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A Level Chemistry > Module 5 Section 2: Energy > Flashcards

Module 5 Section 2: Energy Flashcards

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1
Q

Combine the two half equations

A
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2
Q
A
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3
Q

Why are transition metals good oxidising and reducing agents

A

They are good at changing oxidation state
So they’ll give out or receive electrons
Also change colour so it’s easy to spot when the reaction is finished

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4
Q

How potassium manganate is used as an oxidising agent

A

Contains manganate (VII) ions (MnO4 -) in which manganese has an oxidation state of +7
Can be reduced to Mn2+ ions during a redox reaction

MnO4 - is purple and [Mn(H2O)6]2+ is pale pink, but looks colourless so the colour change in the reaction is purple to colourless

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5
Q

How is potassium dichromate used as an oxidising agent

A

Contains dichromate(VI) ions (Cr2O7 2-) in which chromium has an oxidation state of +6
They can be reduced to Cr+3 ions during a redox reaction

Cr2O7 2- is orange and [Cr(H2O)6]3+ looks green so colour change is from orange to green

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6
Q

How to work out the concentration of a reducing agent

A

Titrate a known volume of it against an oxidising agent of known concentration
Can work the other way

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7
Q

How to carry out redox titration with Fe2+ as a reducing agent and MnO4 - as a oxidising agent

A

Measure out a quantity of the Fe2+ reducing agent
Add dilute sulfuric acid to the flask (this is an excess to make sure there are enough H+ to allow oxidising agent to be reduced)
Gradually add the aqueous MnO4 - to the reducing agent using a burette, swirling the conical flask as you do so
Stop when the mixture in the flask just becomes tainted with the colour of the MnO4 - (end point) and record volume of oxidising agent added
Run more titrations and calculate mean volume of MnO4 -

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8
Q
A
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9
Q

How can iodine - sodium thiosulfate titrations be used

A

They are a way of finding the conc of an oxidising agent
The more concentrated an oxidising agent is, the more ions will be oxidised by a certain volume of it

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10
Q

How to prepare for a titration of iodine and sodium thiosulfate

A

Use a sample of oxidising agent to oxidise as much iodide as possible by measuring out 25cm3 of potassium iodate solution (KIO3 and oxidising agent)
Add this to excess acidified potassium iodide solution (KI) where the iodate(V) ions oxidise some of the iodide ions to iodine

Then find out how many moles of iodine have been produced by titrations the solution with a known conc of sodium thiosulfate (Na2S2O3)

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11
Q

Method of titration of iodine with sodium thiosulfate

A

Take flask containing solution of oxidised iodide ions (iodine and water)
Add sodium thiosulfate from a burette to the flask drop by drop
Colour fades to pale yellow (hard to see at end point) so add 2cm3 of starch solution to detect presence of iodine so solution goes dark blue showing there still iodine there
Add sodium thiosulfate one drop at a time until blue colour disappears
When this happens, all the iodine has just been reacted
Then calculate number of moles of iodine in the solution

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12
Q

How to do titrate calculation to find number of moles of iodine produced

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13
Q

How to calculate the concentration of oxidising agent using the original equation (6.6x10-4 moles of iodine and 25cm3 oxidising agent)

A
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14
Q

Definition of lattice enthalpy

A

Lattice enthalpy is the enthalpy chance when 1 mole of an ionic lattice is formed from its gaseous ions under standard conditions of 298K and 100kPa

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15
Q

What is lattice enthalpy

A

Ionic compounds can form regular structures called giant ionic lattice
The positive and negative ions are held together by electrostatic attractions
Lattice enthalpy is when gaseous ions combine to make a solid lattice

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16
Q

What does lattice enthalpy measure

A

This measures the ionic bond strength
The more negative the lattice enthalpy the stronger the bonding

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17
Q

What makes the lattice enthalpy more exothermic

A

The higher the charge on the ions the more energy is released when an ionic lattice forms
Due to stronger electrostatic forces between the ions
This means that the lattice enthalpy will be more negative
The smaller the ionic radii of the ions involved, the more exothermic the lattice enthalpy
Smaller ions have a higher charge density and the smaller ionic radii means the ions can sit closer together in the lattice

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18
Q

What is the born Haber cycle used for

A

Lattice enthalpy can’t be measured directly so this cycle is used to calculate the enthalpy change if you took another less direct route

