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thermodynamics Flashcards

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1
Q

what is bond dissociation enthalpy

A

enthalpy change when one mole of a bond/ covalent is broken to give separated atoms with everything in the gase state

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2
Q

what is the standard enthalpy of formation

A

the standard enthalpy of formation is the enthalpy change when one mole of a substance is formed from its constituent elements in its standard states under standard conditions

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3
Q

what are ionisation energies

A

the first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms of the elements

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4
Q

what is the enthalpy of lattice formation

A

the enthalpy change when 1 mole of ionic solid is formed its gaseous ions

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5
Q

what is the enthalpy of dissociation

A

the enthalpy change when 1 mole of ionic solid is converted to its gaseous ions

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6
Q

what happens when we know the lattice enthalpy

A

when we know the enthalpy of lattice, we can predict the strength of the ionic bonding

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7
Q

how can we predict which ionic compound is more exothermic

A

It is whichever has the stronger ionic bonding is more exothermic

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8
Q

what are the factors that effect the strength of ionic bonding

A
  • the size of the ions
  • the change of ionic bonding

the strength of ionic bonding increases as:

  • the ions get smaller
  • the ions get more highly charged
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9
Q

how do we find lattice enthalpy

A

we can’t find it directly as it is too difficult but we can do it indirectly using a Hess cycle - which can be neatly drawn as a Born Haber cycle

we can’t find it directly because:

  • there isn’t any equipment capable of forming and reacting gaseous ions
  • heating an ionic compound results in gaseous ion pairs, not individual ions
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10
Q

what are the stages of the Born Haber cycle

A

1) enthalpy of formation
2) atomisation
3) ionisation
4) electron affinities

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11
Q

what is atomisation

A

when 1 mole of gaseous atoms are formed from an element in its standard state

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12
Q

how can we work out the enthalpy of atomisation

A

the enthalpy of atomisation = 1/2 of bond dissociation enthalpy

this applies when:
we are breaking 1/2 mole of a molecule e.g. Cl2

forming separate atoms
everything is in gaseous states in its standard state

only when the elements are covalently bonded in their standard state

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13
Q

what is electron affinity

A

the enthalpy change when 1 mole of electrons are added to 1 mole of gaseous atoms

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14
Q

endothermic or exothermic

A

1) lattice enthalpy = exothermic
2) enthalpy of formation = exothermic
3) ionisation = endothermic
4) electron affinities = EA1 exothermic EA2=endothermic

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15
Q

why is the first electron affinity exothermic

A

the first electron affinity is exothermic because as the electron gets closer to the positive nucleus energy is being given out

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16
Q

why is the second, third e.t.c electron affinity endothermic

A

the electron is being added to a negatively charged ion

therefore, the repulsion outweighs the attraction of the ions nucleus
so the overall force between the anion and the electron is repulsive

so energy is required to add that electron in

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17
Q

what does Hess Law state

A

the enthalpy of a reaction is independent of the route the reaction takes

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18
Q

why are the theoretical and experimental values for lattice enthalpy different

A

theoretical and experiment values enthalpies can be different depending on how “purely ionic” the compound is

the experimental value is derived from Born Haber cycles and take into account the covalent characteristics of the salt/ionic compound

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19
Q

what are the theoretical lattice enthalpies

A

theoretical lattice enthalpies are calculated from data assuming a PERFECTLY IONIC MODEL OF A LATICE

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20
Q

what is the perfectly ionic model

A
  1. ions that are perfectly spherical - there is no distortion
  2. The charge is evenly distributed in this sphere - this is called point charge
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21
Q

what are the practical lattice enthalpies

A

the experimental value is derived from Born Haber cycles and take into account the covalent characteristics of the salt/ionic compound

this number is often different from the theoretical value

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22
Q

what does the difference between the theoretical and practical enthalpy imply

A

This tells us that the compound being experimented on doesn’t follow the perfectly ionic model and has some covalent characteristic

e.g.
most of the time the positive ion distorts the charge distribution in the negative ion. We say the positive ion POLARISES the negative ion

23
Q

the more polarisation(…)

A

the more polarisation of the ionic bonds, the more electron density is shared between ions, the covalent character increases and so the difference between the theoretical and experimental values increases

24
Q

how can we figure out whether a substance is purely ionic

A

comparing theoretical and practical lattice enthalpy we can see how much a substance is purely ionic

in all of the examples, the experimental value is always more exothermic than the theoretical value

25
Q

what does a large difference in theoretical and experimental values show

A

it shows that we have some covalent characteristics being displayed

This is caused by larger distortions in the negative ion

The bigger the difference between the lattice enthalpy, the more polarisation you have
The greater the covalent character

26
Q

what happens when the cation is next to the anion

A

when a cation is next to an anion the electron cloud of the anion is distorted towards the anion

27
Q

As the amount of polarisation increases (…)

A
  • more electron density is shared between ions
  • the amount of covalent character increases
  • the difference between experimental and theoretical enthalpy increases
28
Q

what is the polarising power determined by

A

the charge of the positive ion

the size of the positive ion

Larger amount of polarisation - higher polarising power - higher polarising power of the stated cation and/or higher polarisability of the anion

29
Q

what is common exam question in regards to silver halide

A

although silver cation is large with a 1+ charge, ALL SILVER HALIDES have a large difference between their theoretical and experimental lattice energies

