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Edexcel Chemistry A LEVEL - RM > TOPIC 8 - CHEMICAL ENERGETICS I +II > Flashcards

TOPIC 8 - CHEMICAL ENERGETICS I +II Flashcards

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1
Q

two components make up chemical energy?

A

Kinetic energy - The motion of particles in a substance

Potential energy - How strongly particles in a substance interact with one another (attract and repel)

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2
Q

What is heat energy?

A

The portion of kinetic and potential energy of a substance that is responsible for the temperature of the substance

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3
Q

enthalpy change

A

The heat energy that is transferred between system and surroundings. kJmol-1
_^ H = H[products] - H[reactants]

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4
Q

exothermic

A

When heat energy is transferred from the system to the surroundings and when bonds are formed.
REVERSIBLE.
- freezing h20, neutralisation, combustion, respiration.
- Reaction profile = Enthalpy changes during a reaction
- products are at a lower energy level as _^ is negative. an increase in temp and more energy released in making more bonds.
Incomplete combustion is less exothermic

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5
Q

endothermic

A

When heat energy is transferred from the surroundings to the system and bonds are broken.
IRREVERSIBLE
- melting, photosyn, carbon fix, thermal decomp of limestone, gel pack- energy exchange heat, sherbet - C6H8O7 citric acid and NaHCO3 –> Na3C6H8O7 + H20+ CO2. Hydrated blue cusulphate into white anhydrous cusulphate.

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6
Q

activation energy

A

The minimum energy needed for the reaction to occur

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7
Q

What are and why are standard conditions important?

A

Elements in their standard states
100kPa
298K

Changes in temperature and pressure can affect enthalpy

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8
Q

standard enthalpy change of formation ΔH°f

A

The enthalpy change when one mole of a product is formed from its constituent elements in their standard states under standard conditions

products - reactants

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9
Q

standard enthalpy change of combustion ΔH°c

A

The enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions

reactants - products

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10
Q

standard enthalpy change of neutralisation ΔH°n

A

The enthalpy change when an acid and an alkali react to form one mole of water under standard conditions

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11
Q

Standard enthalpy change of Reaction, ΔH°r

A

The standard enthalpy change of a reaction is the enthalpy change which occurs when molar quantities of materials react under standard conditions, and with everything in its standard state (100kPa, 298K)

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12
Q

How to measure the enthalpy change of neutralisation

A

•React acid and alkali in a polystyrene cup in a beaker
•Make sure the temperature of the acid and alkali are the same.
what is happening is H+ + OH- –> H20
•Measure the change in temperature.
ERROR:
- uncertainties of measurements using pipette or thermometre.
- some heat transf. to thermom or polystyerene cup.

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13
Q

Experiment to measure enthalpy change of combustion

A

1.Weigh the liquid in the burner (alcohol)
2.Light a spirit burner under a beaker containing water
3.Measure the initial temp
4.When the temperature of the water has increased by 20 C extinguish the flame
5.Measure the final temperature
6.Enthalpy change can not be measured directly
(Q=mcΔT)

ERRORS:

  • energy losses from calorimeter
  • incomplete combust of fuel
  • incomplete transfer of energy
  • evaporation of fuel after weighing
  • heat capacity of calorimeter should not be included
  • measurements not carried out under standard conditions as H20 should be gas not liq

In homologous series the no. of C atoms increases so enthalpy of combust increases by constant amount. Calculated value is greater than experimental due to error.

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14
Q

Experiment to measure enthalpy change of combustion

A

1.Weigh the liquid in the burner (alcohol)
2.Light a spirit burner under a beaker containing water
3.Measure the initial temp
4.When the temperature of the water has increased by 20 C extinguish the flame
5.Measure the final temperature
6.Enthalpy change can not be measured directly
(Q=mcΔT)

ERRORS:

  • energy losses from calorimeter
  • incomplete combust of fuel
  • incomplete transfer of energy
  • evaporation of fuel after weighing
  • heat capacity of calorimeter should not be included
  • measurements not carried out under standard conditions as H20 should be gas not liq

In homologous series the no. of C atoms increases so enthalpy of combust increases by constant amount. Calculated value is greater than experimental due to error.

