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as level chemistry > 2.1 thermochemistry > Flashcards

2.1 thermochemistry Flashcards

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1
Q

for exothermic reactions, the products have (more/less) potential energy than the reactants?

A

less

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2
Q

what are exothermic reactions?

A

reactions that give out heat

(changes potential energy into kinetic energy)

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3
Q

for endothermic reactions, the products have (more/less) potential energy than the reactants?

A

more

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4
Q

what are endothermic reactions?

A

reactions that absorb heat

(changes kinetic energy into potential energy)

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5
Q

what is enthalpy? what symbol does it have?

A

the energy at a certain pressure

H

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6
Q

enthalpy change definition

A

heat (or energy) change at constant pressure)

the change in chemical potential energy during a chemical reaction

∆H

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7
Q

if the enthalpy change is negative, the reaction is (endothermic/exothermic)?

A

exothermic

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8
Q

if the enthalpy change is positive, the reaction is (endothermic/exothermic)?

A

endothermic

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9
Q

standard enthalpy change definition

A

enthalpy changes measured under standard conditions

∆H ^θ

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10
Q

what does the enthalpy change of a reaction depend on?

A

the reaction conditions

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11
Q

what are the standard conditions for the measurement of enthalpy changes?

A
  • 1 atm (or 100kPa)
  • 25°C (or 298K)
  • constant pressure
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12
Q

what are enthalpy changes always measured in?

A

KJ mol ^-1

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13
Q

what are the 3 standard enthalpy changes you need to know?

A
  • standard enthalpy of reaction ∆rH ^ θ
  • standard enthalpy of combustion ∆cH^θ
  • standard enthalpy of formation ∆fH^θ
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14
Q

what is the standard enthalpy of reaction?

A

the general name for any reaction with an enthalpy change

( the amount of heat produced by the reaction)

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15
Q

what is the standard enthalpy of combustion? (definition)

A
  • the enthalpy change when 1 mole of substance is completely burned in oxygen under standard conditions, with all reactants and products in their standard states
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16
Q

what are the standard conditions for the standard enthalpy of combustion?

A
  • 298K
  • 100KPa or 1 atm
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17
Q

what is a standard state?

A

the state (solid, liquid or gas) of that substance at standard conditions

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18
Q

at a level what are usually the products for standard enthalpy of combustion reactions?

A

carbon dioxide or water or both

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19
Q

what is the standard enthalpy of formation? (definition)

A

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

(the enthalpy change when 1 mole of a compound is producrd from its elements in their standard states)

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20
Q

what are the 3 ways to calculate an enthalpy change for a reaction?

A
  • indirectly - using Hess’ Law
  • using bond enthalpy data
  • directly - by experiment
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21
Q

what is an example of when it is not possible to determine enthalpy change directly by experiment?

A

when many reactions are occuring at once and it becomes difficult to directly look at one specific reaction

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22
Q

what does Hess’ law state?

A

the enthalpy change for a chemical reaction is the same, independent of the route taken (from reactants to products)

(this is a consequence of the first law of thermodynamics which states that energy cannot be created or destroyed)

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23
Q

which way does the arrows go in a standard enthalpy of formation?

A

arrows up
(from the elements in their standard states that make up the reaction)

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24
Q

which way does the arrows go in a standard enthalpy of combustion?

A

(if CO2 and H2O not given,) put in a box below

arrows down

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25
Q

why is the enthalpy change of formation for elements in their standard states 0?

A

because they are in their standard states, requiring no energy to form

26
Q

bond breaking is (exothermic/endothermic)?

A

endothermic

27
Q

bond making is (exothermic/endothermic)?

A

exothermic

28
Q

overall enthalpy change equation? (bonds)

A

∆Hθ = total bond enthalpy of the reactants - total bond enthalpy of the products

(breaking - making)

29
Q

average bond enthalpy definition

A

the energy required to break 1 mole of a bond in a gaseous substance at 298K (standard conditions)

30
Q

why is the value of an enthalpy change obtained by using bond enthalpy values not always accurate?

A

because the data used for each type of bond is an average bond enthalpy; the actual value for each bond will depend on its environment

31
Q

how can you directly calculate enthalpy changes from experimental data?

A

using calorimetry (measure by experiment)

(a colorimeter is set up)

32
Q

when carrying out the process of directly calculating enthalpy changes from experimental data, what is the aim?

A

to ensure that the temperature change measured is as accurate as possible

33
Q

why is a stirrer used in a calorimeter?

A

to ensure the reaction mixture is fully mixed in order to get the maximum temperature change

34
Q

why is a lid used in a calorimeter?

