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IB chemistry > topic 5/15- energetics/thermochemistry > Flashcards

topic 5/15- energetics/thermochemistry Flashcards

1
Q

what is heat?

A

a form of energy, sometimes referred to as thermal energy

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2
Q

what is temperature?

A

a measure of the average kinetic energy of the particles

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3
Q

total energy is —- in chemical reactions. Why?

A

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

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4
Q

describe an endothermic reaction

A
  • heat is taken into the system from the surroundings
  • the temperature of the surroundings decreases
  • the enthalpy change of the system is positive
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5
Q

describe an exothermic reaction

A
  • heat is given out from the system to the surroundings
  • the temperature of the surroundings increases
  • the enthalpy change of the system is negative.
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6
Q

describe enthalpy

A

the heat stored by a substance

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7
Q

what three things make up a system?

A
  • the chemical reaction
  • the reactants
  • the products
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8
Q

what two things make up the surroundings?

A

the air or solvent around the reactant and product molecules

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9
Q

give the equation for heat change

A

q=mcΔT

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10
Q

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

A

equal to those of water

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11
Q

what is ΔH meaured in?

A

kJ/mol

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12
Q

ΔH of a reaction =

A

ΣΔH(products)-ΣΔH(reactants)

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13
Q

what is Hess’ Law?

A

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

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14
Q

why may Hess’ law be useful?

A

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.

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15
Q

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

A

the sum of the enthalpy changes for the individual steps

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16
Q

define standard enthalpy of formation

A

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

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17
Q

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

A

-

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18
Q

define standard enthalpy of combustion

A

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

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19
Q

what does standard state refer to?

A

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

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20
Q

describe how temperature fits into the definition of standard state

A

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

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21
Q

bond forming —— energy

A

releases

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22
Q

bond breaking ——- energy

A

requires

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23
Q

define average bond enthalpy

A

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

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24
Q

average bond enthalpies are only valid for ——

A

gases

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25
Q

ΔH =

A

bonds broken - bonds made

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26
Q

why may calculations involving bond enthalpies be inaccurate?

A

as they do not take into account intermolecular forces.

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27
Q

draw and describe an energy profile diagram for an exothermic reaction

A
  • 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
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28
Q

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

A

the products are more stable than the reactants

29
Q

draw and describe an energy profile diagram for an endothermic reaction

A
  • 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
30
Q

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

A

the products are less stable than the reactants

31
Q

define the transition state

A

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

32
Q

describe the transition state

A

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

33
Q

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

A

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

34
Q

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

A
  • 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

35
Q

define standard enthalpy change of atomisation,ΔHat

A

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)

36
Q

define first and second ionisation energies

A

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-

37
Q

define first and second electron affinities

A

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)

38
Q

what is a Born Haber cycle?

A

a Hess’ cycle for the formation of an ionic compound

39
Q

define lattice enthalpy of an ionic compound

A

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)

40
Q

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

A

charge and size of the ions

41
Q

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

A

higher

42
Q

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

A

smaller

43
Q

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

A

elements in standard states

atomisation to gaseous atoms

addition or removal of electrons to form ions

formation of lattice

44
Q

define standard enthalpy change of solution

A

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

45
Q

define enthalpy change of hydration

A

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

46
Q

give the two steps of dissolving an ionic compound

A
  • breaking the lattice into gaseous ions (lattice enthalpy) - endothermic
  • hydration of the ions, or the surrounding of the ions by water molecules - exothermic
47
Q

why hydration of ions is exothermic

A

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

48
Q

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

A

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

49
Q

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

A

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

50
Q

Define entropy (S)

A

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

51
Q

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

A

higher

52
Q

solid->liquid->gas effect on entropy

A

increases in entropy

53
Q

gas molecules have —- disorder than those in solution

A

more

54
Q

a solution is —– disordered than a solid

A

more

55
Q

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

A

greater

56
Q

why do liquids have a higher entropy than solids?

A

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

57
Q

why do gases have the highest entropy values?

A

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

58
Q

define standard entropy

A

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

59
Q

how do you calculate entropy change?

A

ΣS(products)-ΣS(reactants)

60
Q

define a spontaneous reaction

A

one which once started continues without the addition of energy

61
Q

for a spontaneous process to occur,

A

the entropy of the universe increases

ΔS(total)= ΔS(system) + ΔS(surroundings)>0

62
Q

a reaction will always be spontaneous if

a reaction will never be spontaneous if

otherwise reactions are spontaneous or not depending on

A

it is exothermic and ΔS(system) increases

it is endothermic and ΔS(system) decreases

temperature

63
Q

what does Gibbs free energy (G) do?

A

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

64
Q

give the equation for Gibbs free energy

A

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)

65
Q

a reaction occurs spontaneously when ΔG is ….

A

negative

66
Q

give an alternative equation for Gibbs free energy

A

ΔGꝋ = ΣΔGproductsꝋ – ΣΔGreactantsꝋ

67
Q

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

A
  • 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

68
Q

enthalpy change of reaction if you only have combustion values

A

enthalpy of combustion of reactants - enthalpy of combustion of products

69
Q

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

A

0

IB chemistry (29 decks)

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