made from the element in its standard states under standard conditions eg. 1/2Cl2 (g) -> Cl (g)

Question 36

define first and second ionisation energies

Answer 36

1st- the enthalpy change when one mole of gaseous +1 ion are formed from gaseous atoms under standard conditions eg. X(g) -> K+ (g) + e- 2nd- the enthalpy change when one mole of gaseous +2 ions are formed from gaseous +1 ions under standard conditions eg. X+ (g) -> X2+(g) + e-

Question 37

define first and second electron affinities

Answer 37

1st- the enthalpy change when one mole of gaseous -1 ions are formed from gaseous atoms under standard conditions eg. X(g) + e- -> X- (g) 2nd- the enthalpy change when one mole of gaseous -2 ions are formed from gaseous -1 ions under standard conditions. eg. X- (g) + e- -> X2- (g)

Question 38

what is a Born Haber cycle?

Answer 38

a Hess' cycle for the formation of an ionic compound

Question 39

define lattice enthalpy of an ionic compound

Answer 39

the enthalpy change that occurs when 1 mole of an ionic compound is separated into its constituent gaseous ions under standard conditions - it is always endothermic - eg NaCl (s) -> Na+ (g) + Cl- (g)

Question 40

the value of the lattice enthalpy of a compound depends on the

Answer 40

charge and size of the ions

Question 41

the ----- the charge the stronger the attraction between the ions so more energy is needed to separate them

Answer 41

higher

Question 42

the ----- the ion the greater its charge density so the stronger the attraction for oppositely charged ions

Answer 42

smaller

Question 43

describe the formation of an ionic lattice in steps (Born Haber cycle)

Answer 43

elements in standard states atomisation to gaseous atoms addition or removal of electrons to form ions formation of lattice

Question 44

define standard enthalpy change of solution

Answer 44

the enthalpy change when 1 mole of solute is dissolved in excess solution to form a solution of 'infinite dilution'

Question 45

define enthalpy change of hydration

Answer 45

the enthalpy change when 1 mole of gaseous ions are added to water to form an infinitely dilute solution

Question 46

give the two steps of dissolving an ionic compound

Answer 46

- breaking the lattice into gaseous ions (lattice enthalpy) - endothermic - hydration of the ions, or the surrounding of the ions by water molecules - exothermic

Question 47

why hydration of ions is exothermic

Answer 47

energy is given out as water molecules bond to the metal ions

Question 48

enthalpy of hydration of an ion becomes ----- exothermic as the charge on the ion increases

Answer 48

more; The greater the ionic charge the stronger the ion-dipole interactions that form between the ions in the ionic lattice and water molecules.

Question 49

The greater the ionic charge the stronger the ion-dipole interactions that form between the ions in the ionic lattice and water molecules.

Answer 49

less; lower charge density-> weaker ion-dipole interactions

Question 50

Define entropy (S)

Answer 50

refers to the distribution of available energy among the particles in a system

Question 51

the more ways the energy can be distributed, the ------ the entropy

Answer 51

higher

Question 52

solid->liquid->gas effect on entropy

Answer 52

increases in entropy

Question 53

gas molecules have ---- disorder than those in solution

Answer 53

more

Question 54

a solution is ----- disordered than a solid

Answer 54

more

Question 55

the side of the equation with the greater number of molecules will have ------ entropy

Answer 55

greater

Question 56

why do liquids have a higher entropy than solids?

Answer 56

the particles are free to move so there are more ways of distributing the particles and energy

Question 57

why do gases have the highest entropy values?

Answer 57

they are free to move, widely spaced, and there are therefore many more ways of distributing energy between the particles

Question 58

define standard entropy

Answer 58

the entropy of one mole of substance under standard conditions, units J/K/mol

Question 59

how do you calculate entropy change?

Answer 59

ΣS(products)-ΣS(reactants)

Question 60

define a spontaneous reaction

Answer 60

one which once started continues without the addition of energy

Question 61

for a spontaneous process to occur,

Answer 61

the entropy of the universe increases ΔS(total)= ΔS(system) + ΔS(surroundings)>0

Question 62

a reaction will always be spontaneous if a reaction will never be spontaneous if otherwise reactions are spontaneous or not depending on

Answer 62

it is exothermic and ΔS(system) increases it is endothermic and ΔS(system) decreases temperature

Question 63

what does Gibbs free energy (G) do?

Answer 63

it relates the energy that can be obtained from a chemical reaction to the change in enthalpy, change in entropy, and absolute temperature

Question 64

give the equation for Gibbs free energy

Answer 64

ΔG=ΔH-TΔS The units of ΔGꝋ are in kJ mol–1 The units of ΔHreactionꝋ are in kJ mol–1 The units of T are in K The units of ΔSsystemꝋ are in J K-1 mol–1(and must therefore be converted to kJ K–1 mol–1 by dividing by 1000)

Question 65

a reaction occurs spontaneously when ΔG is ....

Answer 65

negative

Question 66

give an alternative equation for Gibbs free energy

Answer 66

ΔGꝋ = ΣΔGproductsꝋ – ΣΔGreactantsꝋ

Question 67

describe the relationship between Gibbs free energy and equilibrium using a graph

Answer 67

- from the time a reversible reaction commences to the point it reaches equilibrium, the Gibbs free energy is changing as the ratio of reactants to products alters As the amount of product increases and the reaction moves towards completion (non reversible) or equilibrium (reversible): - Gibbs free energy decreases - forward reaction is favoured - minimum GFE = equilibrium Past the equilibrium: - forward reaction becomes non-spontaneous - reverse reaction becomes spontaneous

Question 68

enthalpy change of reaction if you only have combustion values

Answer 68

enthalpy of combustion of reactants - enthalpy of combustion of products

Question 69

the standard enthalpy change of any element in its standard state is

Answer 69

0