A

to minimise heat loss or absorption

35
Q

why is the cup well insulated in a calorimeter?

A

so that the heat is not lost to, or absorbed from the surroundings

36
Q

what is the equation used to measure the enthalpy change of any reaction?

A

∆H = -mc∆T / n

∆H = the enthalpy change
m = total mass of the solution (g) (we assume 1cm^3=1g)
c = a constant 4.18 Jg^-1K^-1
∆T = the temperature change
n = moles

37
Q

once we know the energy given out (by q = mc ∆T), what equation can be used to convert it to an enthalpy?

A

∆H = - q/n

q = heat energy (J or KJ)
n = number or moles (always lowest no of moles)

38
Q

neutralisation reactions are (endothermic/exothermic)?

A

exothermic

39
Q

displacement reactions are (endothermic/exothermic)?

A

exothermic

40
Q

why isn’t your value when doing the calorimetry experiment exactly right?

A

due to heat loss to the surroundings

41
Q

how can you minimise heat loss to the surroundings when doing thr calorimeter experiment?

A
  • insultate the polystyrene cup
  • add a windshield
  • add a lid to the calorimetrr
42
Q

how can you obtain thr maximum value of ∆T when doing the calorimeter experiment?

A

if you have a solid you are adding to a solution, you need to make sure the solid is crushed so the reaction can happen immediately as well as ensure the heat loss is minimised

  • crush into POWDER
43
Q

what is the unit for the density of water?

A

gcm^-3

44
Q

how do you measure the temperature change of a reaction by a graph?

A

extrapolate both lines to the point where the two products were added together

45
Q

why is the temperature of the solution measured for 3 minutes before adding thr (e.g magesium)?

A

to ensure that the (initial) temperature is constant

46
Q

explain why extrapolation gives a more accurate tempetature change than using the maximum temperature recorded in the experiment?

A

extrapolation gives the temperature that would’ve been reached if the reaction occured instantly (/to allow for heat loss during the experiment)

47
Q

explain why you don’t need to accurately weigh the (e.g powderd magnesium) in a calorimeter experiment?

A

because it is in excess

48
Q

at what pressure does changes in enthalpy occur?

A

constant pressure

49
Q

what is the first law of thermodynamics? (or the law of conservation of energy)

A

energy cannot be created or destroyed only transformed from one form to another

50
Q

what is the combined equation for working out the enthalpy change?

A

∆H = -mC∆T / n

m = mass (usually liquid)
C = specific heat capacity
∆T = temperature change
n = no of moles

51
Q

what is the unit for enthalpy changes?

A

kJ mol^-1

52
Q

bond enthalpy definition

A

the amount of energy that is needed to break a bond

53
Q

standard molar enthalpy change of formation definition

A

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

54
Q

why are energy change reactions carries out in an insulated container?

A

to prevent heat energy loss to the surroundings

55
Q

give two reasons for the changing amounts of CO2 in the atmosphere after 1900? [2]

A

more burning of fossil fuels / more industries / more transportation / deforestation

(any two)

56
Q

use the temperature rise from the graph (9.6°C) to calculate the amount of heat given out during the experiment. (assume that the density of the solution is 1.00gcm^-1 and that its specific heat capacity is 4.18JK^-1g^-1)

(mass of 50 cm^3 of substance)

A

heat = 50 x 4.18 x 9.6
= 2006J

57
Q

ethanol, C2H5OH is a liquid at room temp. an equation to represent the standard molar enthalpy change of formation of ethanol is:
2C (g) + 3H2 (g) + 1/2O2 (g) -> C2H5OH (l)

suggest why this enthalpy change cannot be measured directly? [1]

A

(if these elements were reacted together) other products would form/carbon does not react with hydrogen AND oxygen under standard conditions

58
Q

in terms of bond strengths, explain why enthalpy changes of combustion are negative [1]

A

negative enthalpy change means energy in bonds broken is less than that in bonds made

59
Q

explain why the enthalpy change of combustion of propanol is more negative than that of ethanol [1]

A

more bonds broken and made in propanol and therefore more energy released

60
Q

explain why the temperature falls on continuing to add HCl AFTER the maximum temperature has been reached? [2]

A
  • no further reaction occurs
  • the excess acid cools the solution
61
Q

state what is meant by molar mass [1]

A

the mass of one mole of compound

62
Q

CO(g) + 2H2(g) —> CH3OH(l) is more exothermic than CO(g) + 2H2(g) —> CH3OH(g).
state why these two values are different, explaining your answer [1]

A

in the reaction, methanol is produced as a liquid - it is more exothermic because heat is given out when a gas becomes a liquid

as level chemistry (16 decks)

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