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19
Q

Meaning of the born Haber cycle diagram

A

Route 1:
Formation of gaseous atoms
• changing the elements in their standard states into gaseous atoms (endothermic process)
Formation of gaseous ions
• changing gaseous atoms into positive and negative gaseous ions (endothermic)
Lattice formation
• changing gaseous ions into the solid ionic lattice (this is lattice enthalpy and is exothermic)

Route 2
Converts elements in their standard states directly to the ionic lattice
Includes one enthalpy change which is enthalpy or formation and is exothermic

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20
Q

Standard enthalpy change of atomisation

A

The enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in it’s standard state under standard conditions

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21
Q

What sign is in front of the enthalpy of atomisation

A

Standard enthalpy of atomisation is always an endothermic process so + value
Bonds are broken to form gaseous atoms

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22
Q

What is electron affinity

A

This is the opposite of ionisation energy
Electron affinity measures the energy to gain electrons
Always exothermic as the electron being added is attracted to nucleus

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23
Q

Definition of the first electron affinity

A

The enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions

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24
Q

Diagram example of successive electron affinities of oxygen

A
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25
Q

Why are second electron affinities endothermic

A

Second electron affinities are endothermic
This is because an electron is being gained by a negative ion which repels the electron away
Means that energy must be put in to force the negative electron onto the negative ion

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26
Q
A
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27
Q
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28
Q

How can electrochemical cells be made

A

Electrochemical cells can be made from two different metals dipped in salt solutions of their own ions and connected by a wire (the external circuit)

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29
Q

What are the two reactions taking place in an electrochemical cell

A

Oxidation and reduction

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30
Q

What happens in a zinc/copper electrochemical cell

A

Zinc loses electrons more easily than copper. So in the half-cell on the left, zinc (from the zinc electrode) is oxidised to form Zn2+ ions.
This releases electrons into the external circuit.
In the other half-cell, the same number of electrons are taken from the external circuit, reducing the Cu2+ ions to copper atoms

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31
Q

Which way do electrons flow in electrochemical cells

A

Flow from more reactive metal to least reactive metal

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32
Q

What would adding a volt meter into the circuit allow you to find

A

Shows the voltage between the two half-cells.
This is the cell potential or e.m.f.
Add a high resistance voltmeter

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33
Q

Other types of electrochemical cells

A

You can also have half-cells involving solutions of two aqueous ions of the same element, such as Fe2+/Fe3+
The conversion from Fe2+ to Fe3+, or vice versa, happens on the surface of a platinum electrode.

Electrochemical cells can also be made from non-metals.
For systems involving a gas (e.g. chlorine), the gas can be bubbled over a platinum electrode sitting in a solution of its aqueous ions (e.g. Cl-)

Platinum used as its inert and conducts electricity

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34
Q

What determines what direction the reaction goes in at each electrode

A

The reversible arrows show that both reactions can go in either direction.
Which direction each reaction goes in depends on how easily each metal loses electrons (how easily it’s oxidised).

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35
Q

What are electrode potentials

A

This measures how easily a metal is oxidised
A metal that’s easily oxidised (losing electrons) has a very negative electrode potential
One that’s harder to oxidise has a less negative or a positive electrode potential.

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36
Q

Example of how electrode potentials can determine the direction of the reaction in the cell

A

The zinc half-cell has a more negative electrode potential, so zinc is oxidised (the reaction goes backwards)
Copper is reduced (the reaction goes forwards)

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37
Q

What does the Θ symbol mean

A

The Θ symbol next to the E means it’s under standard conditions - 298K and 100 kPa.

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38
Q

Shorthand for writing electrochemical cell
Use example of Zn/Cu cell

A

The half-cell with the more negative potential goes on the left.
The oxidised forms go in the centre of the cell diagram

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39
Q

How to calculate overall cell potential

A
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40
Q

What voltage will the cell potential always be

A

The cell potential will always be a positive voltage
Because the more negative E° value is being subtracted from the more positive E° value.
E.g. the cell potential for the Zn/Cu cell - +0.34 - (-0.76) = +1.1 v

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41
Q

How do you measure the electrode potential

A

You measure the electrode potential of a half-cell against a standard hydrogen electrode

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42
Q

Definition of the standard electrode potential

A

The standard electrode potential, E°, of a half-cell is the voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode.

