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Inorganic Chemistry Year 12 > 3.2.1 Enthalpy > Flashcards

3.2.1 Enthalpy Flashcards

1
Q

what is the symbol of enthalpy

A

H

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2
Q

what is enthalpy

A

the measure of the heat energy within a chemical system

  • thought of as the energy stored within bonds
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3
Q

what can you measure with enthalpy

A
  • can measure enthalpy change, NOT just enthalpy
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4
Q

what is a chemical system

A

atoms, ions and molecules making up the chemicals

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5
Q

why is their enthalpy change in reactions, and how is this measured

A

the reactants and products in a chemical reaction are likely to have different enthalpies
- H(reactants) - H(products)
- can be negative or positive, depending on whether the products contain more or less energy that the reactants

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6
Q

what is the conservation of energy, and how does this correlate to enthalpy changes

A

energy cannot be created or destroyed, only transferred
- so, when an enthalpy change occurs, it means that heat energy has been transferred to or from the surroundings

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7
Q

what is system, surroundings and universe

A

system= chemicals (reactants and products)
surroundings= apparatus (thermometer and lab)
universe= both

  • measure enthalpy change by measuring energy transfer between system and the surroundings
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8
Q

what is exothermic

A

where energy is transferred from the system to the surroundings
- rise in temperature

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9
Q

what is endothermic

A

where energy is transferred from the surroundings to the system
- drop in temperature

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10
Q

what do enthalpy profile diagrams show

A
  • the relative enthalpies of reactants and products
  • the enthalpy change
  • shown via progress of reaction on x axis, and enthalpy H on y axis
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11
Q

why is exothermic enthalpy change always negative

A
  • chemical system releases heat energy to surroundings, and energy LOST by system = energy gained by surroundings
  • temp of surroundings increase, as they gain energy
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12
Q

how would you draw an enthalpy profile diagram for exothermic reactions

A
  • reactants above products
  • arrow points downwards
  • negative enthalpy change
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13
Q

why is enthalpy change of endothermic reactions always positive

A

energy is transferred from surroundings to the system
- so positive enthalpy change, as system is taking in energy
- temperature of surroundings is decreasing, as losing heat energy

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14
Q

how would you draw an enthalpy profile diagram for endothermic reactions

A
  • products have more energy that reactants
  • arrow points upwards, with positive enthalpy
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15
Q

what is activation energy

A

the minimum amount of energy required for a reaction to take place
- Ea

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16
Q

what does a big or small activation energy means

A

SMALL= reaction takes place rapidly, as the little amount of energy needed to break bonds is readily available from surroundings
LARGE= large energy barrier, so slow reaction, or doesn’t occur at all

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17
Q

how would you label activation energy on an exothermic enthalpy profile diagram

A

small rise from reactants up, and then big drop
- extend line of reactants to label Ea, pointing up

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18
Q

how would you label activation energy on an endothermic enthalpy profile diagram

A

big jump up from reactants, then small line down to products

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19
Q

where is the activation energy arrow always on diagrams

A

always from reactants line to the peak of the curve

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20
Q

why can ΔH values vary for the same reaction

A

dependent on the conditions used

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21
Q

which values do chemists use when recording ΔH

A
  • standard conditions, that are close to typical working conditions
  • shown in data symbol using little theater symbol on top : ΔH°
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22
Q

what are the standard condition values

A

pressure = 100kPa (similar to one atm)
temperature= 298K (25C)
concentration= 1moldm-3 (if relevant)

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23
Q

what is meant by standard state

A

physical state of a substance under standard conditions

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24
Q

what is the standard enthalpy change on a reaction

A

the enthalpy change that accompanies a reaction in the molar quantities shown in a chemical reaction under standard conditions (all reactants and products in their standard states)

