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A-Level Chemistry > Chemistry 5 - Physical Chemistry and Transition Elements > Flashcards

Chemistry 5 - Physical Chemistry and Transition Elements Flashcards

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1
Q

Define rate of reaction.

A

The change in the concentration of a reactant or a product per unit time.

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2
Q

At a given instant, how could you calculate the rate of reaction?

A

Rate of reaction=change in concentration of reactant or product/time

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3
Q

If the order is zero with respect to a reactant, what does that mean?

A

Changing the concentration of the reactant has no effect on the rate.

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4
Q

What is the effect of concentration on rate in a first order reaction?

A

Rate is directly proportional to the concentration.

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5
Q

What is the effect of concentration on rate in a second order reaction?

A

Change in rate=Change in concentration squared

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6
Q

Write a generic rate equation and state what each term means.

A

Rate=k[X]ˣ[Y]ʸ
Where:
k=The rate constant for the reaction
[X[=Concentration of reactant species X
[Y]=Concentration of reactant species Y
x=The order of reaction with respect to X
y=The order of reaction with respect to Y

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7
Q

Do zero order reactants appear on a rate equation? Why?

A

No, because they have no effect on the rate.

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8
Q

How is the overall order of a reaction calculated?

A

The sum of individual order.

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9
Q

How would you calculate the units of the rate constant?

A

Units of rate are moldm⁻³s⁻¹ and the units of concentration are moldm⁻³.
Rearrange the rate equation to get k=.
Sub in the units and cancel out to get the units of k.

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10
Q

When the overall order is 3, what would the units be for the rate constant?

A

mol⁻²dm⁶s⁻¹

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11
Q

How could you measure the rate of reaction experimentally?

A

If there is a colour change use a colorimeter at suitable intervals.
If gas is evolved, use a gas syringe to collect the volume of gas evolved, or measure the change in the mass of the reaction mixture.

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12
Q

What does half-life mean?

A

The time taken for concentration of a reactant to decrease by half.

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13
Q

What is the symbol for half-life?

A

t₁/₂

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14
Q

What is the relationship between first order reactions and half-life?

A

First order reactants have constant half-lives.

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15
Q

What is the relationship between zero order reactions and half-life?

A

Zero order reactants have half-lives that decrease by half each time.

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16
Q

What is the equation that is used to determine the rate constant using the half-life in the first order reaction?

A

k=ln(2)/t₁/₂

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17
Q

How would you draw a rate concentration graph?

A

Plot [A] against time, draw tangents are different values then draw a secondary graph of rate against [A].

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18
Q

What is the relationship between rate and time?

A

Rate∝1/t

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19
Q

How would you determine the rate constant from a rate concentration graph of first order?

A

k=rate/concentration

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20
Q

What is a rate determining step?

A

The slowest step in a reaction with multiple steps.

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21
Q

How does the rate determining step link to the species involved in the rate equation?

A

Any species involved in the rate determining step appear in the rate equation. Species only involved after the rate determining step do not appear in the rate equation.

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22
Q

For a reactant in the rate equation, what indicates how many molecules of that reactant are involved in the rate determining step?

A

The order of the reactant.

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23
Q

What affects the value of the rate constant for a given reaction?

A

Temperature and nothing else.

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24
Q

What is the effect of a 10°C temperature increase on the rate of reaction, roughly?

A

It doubles the rate of reaction.

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25
Q

What is the Arrhenius equation? What does each term mean?

A 
k=Ae(-Eₐ/RT)
Where:
k=Rate constant for reaction
A=Pre-exponential factor (the number of collisions between reactant molecules)
e=Mathematical quantity
R=Gas constant
T=Temperature in Kelvin
Eₐ=Activation energy for the reaction in Joules
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26
Q

How can you convert the Arrhenius equation into a useful form for plotting a graph?

A

In(k)=(-Eₐ/RT)+ln(A)

Graph of ln(k) against 1/T is a straight line, with the gradient -Eₐ/R and a y-intercept of ln(A).

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27
Q

What does it mean when a reaction at equilibrium?

A

The rate of forward and reverse reactions are equal.

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28
Q

What are the methods that can be used to measure equilibrium?

A

Measure the change in colour or colour intensity using a colorimeter.
Use a pH probe.
Measure electrical conductivity.
Titration- this is used when equilibrium is slow so that titration can be completed before there is much change in equilibrium mixture.

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29
Q

What does mole fraction mean?

A

The amount of a given component in a given reaction mixture.

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30
Q

Write the equation used to calculate mole fraction.