SO WHAT WE NEED TO SAY IS:
AgBr has a larger difference in theoretical and experimental lattice energies
AgBr has a large amount of covalent character

30
Q

why are covalent characteristics good

A

covalent character can increase the strength of bonding in an ionic compound

since experimental lattice enthalpies account for covalent character while theoretical lattice enthalpies don’t, experimental lattice enthalpies are always more exothermic or negative than the theoretical ones

31
Q

what are factors that affect polarisation

A

1) polarising power

2) polarisablity of the anion

32
Q

if an anion has a high polarisability (…)

A

some anions electron clouds are more easier to distort than others this is because(…)

the larger and more highly charged the ion is, the easier it is to distort/ higher polarisability

sometimes we are unable to predict the polarisability

33
Q

what is the enthalpy of the solution

A

the change when 1 mole ionic substance is dissolved in the minimum amount of solvent to ensure no further enthalpy change is observed upon further dilution

-dissolves in enough solvent to form an infinitely dilute solution
e.g.
take a solid, add water until we have just dissolved the ionic compound

34
Q

what must happen before something dissolves

A
  1. substances bonds must break (endothermic)
  2. new bonds formed between the solvent and substance (exothermic)
    e. g. adding an ionic salt to water

ionic lattice in solid form

substance bonds are broken to create free moving ions

bonds are formed between ions and water. The ions are therefore hydrated

35
Q

what do most ionic compounds dissolve in

A

polar solvents like H2O

The delta positive hydrogen is attracted to the negative ions and the delta negative oxygen is attracted to positive ions

The structure, therefore, begins to break down

36
Q

what must happen before the ionic structure is broken down and hydration occurs

A

for this to happen, new bonds formed must be the same strength or greater than those broken

If not, it is unlikely to dissolve

therefore, enthalpies for soluble substances are exothermic

essentially, if the energy for the first process (endothermic - enthalpy of solution) is greater than the energy in the second process (exothermic- enthalpy of hydration ) the substance will not dissolve

37
Q

what is entropy

A

entropy is the measure of disorder in a system

entropy (S) is the number of ways energy can be shared out between particles
It also tells us the ways in which particles can be arranged

by disorder we mean randomness

38
Q

the more disorder there is ….

A

the more disorder there is the higher the level of entropy

e.g. disorder increases from solid, liquid to a gas

39
Q

why do solids have the lowest level of disorder

A

solids have the lowest level of disorder, particles are arranges neatly in rows

40
Q

why do liquids and gases have higher entropy than solids

A

liquids and gases are more disordered

to put liquids into a state of order you need to apply energy. Particles want to be the lowest energy possible therefore entropy requires little energy

41
Q

what affects entropy

A

the number of particles also affect entropy change

If a reaction is in the same state but more moles are produced, the entropy increases

This is because there are more ways energy can be distributed

42
Q

how can a reaction occur even when it is enthalpically unfavourable

A

a reaction can be spontaneous (feasible) even if it is enthalpically unfavourable (e.g. endothermic) because a reaction will tend towards more disorder and hence increase entropy

-changes in entropy can sometimes change in enthalpy

e.g.
a reaction is very enthalpically unfavourable as it is very endothermic

the number of mole increases and so the reaction is entropically favourable

starting with 2 solids but making gas and a liquid

43
Q

how can we work out entropy

A

the entropy of products - entropy of reactants

most reactions will not occur unless the entropy change is positive
entropy is represented by symbol S

the units are in joules per kelvin per mol

44
Q

what is Gibbs free energy

A

Gibbs free energy tells us if a reaction is feasibly or not

45
Q

how do we work out Gibbs free energy

A

enthalpy - (entropy change x temperature)

46
Q

when is a reaction feasible

A

a reaction is feasible in theory if Gibbs is negative or

zero

47
Q

even if a reaction is feasible, you may not observe a reaction occurring why?

A

this is because the activation energy is too high or the rate of the reaction is very slow

48
Q

when is a reaction feasible at any temperature

A

if a reaction is exothermic (negative enthalpy - delta H) and has a positive entropy change, the delta G (Gibbs free energy) is always negative since:

G = H – TS

these reactions are feasible at any temperature

49
Q

when is a reaction never feasible at any temperature

A

if a reaction is endothermic (positive enthalpy - delta H) and has a negative entropy change, the delta G (Gibbs free energy) is always positive since:

G = H – TS

these reactions are never feasible at any temperature

50
Q

when is a reaction feasible only at certain temperatures

A

if delta H is positive (endothermic) and S is also positive then the reaction will only be feasible at a certain temperature. This is because:

G = H – TS

51
Q

how can we find if the temperature when gibbs equals 0

A

by rearranging this equation:

G = H – TS
to this

T= H/S

This is because when G = 0
H=TS

52
Q

What is on the y axis and what is on the x-axis in a Gibbs free energy graph

A

temperature (K) is on the x

free energy change is on the y

the y-intercept of this graph is delta H

the gradient of this graph is -S

53
Q

how do you work out the enthalpy of a solution

A

lattice dissociation + enthalpy of hydration of positive ios

54
Q

what is the enthalpy of hydration

A

the enthalpy change when 1 mole of aqueous ions is formed from gaseous ions

chemistry a level (24 decks)

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