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15
Q

equation for energy change and enthalpy change

A

Q (J) =m(g)cΔT(K)
energy change = mass of subs heatedspecific heat capatemp
ΔH (kJ/mol) = Q/n

ALWAYS USE LIMITING REACTANT TO FIND RATIO

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16
Q

experiment for measure enthalpy change of reaction

A

. put polystyrene cup in a beaker for insulation and support
. Measure out desired volumes of solutions with volumetric pipettes and transfer to insulated cup
. measure the initial temperatures of the solution or both solutions if 2 are used. Do this every minute for 2-3 minutes.
. At minute 3 transfer second reagent to cup. If a solid reagent is used then add the solution to the cup first and then add the solid weighed out on a balance.
If using a solid reagent then use ‘before and after’ weighing method.
. stirs mixture (ensures that all of the solution is at the same temperature)
. Record temperature every minute after addition for several minutes.

ERROR
• energy transfer from surroundings (usually loss)
• approximation in specific heat capacity of solution. The method assumes all
solutions have the heat capacity of water.
• neglecting the specific heat capacity of the calorimeter- we ignore any
energy absorbed by the apparatus.
• reaction or dissolving may be incomplete or slow.
• density of solution is taken to be the same as water.

Extrapolate the temperature curve/line back to the time the reactants were added together.
As when adding 2nd subst there is heat loss/ gain straight away so temp does not rise/fall too much. We consider heat loss/gain to be constant in time so we extrapolate to get true value of change in temp.

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17
Q

Hess’s law

A

Enthalpy change of a reaction which is independent of the route by which chemical change occurs.

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18
Q

impossible to find enthalpy change of a reaction

A
  • C and H do not react under normal conditions
  • C and 02 do not react to form CO2 solely but others CO.
    . Hess’s law is used to work out the enthalpy change to form a hydrated salt from an anhydrous salt.
    This cannot be done experimentally because it is impossible to add the exact amount of water and it is not easy to measure the temperature change of
    a solid. Instead both salts are dissolved in excess
    water to form a solution of copper sulfate.
    The temperature changes can be measured
  • ADD aq to the cycle inside
    . Hess’s law is used to work out the enthalpy change for the thermal decomposition of calcium carbonate.
    This cannot be done experimentally because it is impossible to add the heat required to decompose the solid and to measure the temperature change of a
    solid at the same time. Instead both calcium carbonate and calcium oxide are reacted with hydrochloric acid to form a solution of calcium chloride. The temperature changes can be measured.
  • ADD HCl aq to reactant end and product end to give 2 DIFF REACTIONS forming products down and arrow points down too.
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19
Q

Hess’s Law calc

A
  • always clockwise
  • balance eq
    FORMATION
    ΔH reaction = Σ ΔfH products - Σ ΔfH reac
    all elements = 0
  • arrow up from elements to reaction formed
    COMBUSTION
    ΔH reaction = Σ ΔcH reactants - Σ ΔcH products
  • arrow down to combust products
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20
Q

Bond enthalpy

A

Enthalpy change when one mole of a bond in the gaseous state is broken.

NO LIQ OR SOLID - ALL GASES EQ
ΔH = Δ bond enthalpies broken - Δ bond enthalpies made
all reactants bonds broken - all product bonds made

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21
Q

Mean bond enthalpy

A

Enthalpy change when one mole of a bond, averaged over many different molecules, is broken.