A

Mole fraction=number of moles of substance A/total number of moles of all substances

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31
Q

What does partial pressure mean?

A

The pressure exerted by a single species in a reaction vessel.

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32
Q

What is the symbol for partial pressure?

A

Ρ

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33
Q

Write the equation used to calculate partial pressure.

A

Partial pressure=mole fraction×total pressure

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34
Q

What is the relationship between concentration of a substance and its partial pressure?

A

Concentration of a substance is proportional to its partial pressure.

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35
Q

What are the three common units of pressure used in chemistry?

A

Pascals (Pa)
Atmospheres (Atm)
Newtons per square metre (Nm⁻²)

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36
Q

What type of brackets are used in K꜀ expressions?

A

Square brackets.

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37
Q

What type of brackets are used in Kₚ expressions?

A

Round brackets.

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38
Q

A reaction is represented by aA(g)+bB(g)⇌cC(g)+dD(g), how would you calculate Kₚ for the system?

A

Kₚ=(pCᶜpDᵈ)/(pAᵃpBᵇ)
Where
pA=Partial pressure of A
a=Number of moles of A

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39
Q

How would you calculate the units for Kₚ?

A

Write out the units for the partial pressures in the same arrangement as the Kₚ equation then cancel out/multiply together.
It is usually in Pa, kPa, atm etcetera, do not change units.

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40
Q

What is the symbol of equilibrium constant?

A

K

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41
Q

What does it mean when K is greater than 1?

A

Reaction favours products.

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42
Q

What does it mean when K is less than 1?

A

Reaction favours reactants.

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43
Q

What is the effect of increasing temperature on K?

A

Equilibrium shifts to the direction of endothermic reaction.

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44
Q

What is the effect of decreasing temperature on K?

A

Equilibrium shifts to the direction of exothermic reaction.

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45
Q

What is the only factor that affects K?

A

Temperature.

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46
Q

What is the effect of increasing temperature on K for a forward endothermic reaction?

A

K increases as temperature increases.

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47
Q

What is the effect of increasing temperature on K for a forward exothermic reaction?

A

K decreases as temperature increases.

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48
Q

What effect does catalyst have on K?

A

Catalysts does not affect the position of equilibrium, they only have an effect on the rate of reaction.

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49
Q

Define a Bronsted-Lowry acid.

A

A proton donor.

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50
Q

Define a Bronsted-Lowry base.

A

A proton acceptor.

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51
Q

Define a Lewis acid.

A

An electron pair acceptor.

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52
Q

Define a Lewis base.

A

Electron pair donor.

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53
Q

What ions cause a solution to become acidic? Give their names and formula.

A

H⁺ (Hydrogen ion) or, more accurately, H₃O⁺ (Oxonium ion), as protons react with H₂O to form it.

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54
Q

What ion causes a solution to be alkaline?

A

⁻OH (hydroxide ion)

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55
Q

Write an equation for the ionisation of water.

A

2H₂O(l)⇌H₃O⁺(aq)+⁻OH(aq)

Or H₂O(l)⇌H⁺(aq)+⁻OH(aq)

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56
Q

Give an example of a monobasic acid.

A

HCl

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57
Q

Give an example of a dibasic acid.

A

H₂SO₄

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58
Q

Give an example of a tribasic acid.

A

H₃PO₄

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59
Q

Identify the acid base pairs for the reaction below;

CH₃COOH+H₂O⇌CH₃COO⁻+H₃O⁺

A 
Pair 1:
•Acid: CH₃COOH
•Base: CH₃COO⁻
Pair 2:
•Acid: H₃O⁺
•Base: H₂O
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60
Q

Define strong acid.

A

Acids that dissociate completely.

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61
Q

Give some examples of strong acids.

A

Hydrochloric acid
Sulfuric acid
Nitric acid

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62
Q

What is the difference between concentrated and strong?

A

Concentrated means that there are many moles per dm³, strong refers to the amount of dissociation.

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63
Q

Define weak acids.

A

Acids that only partially dissociate.

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64
Q

Give some examples of weak acids.

A

Methanoic acid

Any organic acid

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65
Q

What is constant that is used to measure the extent of acid dissociation called?

A

Acid dissociation constant.

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66
Q

What is the symbol of the acid dissociation constant?

A

Kₐ

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67
Q

Write the acid dissociation constant expression.

A

For acid HA, HA⇌H⁺+A⁻

Kₐ=([H⁺][A⁻])/[HA]

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68
Q

What does a large Kₐ value mean?