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25
what is important to remember when looking at the standard enthalpy change of reaction values
- ΔH°r - refers to a specific equation, and value of ΔH° depends on the balancing equation - e.g. if you were to halve the moles of all given values, the ΔH° would also halve
26
what is the enthalpy change of formation
- ΔfH⦵ - the enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions, with all reactants and products in their standard states - ALWAYS ONE MOLE OF PRODUCT!
27
what is the ΔfH⦵ of elements
always equal to 0 - as forming an element from element, so no change
28
what is the enthalpy change of combustion
- Δ𝐻⦵c the enthalpy change that takes place when ONE MOLE of substance/fuel reacts COMPLETELY with oxygen under standard conditions, with all reactants and products in their standard states
29
what will products always be in combustion reactions
oxides of elements in the substance
30
what is important to remember for the Δ𝐻⦵c
- always a negative enthalpy change, as an exothermic reaction - fractions may be used in equation, to ensure that the fuel stays at one mole
31
what is the enthalpy change of neutralisation
- ΔneutH⦵ the enthalpy change that accompanies a reaction of an acid by base to form ONE MOLE of H2O, under standard conditions, with all reactants and products in their standard states
32
why is the ΔneutH⦵ same for all reactions
- -57kJmol- - all reactions involve reaction of H+ and OH- to form one mole of water
33
what are you measuring the temperature change of when measuring ΔH
- the temperature change of surroundings
34
what are the values required when measuring ΔH
- mass - specific heat capacity - temperature change
35
what is the equation for ΔH
ΔH= mcΔT - if per mole, divide by n - if J, not kJ, divide by 1000 too
36
how do you find mass value for ΔH equation
- measure by weighing, in g - measure the mass of material changing temperature, usually the solution
37
how do you find the SHC value of ΔH equation
- energy required to raise the temperature on 1kg of substance by 1C/K - usually use the value of water = 4.18
38
how do you measure the temperature change in the ΔH equation
determined by thermometer reading - final temp-initial temp
39
what unit does the basic ΔH equation give you
just J (joules)
40
how can you determine the enthalpy change of combustion of methanol
- use measuring cylinder to measure out 150cm3 of water and add to beaker - record the initial temperature of the water, to the nearest 0.5C, using thermometer - add methanol to spirit burner, making it easy to burn the fuel, and record initial mass - place spirit burner under beaker, and light the burner to burn the methanol - stir the water using the thermometer/glass rod at same time - after 3 minutes, extinguish the flame - immediately record the max temperature reached by the water - reweigh the spirit burner containing the methanol, whilst assuming that the wick has not been burned
41
how would you calculate the ΔH for the combustion practical
- m= of the water, as density means that volume is equal to mass - c and ΔT are of water too - moles at bottom is of methanol, worked out by dividing mass taken by molar mass
42
why is the value you calculate for ΔHc usually less negative than what is should be
- you get a less exothermic value 1) heat is lost to surroundings other than water (beaker and air surrounding flame) 2) methanol has been incompletely combusted (CO/C is produced) 3) evaporation of methanol from the wick - need to weigh burner as soon as flame is extinguished, or methanol may disappear (usually spirit burners have cover to stop this) 4) non-standard conditions
43
how can you minimise inaccurate ΔHc results
- input more O2, draught screens, constantly stir
44
how can you determine the ΔHr using a practical
- e.g. a solid and a solution - use a polystyrene cup (cheap, waterproof, lightweight, offers some insulation) - solution acts as the immediate surroundings (chemical particles IN the solution may react when they collide, and the ΔH is between the chemical particles and the water molecules
45
what are things to remember when calculating ΔHr
- you use the moles of whichever reactant isn't in excess (usually the solution, as powder is added in excess) - use the SHC of water - use the density of water
46
explain the practical for adding powder and solution together for determining ΔHr
1) add solution into polystyrene cup, and take record of mass 2) weigh out excess zinc powder 3) start stopwatch, and take temperature of solution every 30 seconds, until a constant temperature is reached 4) add zinc into solution, and take temperature every 30 seconds until temperature has fallen for several minutes 5) plot graph of temperature against time 6) to correct for cooling: EXTRAPOLATE DATA: - extend both the lines of best fit until you reach the time that the zinc was added - draw a vertical line from the time at the 2 lines going down - difference in temp is the ΔT you will use in the equation
47
how do you calculate the ΔHneut
- very similar to ΔHr, BUT: - to find the mass, you add two values together - for moles, can use either quantities (usually using concentration equation), as usually equal, but even not, will need to use equation to equal to water - remember, in final equation, need to use value for ONE MOLE of H2O burned - the moles of H2O FORMED in the value of n
48
what is average bond enthalpy
the energy required to break one mole of a specified type of bond in a gaseous molecule
49
why are average bond enthalpies always endothermic
- energy is always required to break a bond - so always has to have a positive enthalpy value
50
why are bond enthalpies an estimated average
1) actual bond enthalpies can vary, depending on the chemical environment of the bond. the average is calculated from the actual bond enthalpies, and as they are all very similar, you can use the average for all. 2) all species need to be gaseous molecules, which may mean that the answer is not under standard conditions (e.g. if producing H2O (g), instead of actual liquid (therefore would have to add on the ΔH of condensation of H2O too)
51
what process is bond breaking
ENDOthermic: - required energy, so positive
52
what process is bond making
EXOthermic: - releasing energy, so negative
53
when can a reaction be classed as endothermic
more energy is needed to break bonds than is released when making bonds
54
when can a reaction be classed as exothermic
more energy is released when making bonds than is needed to break bonds
55
where is bond breaking and bond making seen on an enthalpy change profile
- bond breaking = between reactants and activation energy - bond making = between top of activation energy peak to products
56
how would you calculate ΔH from average bond enthalpies
- sum of (bond enthalpies of reactants) - sum of (bond enthalpies of products)
57
what is an overview of Hess' law
if a reaction can take place by 2 routes, and the starting and finishing conditions are the same, the total ΔH is the same for each route (A+B=C) - many values of ΔH are in data book, and can be used to find unknown values
58
how can you use the values of ΔHf in Hess law
- you know values to get from elements to reactants and products - so can make triangle out of this and work out
59
how can you use the values of ΔHc in Hess law
- can combust both reactants and products - may see that they have similar products - therefore construct triangle
60
key things to remember in Hess Law
- elements have ΔH value of 0, so can be ignored - always remember directions of arrows - reactants or products that are common on both sides can also be cancelled out (or can add certain substances to either side to fit the values known)

Inorganic Chemistry Year 12 (7 decks)

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