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43
Q

What are the standard conditions for electrochemical cells

A

Any solutions must have a concentration of 1.00 mol dm3 or be equimolar (i.e. contain the same number of moles of ions).
The temperature must be 298 K (25 °C).
The pressure must be 100 kPa

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44
Q

What side is the standard hydrogen electrode shown

A

The standard hydrogen electrode is always shown on the left
Doesn’t matter whether or not the other half-cell has a more positive value.
The standard hydrogen electrode half-cell has a value of 0.00 V.

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45
Q

How to work out whole cell potential while using standard hydrogen electrode

A

Whole cell potential = E°RHS -E°LHS
E°LHS is 0V so voltage reading will be equal to E°RHS
Reading could be positive or negative depending on the way the electrons flow

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46
Q

How can conditions effect the value of the electrode potential

A

Like regular reversible reactions the equilibrium position in a half-cell is affected by changes in temperature, pressure and concentration.
Changing the equilibrium position changes the cell potential.
Standard conditions are used to measure electrode potentials
Using these conditions means you always get the same value for the electrode potential and you can compare values for different cells

Shifting a position of equilibrium can increase or decrease the cell potential depending on whether the position shifts to produce more or less electrons

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47
Q

Overall equation for oxidation of iodide to iodine using iodate ions

A
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48
Q

Ionic equation for iodine reacting with sodium thiosulfate

A
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49
Q

What is a salt bridge

A

This is filter paper soaked in super saturated KNO3
Allows ions to flow through and balance out charges

50
Q

What type of electrode potential do more reactive metals have

A

The more reactive a metal is the more it wants to lose electrons to form a positive ion
More reactive metals have more negative standard electrode potentials

51
Q

What would happen in a cell containing magnesium and zinc

A

Magnesium is more reactive than zinc so it’s more eager to form 2+ ions than zinc is.
The list of standard electrode potentials shows that Mg2+/Mg has a more negative value than Zn2+/ Zn
In terms of oxidation and reduction, magnesium would reduce Zn2+ (or Zn2+ would oxidise Mg).

52
Q

What type of electrode would a more reactive non metal have

A

The more reactive a non-metal is, the more it wants to gain electrons to form a negative ion.
More reactive non-metals have more positive standard electrode potentials

53
Q

What would happen in a cell containing chlorine and bromine

A

Chlorine is more reactive than bromine so it’s more eager to form a negative ion than bromine is.
The list of standard electrode potentials shows that ½Cl2 /Cl- is more positive than ½Br /Br-.
In terms of oxidation and reduction, chlorine would oxidise Br (or Br would reduce CI)

54
Q

What does a more positive electrode potential mean on an electrochemical series

A

The left hand substances are more easily reduced
The right hand substance are more stable

55
Q

What does a more negative electrode potential mean on an electrochemical series

A

The right hand substances are more easily oxidised
The left hand substances are more stable

56
Q

Predict whether zinc will react with aqueous copper ions using the half equations in the series

A
57
Q

Why do you flip one of the equations round when balancing half equations
Which one do you flip

A

Half equations are written as reduction equations (e.g. in electrochemical series)
One will always need to switch in the direction of oxidation
E.g. Zn/ Zn2+ has a more negative electrode potential than Cu2+/Cu so the zinc will be oxidised (equation flips round)

58
Q

Predict whether copper will react with dilute sulfuric acid using the half equations in the series

A
59
Q

What do predictions using E° state

A

This only states if the reaction is possible under standard conditions

60
Q

Why may a prediction of whether two metals will react be wrong if the conditions are not standard

A

Changing the conc (or temp) of the solution can cause the electrode potential to change
E.g:

61
Q

Why may a prediction of whether two metals will react be wrong if the reaction kinetics are not favourable

A

The rate of reaction may be so slow that the reaction might appear not to happen
If a reaction has a high activation energy then this may stop it happening

62
Q

Why are standard conditions required at half cell reactions

A

The position of the redox equilibrium may change with conditions

63
Q

What will decreasing the conc of Mn+ do to the reaction

A

Decreasing the concentration of Mn+ shifts equilibrium to the left to make more electrons
This makes electrode potential more negative as there is more electrons

64
Q

Diagram of a standard hydrogen electrode

A

Monoprotic acid is used e.g. HCl or NO3
Electrode made of finely divided pieces of platinum (platinum black) which increases surface area