A

The larger the Kₐ, the greater the extent of dissociation.

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69
Q

Write the equation used to convert Kₐ into pKₐ.

A

pKₐ=-log₁₀(Kₐ)

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70
Q

Write the equation used to convert pKₐ into Kₐ.

A

Kₐ=10⁻ᵖᴷᵃ

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71
Q

What is the relationship between pKₐ and the strength of the acid?

A

The smaller the pKₐ, the stronger the acid.

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72
Q

Write the equation used to convert the concentration of H⁺ into pH.

A

pH=-log₁₀[H⁺]

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73
Q

Write the equation used to convert pH into the concentration of H⁺.

A

[H⁺]=10⁻ᵖᴴ

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74
Q

Why is a pH scale useful compared to concentration of H⁺?

A

The pH scale allows a wide range of H⁺ concentration to be expressed as simple positive values.

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75
Q

What is the relationship between pH and [H⁺]?

A

A high pH value means a small [H⁺].

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76
Q

If two solutions have a pH difference of 1, what is the difference in [H⁺[?

A

A factor of 10.

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77
Q

[H⁺] of a strong acid is equal to what?

A

[H⁺]=[HA]

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78
Q

Write the equation used to calculate [H⁺] of weak acids.

A

[H⁺]=√(Kₐ×[HA])

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79
Q

What is the assumption made when calculating pH of weak acids?

A

it is assumed that the concentration of acid at equilibrium is equal to the concentration of acid after dissociation. This is because only very little of the acid dissociates.

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80
Q

Write the expression for ionic product of water, Kᵥᵥ.

A

Kᵥᵥ=[H⁺][OH⁻]

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81
Q

What is the units for Kᵥᵥ?

A

mol²dm⁻⁶

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82
Q

What is the value of Kᵥᵥ at 298K?

A

1.0×10⁻¹⁴

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83
Q

What physical factors affect the value of Kᵥᵥ? How do they affect it?

A

Temperature only- if temperature is increased, the equilibrium moves to the right so Kᵥᵥ increases and the pH of pure water decreases.

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84
Q

ndices of [H⁺] and [OH⁻] always adds up to what value?

A

-14

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85
Q

Define the term strong base.

A

A base that dissociates fully in water.

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86
Q

Give examples of some strong bases.

A

NaOH
KOH
Ca(OH)₂

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87
Q

Give an example of a weak base.

A

Ammonia

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88
Q

Write the equation used to calculate [H⁺] of strong bases.

A

[H⁺]=Kᵥᵥ/[OH⁻]

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89
Q

Define a buffer solution.

A

A mixture that minimises pH change on addition of small amounts of an acid or a base.

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90
Q

What are the two ways in which buffers can be made.

A
  • A weak acid and its conjugate base

* A weak acid and a strong alkali.

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91
Q

In which direction does the equilibrium shift when an acid is added to a buffer solution? Why?

A

Equilibrium shifts to the left because [H⁺] increases and the conjugate base reacts with the H⁺ to remove most of the H⁺.

92
Q

In which direction does the equilibrium shift when an alkali is added to a buffer solution? Why?

A

Equilibrium shifts to the right, because [OH⁻] increases and the small concentration of H⁺ reacts with OH⁻. To restore the H⁺ ions HA dissociates, shifting the equilibrium.

93
Q

Write the equation used to calculate [H⁺] of buffer solution.

A

[H⁺]=Kₐ×[HA]/[A⁻]

94
Q

Which buffer system maintains blood pH at 7.4? What happens when acid/alkali is added to blood?

A

H⁺+HCO₃⁻=CO₂+H₂O
Added OH⁻ reacts with H⁺ to form H₂O, then the equilibrium shifts to the left to restore lost H⁺.
Added H⁺ causes the equilibrium to shift to the right, removing excess H⁺.

95
Q

What is a titration?

A

The addition of an acid/base of known concentration to a base/acid to determine the concentration. An indicator is used to show that neutralization has occurred, as is a pH meter.

96
Q

Define the term equivalence point.

A

The point at which the exact necessary volume of base to just neutralise the acid, or vice-versa, has been added.

97
Q

What is the end point?

A

The point at which pH changes rapidly.

98
Q

What are the properties of a good indicator for a reaction?

A

A distinct sharp (not gradual) colour change- no more than one drop of acid/alkali needed for colour change, with a distinct colour change so it is obvious when the end point has been reached.
The end point must be the same as the equivalence point, otherwise the titration will give a wrong answer.