Bond enthalpy value is diff from value of data book coz data book averages out over a range of different molecules. These _ bonds will be in a diff environment from compound which will affect bond strength

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22
Q

enthalpy of lattice formation (ionic)
/\ lef H
EXO

A

is the energy change when 1 mole of ionic solid is formed from it’s gaseous ions under standard conditions

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23
Q

enthalpy of dissociation

/\led H ENDO

A

when 1 mole of a solid ionic compound is broken up into its constituent ions in gas phase

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24
Q

enthalpy of atomisation
/\atH
ENDO always

A

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

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25
Q

first ionisation enthalpy
/\IE H^o
ENDO

A

is the enthalpy change required to remove 1 mole of e-s from 1 mole of gaseous atoms to form 1 mole of gaseous ions of 1+ charge

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26
Q

second ionisation enthalpy
/\IE H
ENDO

A

is the enthalpy change required to remove 1 mole of e-s from 1 mole of gaseous +1 ions to form 1 mole of gaseous ions of 2+ charge

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27
Q

first electron affinity
/\eaH
EXO

A

the enthalpy change that occurs when 1 mole of gaseous atoms gain 1 mole of e- to form1 mole of gaseous ions with -1 charge

  • ion is more stable that the atom and there is an attraction btw nucleus and the e-
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28
Q

second electron affinity
/\eaH
ENDO

A

the enthalpy change when 1 mole of gaseous 1- ions gains 1 e- per ion to produce 1 mole of gaseous 2- ions

  • endo because it takes energy to overcome the repulsive forces btw the -ve ion and the e-
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29
Q

first electron affinity values become less exo when going down

A

gp 7 as atoms get bigger, more sheilding so less easily attract e-s to form the -ve ion

30
Q

enthalpy of hydration
hyd

EXO

A

enthalpy change when one mole of gaseous ions become hydrated in water to infinite dilution - water molec totally surround ions

31
Q

enthalpy of solution

varies

A

enthalpy change when one mole of solute dissolves completely in a solvent to infinite dilution

32
Q

strength of lattice enthalpy depends on

A
  1. size of ion - larger the ion, less neg the enthalpy - further apart the ions are so have weaker forces of attractions
  2. charges on the ion - bigger the charge, greater the forces of attraction so more -ve values
33
Q

eq for /\LE H in born harber cycle

A

= /\fH - (sum of all /)

34
Q

eq for /\LE H in born harber cycle

A

= /\fH - (sum of all /)

in order to measure the whole ionic bond strength

35
Q

theoretical lattice enthalpies assumes

A

a perfect ionic model - ions are 100% and spherical and attractions are purely electrostatic
- charge evenly distributed around centre so each ion can be considered as point charges

36
Q

the more covalent character

A

the bigger the difference btw born Haber lattice enthalpies and theoretical values
theoretical < born harber

37
Q

polarising power of cation increases

A
  • +ve ion is small

- +ve ion has multiple charges

38
Q

polarizability of an anion depends on

A

size

the bigger the anion the more easily its distorted

39
Q

how to draw a born harber cycle

A
  1. /\fH
  2. /\ attH of each element metal
  3. /\IE H of metal and successive IE if applicable
  4. /\att H of non-metals
  5. /\Ea H of non-metals and successive if applicable
  6. /\LE H
    elements in standard states to gaseous atoms to gaseous ions to ionic lattice
40
Q

if there is a big incr in ionsiation enthalpy to remove 3rd e- Ca3+ this

A

does not compensate for by the stronger lattice enthalpy of formation
the enthalpy of formation is ENDO - CaCl3 is the least stable form

41
Q

we cannot measure lattice enthalpy directly

A
  • use experimental values via born harber cycle

- theroretical values via physical calc

42
Q

NaF / NaCl which lattice enthalpy is more exo

A
  • cations same size and charge

- anions same charge but F- is smaller than Cl- so NaF is more exo

43
Q

NaCl / MgCl2

A
  • anions same size and charge
    -Mg2+ smaller than Na+
    and have higher charge density
44
Q

MgO / MgCl2

A
  • cations same size and charge

- O2- smaller than Cl- and has higher charge density so MgO

45
Q

eq of /\sol H

A

/\sol H = -/\LE H + sum of /\hyd H

46
Q

hydration enthalpies are EXO as energy is given out as water molec bond to metal ions