65
Q

Definition of standard electrode potential

A

The emf of a half cell connect to a standard hydrogen half cell under standard conditions of 298kPa, solution concentrations of 1.00 mol dm-3 and a pressure of 100kPa

66
Q

What we can tell from the metals using their electrode potential

A

If electrode potential positive then we know the metal is better at gaining electrons than hydrogen
If electrode potential is negative then the metal is better at releasing electrons than hydrogen

More reactive metal = more oxidised (better reducing agent) = more negative terminal

67
Q

How do energy storage cells work

A

Energy storage cells have been around for ages and modern ones work just like an electrochemical cell
Other name for a battery

68
Q

What was the nickel-iron cell used for

A

The nickel-iron cell was developed at the start of the 1900s and is often used as a back-up power supply because it can be repeatedly charged and is very robust

69
Q

How to work out voltage for energy storage cells

A

You can work out the voltage produced by these cells by using the electrode potentials of the substances used in the cell.

70
Q

Advantages of electrochemical cells

A

They are more efficient at producing energy than conventional combustion engines.
This is because energy is wasted during combustion as heat.

They produce a lot less pollution (such as CO2).
For hydrogen fuel cells, the only waste product is water.

71
Q

Drawbacks of electrochemical cells

A

The production of the cells involves the use of toxic chemicals, which need to be disposed of once the cell has reached the end of its life span.
The chemicals used to make the cells are also often very flammable.
E.g. lithium (commonly used in rechargeable batteries), is highly reactive and will catch fire if a fault causes it to overheat

72
Q
A

Have to double 2nd equation to the electrons balance those in the first equation

73
Q

Draw diagram for how fuel cells generate electricity from reacting a fuel with an oxidant

A
74
Q

What are primary cells

A

These are non rechargeable
The electrical charge is produced by oxidation and reductions at the electrodes
The reactions cannot be reversed
Eventually the chemicals will be used up, voltage will fall and the battery will go flat

75
Q

What are secondary cells

A

these cells are rechargeable
The cell reaction producing electrical energy van be reversed during recharging
The chemicals in the cell are then regenerated and the cell can be used again

76
Q

Common examples of secondary cells

A

Lead acid batteries used in car batteries
Lithium ion and lithium ion polymer cells used in appliances like laptops and mobile phones

77
Q

What are fuel cells

A

A fuel cell uses the energy from the reaction of a fuel with oxygen to create voltage
The fuel and oxygen flow into the fuel cell and products flow out, electrolyte remains in the cell
Fuels can operate continuously provided that the fuel and oxygen are supplied into the cell
Do not have be recharged

78
Q

What happens when an ionic lattice dissolves in water

A

The bonds between the ions break to give gaseous ions - this is endothermic.
The enthalpy change is the opposite of the lattice enthalpy

Bonds between the gaseous ions and the water are made - this is exothermic.
The enthalpy change here is called the enthalpy change of hydration.

79
Q

Definition of the enthalpy of hydration

A

Is the enthalpy change when 1 mole of gaseous ions dissolves in water

80
Q

Definition of the enthalpy change of solution

A

Is the enthalpy change when 1 mole of solute dissolves in water

81
Q

What are the energy changes in order for substances to dissolve

A

Substances generally only dissolve if the energy released is roughly the same, or greater than the energy taken in.
So soluble substances tend to have exothermic enthalpies of solution

82
Q

How to set out a born haber cycle for enthalpy change of solution

A

Put the ionic lattice and the dissolved ions on the top connect them by the enthalpy change of solution - this is the direct route

Connect the ionic lattice to the gaseous ions by the reverse of the lattice enthalpy.
The breakdown of the lattice has the opposite enthalpy change to the formation of the lattice

Connect the gaseous ions to the dissolved ions by the hydration enthalpies of each ion - this completes the indirect route

83
Q

Work out enthalpy change of solution for NaCl
Lattice enthalpy NaCl: -787 kJmol-1
Enthalpy hydration Na+: -406 kJmol-1
Enthalpy hydration Cl-: -364 kJmol-1

A
84
Q

Work out enthalpy change of solution for AgCl
Lattice enthalpy AgCl: -905 kJmol-1
Enthalpy hydration Ag+: -464 kJmol-1
Enthalpy hydration Cl-: -364 kJmol-1