99
Q

What indicator would you use for a strong acid-strong base titration?

A

Phenolphthalein or methyl orange, but phenolphthalein is usually used as a clearer colour change.

100
Q

What indicator would you use for a strong acid-weak base titration?

A

Methyl orange.

101
Q

What indicator would you use for a strong base-weak acid titration?

A

Phenolphthalein.

102
Q

What indicator would you use from a weak acid-weak base titration?

A

Neither methyl orange or phenolphthalein is suitable, as neither give a sharp change at the end point.

103
Q

What colour is methyl orange in acid or in alkali?

A

Red in acid; yellow in alkali.

104
Q

What colour is phenolphthalein in acid or in alkali?

A

Colourless in acid; pink in alkali.

105
Q

What colour is bromothymol blue in acid or in alkali?

A

Yellow in acid; blue in alkali.

106
Q

Describe how to use a pH metre.

A

Remove the pH probe from the storage solution and rinse it with distilled water.
Dry the probe and place it into the solution with unknown pH.
Let the probe stay in the solution until it gives a settled reading.

107
Q

Define lattice enthalpy.

A

The formation of 1 mole of ionic lattice from gaseous ions under standard conditions.

108
Q

What does a more exothermic lattice enthalpy mean?

A

More exothermic means stronger ionic bonds.

109
Q

Why is it not possible to measure lattice enthalpy directly?

A

It is not possible to form 1 mole of ionic solid from its gaseous ions.

110
Q

Define enthalpy change of solurion.

A

The enthalpy change that takes place when 1 mole of a solute is completely dissolved in water under standard conditions.

111
Q

Define enthalpy change of hydration.

A

The enthalpy change that takes place when dissolving one mole of gaseous ions in water.

112
Q

What are the factors that impact the size of lattice enthalpy?

A
  • The size of ions involved
  • The charges on the ions
  • The ionic bond strength
113
Q

Which ions have more negative lattice enthalpy values; smaller or larger ions? Why?

A

Smaller ions because they can get closer hence more stronger attraction.

114
Q

Describe hydration.

A

When an ionic lattice is broken, the ions become part of the solution.
Positive ions get attracted towards slightly negative oxygen and negative ions get attracted towards slightly positive hydrogen.

115
Q

What are the factors that impact the magnitude of the enthalpy of hydration?

A
  • The size of the ion

* The charge on the ions

116
Q

Define entropy.

A

A measure of the dispersal of energy in a system which is greater when the system is more disordered.

117
Q

What is the symbol of entropy?

A

S

118
Q

Which is more disordered; solid or gas?

A

Gas.

119
Q

What is the unit of standard entropy?

A

JK⁻¹mol⁻¹

120
Q

How does temperature affect entropy?

A

The greater temperature particles have, the more energy they have and the more they move. Thus the arrangement of particles becomes more random. The more random the particle arrangement, the higher the entropy.

121
Q

When a solid ionic lattice is dissolved in solution, what happens to entropy?

A

Entropy increases because the ions are more disordered.

122
Q

How does a change in the number of gas molecules in a reaction affect entropy?

A

An increase in the number of gas molecules leads to an increase in entropy.
A decrease in the number of gas molecules leads to a decrease in entropy.

123
Q

Write the equation used to calculate entropy change.

A

ΔSᶿᵣₑₐ꜀ₜᵢₒₙ=ΣSᶿₚᵣₒdᵤ꜀ₜₛ-ΣSᶿᵣₑₐ꜀ₜₐₙₜₛ

124
Q

Write the Gibbs’ free energy equation and state what each symbol means.

A 
ΔG=ΔH-TΔS
Where:
ΔG=Gibbs' free energy (also called available energy)
ΔH=Enthalpy change
T=Temperature in kelvin
ΔS=Entropy change
125
Q

For a negative enthalpy change, and a positive entropy change, describe the Gibbs’ free energy and the feasibility of spontaneous change.

A

The Gibbs’ free energy is always negative and the reaction is feasible.

126
Q

For a positive enthalpy change, and a negative entropy change, describe the Gibbs’ free energy and the feasibility of spontaneous change.

A

The Gibbs’ free energy is always positive and the reaction is never feasible.

127
Q

For a positive enthalpy change, and a positive entropy change, describe the Gibbs’ free energy and the feasibility of spontaneous change.

A

The Gibbs’ free energy is negative at high temperatures and the reaction is feasible at high temperatures.

128
Q

For a negative enthalpy change, and a negative entropy change, describe the Gibbs’ free energy and the feasibility of spontaneous change.