A

higher the hydration enthalpy - the smaller the ions and greater the charge

47
Q

/\sol H tells us

A

if exothermic - substance is soluble - gives energy

48
Q

if subst is insoluble

A

lattice enthalpy > than hydration enthalpy and is not energetically favourable to break up lattice so /\solH is endothermic

49
Q

Entropy (S)

A

is a measure of disorder in a system

50
Q

Entropy (S)

A

is a measure of disorder in a system

ALWAYS POSITIVE

51
Q

more spreading out of energy

A

incr in entropy

52
Q

more random arrangement of particles

A

incr in entropy

2Cus + 02g -> 2CuOs = decr in entropy

53
Q

signs in chemistry

A

gives info abt a reaction

54
Q

how entropy is affected by temp

A

entropy of pure subst increases with incr energy
1. particles at higher temp have higher energy and move more
2. the arrangement of particles at higher temp becomes more random
solid

55
Q

when water boils and becomes gaseous its entropy rises

A

as temp is constant at 100’ there is a change in state

56
Q

entropy when dissolving ionic solids

A

if a reaction results in products tht allow more disorder then entropy incr
e.g. hydrated CuSO4.5H2O = Cu2+aq + SO42- aq + 5H2Ol

57
Q

how does the no of gas molec affect entropy

A

rise in no. gas molec during a reaction causes an incr in entropy
e.g. decr in entropy 2Mgs + O2 g -> 2MgO s

58
Q

entropy of salt

A

0 - ionic as at 0K = 0 entropy there is no disorder as particles are stationary

59
Q

/\syst* =

A

S* products - S*reactantts

60
Q

/\syst* =

A

S* products - S*reactants

times by moles present

61
Q

spontaneous processes/ changes

A

for a process to happen spontaneously entropy must incr. if reaction is spont. and feasible it will take place of its own accord and does not take into account of ROR
e.g. metal rusting in O2
2Fes + O2g -> 2FeOs
entropy is low but still spontaneous react occurs.

  • ENTHALPY will be NEG
62
Q

spontaneous processes/ changes

A

for a process to happen spontaneously entropy must incr. if reaction is spont. and feasible it will take place of its own accord and does not take into account of ROR
e.g. metal rusting in O2
2Fes + O2g -> 2FeOs
entropy is low but still spontaneous react occurs.
- /\S tot must be POSIT
- ENTHALPY will be NEG

63
Q

Entropy change in any reaction eq

A

/\S tot = /\S syst + /\S surround

64
Q

entropy change in surroundings

A

/\S surr = - /\H (J/mol)/ T (K)

65
Q

Gibbs free energy

A

using entropy and enthalpy used to determine the feasibility of the reaction and temp

/\G (KJ/mol) = /\H - T/\S syst need to convert J/K/mol into (KJ/mol)

66
Q

entropy of a reaction is always +ve

A

change still may only be spontaneous at certain temps. as temp affects the -/+ value

67
Q

for a reaction to be spontaneous

A

/\S tot = + ve

/\G must be -ve < 0

68
Q

At minimal temp or above the reaction is feasible

A
  • /\G = 0
    0 = /\H - T/\S
    T = /\H / /\S
69
Q

/\G links to gas constant

A

= - TRlnk
T - temp
R - gas constant
k - equilib constant

rearrange
k = e- ( /\G (J) / TR)

if /\G is -be then K>1 more equilib lies towards products

70
Q

if /\G is +ve than exponent is +ve and K<1 so

A

equilib lies towards reactants and never get 100% products. /\G is positive it does not mean reaction is not taking place but it’s impossible for all reactants to converted into products

71
Q

rough rule of thumb

A

if /\G < -40kJ/mol all reactants are converted to products

if /\G > +40 very few reactants converted to products

/\G = - RTlnk = /\H - T/\S

lnk = - /\H/RT + /\S/R
- incr in temp = -/\H/RT - less +ve so overall total lnk = smaller - k = smaller this shifts equilib to counteract the change and shift to the left

Edexcel Chemistry A LEVEL - RM (12 decks)

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