A
85
Q

Is the enthalpy of hydration more or less for ions with a greater charge

A

Ions with a greater charge have a greater enthalpy of hydration.
lons with a higher charge are better at attracting water molecules than those with lower charges
Because the electrostatic attraction between the ion and the water molecules is stronger.
This means more energy is released when the bonds are made giving them a more exothermic enthalpy of hydration

86
Q

Is the enthalpy of hydration more or less for ions which are smaller

A

Smaller ions have a greater enthalpy of hydration.
Smaller ions have a higher charge density than bigger ions.
They attract the water molecules better and have a more exothermic enthalpy of hydration

87
Q
A
88
Q
A
89
Q

Alternative way of setting out born-Haber cycle

A
90
Q

How does changing ionic size and ionic charge affect lattice enthalpy

A

As you go down the group the ionic radius increases
This means the attraction between ions decreases so lattice is less negative
Melting point decreases

If ionic charge increases then the attraction between ions increases
This means that lattice enthalpy becomes more negative
Melting point increases

91
Q

What is entropy

A

Entropy is a measure of the number of ways that particles can be arranged and the number of ways that the energy can be shared out between the particles.

92
Q

What results in a higher entropy

A

The more disordered the particles are, the higher the entropy is.
A large, positive value of entropy shows a high level of disorder

93
Q

What things can affect entropy

A

Physical state
Number of particles

94
Q

How does physical state affect entropy

A

Solid particles are fixed in place so there is less randomness and lowest entropy
Gas particles are constantly moving and have the most random arrangements of particles so have the highest entropy

95
Q

How does physical state affect entropy

A

Solid particles are fixed in place so there is less randomness and lowest entropy
Gas particles are constantly moving and have the most random arrangements of particles so have the highest entropy

96
Q

How does the number of particles effect entropy

A

The more particles you’ve got, the more ways they and their energy can be arranged
In the reaction: N2O4(g) -> 2NO2(g) entropy increases because the number of moles increases

97
Q

What sort of state will substances usually react towards

A

Substances are more energetically stable when there’s more disorder
Particles will move to try and increase entropy

This is why some reactions are feasible even when the enthalpy change is endothermic

98
Q

Why is the reaction between sodium hydrogencarbonate and hydrochloric acid still feasible even though it is endothermic

A

This is due to an increase in entropy as the reaction produces CO2 gas and water
Liquids and gases are more disordered than solids and so have a higher entropy
This increase in entropy overcomes the change in enthalpy
Also favoured because it increases the number of moles

99
Q

When is a substance thermodynamically stable

A

When a substance reaches its maximum entropy state (its lowest energy state), it’s said to be thermodynamically stable.
This means it won’t react any further without the input of energy

100
Q

How to calculate entropy change

A

Must know the entropies of the products and reactants

101
Q

Units of entropy

A

J K-1 mol-1

102
Q

What entropy are more feasible reactions

A

Reactions with a positive value of △S are more likely to be feasible

103
Q
A
104
Q

Definition of standard entropy

A

The standard entropy S° of a substance is the entropy of one mole of a substance under standard conditions (100kPa and 298K)

105
Q

What does the tendency for a reaction to happen depend on

A

The entropy: △S
The enthalpy: △H
The temperature T
These combine to make the free energy change △G

106
Q

What does △G tell you

A

Says if a reaction is feasible or not
The more negative the value of △G the more feasible the reaction

107
Q

Formula for free energy change

A
108
Q
A
109
Q

What value does △G have to be in order for the reaction to be just feasible

A

△G = 0

110
Q

How to calculate the minimum temperature needed to make a reaction feasible

A
111
Q
A
112
Q

What are the limitations of predictions made for feasibility

A

A negative △G does not guarantee a reaction
It can indicate thermodynamic feasibility but does not take into account kinetics or rate or reaction
It may have a high activation energy so it will happen very slowly

113
Q
A
114
Q

What value will △G be when △H is negative and △S is positive

A

△G is always negative and the reaction is feasible

115
Q

What value will △G be when △H is positive and △S is negative

A

△G will always be positive and the reaction is not feasible

116
Q
A
117
Q
A
118
Q
A
119
Q

What does a reaction being feasible mean

A

They just happen by themselves without the addition of energy

120
Q

Explain the colour change of MnO4- to Mn2+

A

The Mn2+ is being made in the reaction and is so pale pink that it looks colourless
When the limiting reactant is used up, the MnO4- is not longer reduced to Mn2+ and remains in solution
This looks pale pink

A Level Chemistry (15 decks)

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