A

The Gibbs’ free energy is negative at low temperatures and the reaction is feasible at low temperatures.

129
Q

For a reaction to occur spontaneously, does ΔG have to be positive or negative?

A

Negative.

130
Q

What are the limitations of the predictions of feasibility made by using ΔG?

A
  • The reaction may have a high activation energy.

* The rate of reaction may be very slow.

131
Q

Define oxidising agent.

A

A species that is reduced in a reaction and causes another species to be oxidised.

132
Q

Define reducing agent.

A

A species that is oxidised in a reaction and causes another species to be reduced.

133
Q

Define oxidation.

A

The loss of electrons.

An increase in the oxidation number.

134
Q

Define reduction.

A

The gain of electrons.

A decrease in the oxidation number.

135
Q

What happens in a redox reaction?

A

Electrons are transferred from one species to another.

One element is oxidised whilst the other is reduced.

136
Q

Why is 2CrO₄²⁻+2H⁺→Cr₂O₇²⁻+H₂O not a redox reaction?

A

Chromium is oxidised whereas hydrogen remains the same oxidation state (so no element is reduced).

137
Q

What are the half equations and the ionic equation for: SnO+Zn→ZnO+Sn?

A 
Half equations:
•Sn²⁺+2e⁻→Sn
•Zn→Zn²⁺+2e⁻
Ionic equation:
•Sn²⁺+Zn→Sn+Zn²⁺
138
Q

Define standard electrode potential.

A

The e.m.f. of a half cell compared with a standard hydrogen half cell measured at 298K with solution concentration of 1moldm⁻³ and a gas pressure of 100kPa.

139
Q

What happens when a rod of metal is dipped into a solution of its own ions?

A

An equilibrium is set up between the solid metal and the aqueous metal ions.

140
Q

Write a half-equation for zinc(s) to zinc(II).

A

Zn(s)⇌Zn²⁺(aq)+2e⁻

141
Q

Write a half-equation for copper(II) to copper(III).

A

Cu²⁺(aq)⇌Cu³⁺(aq)+e⁻

142
Q

What is a standard hydrogen half cell made of?

A
  • A solution containing 1moldm⁻³ (often an acid, like hydrochloric acid)
  • Hydrogen gas at 100kPa
  • Inert platinum electrode
143
Q

Why are hydrogen half cells used as a standard half cell?

A

It is easy to control its purity and reproducibility.

144
Q

Describe how to make a simple salt bridge?

A

Soak a piece of filter paper in an aqueous solution of KNO₃ or NH₄NO₃.

145
Q

Why are salt bridges necessary?

A

To complete the circuit by connecting the two solutions. This enables charge to be transferred between the half cells. They do not react with the electrodes.

146
Q

Why might you use other standard electrodes occasionally?

A

They are cheaper/easier/quicker to use and can provide just as good a reference.
Platinum is expensive.

147
Q

If an E° value is more negative, what does it mean in terms of oxidising/reducing power?

A

It is a better reducing agent (easier to oxidise).

148
Q

If an E° value is more positive, what does it mean in terms of oxidising/reducing power?

A

It is a better oxidising agent (easier to reduce).

149
Q

How do you calculate the e.m.f. of a cell from E° values?

A

E°꜀ₑₗₗ=E°ₚₒₛᵢₜᵢᵥₑ-E°ₙₑ₉ₐₜᵢᵥₑ

150
Q

When would you use a Platinum electrode? Why?

A

When both the oxidised and reduced forms of the metal are in aqueous solution. It is a good choice since Platinum is inert and a good conductor to complete the circuit.

151
Q

How would you predict if a reaction would occur?

A

Take the 2 half equations.
Find the species that is being reduced.
Calculate its E° value minus the E° value of the species being oxidised.
If E° overall is greater than 0.4V, the reaction will occur.

152
Q

What are the three main types of electrochemical cells?

A
  • Non rechargeable cells
  • Rechargeable cells
  • Fuel cells
153
Q

Describe how non rechargeable cells work.

A

They provide electrical energy until all the chemicals have reacted.

154
Q

Describe how rechargeable cells work.

A

Chemicals in the cell provide electrical energy. When recharging the reactions of the cells can be reversed.

155
Q

Give some examples of rechargeable cells.

A
  • Nickle and cadmium batteries
  • Lithium ion batteries
  • Lithium polymer batteries
156
Q

Explain why lithium is used in laptop batteries.

A

Lithium has a low density so the electrode is light and it is very reactive.

157
Q

What are the drawbacks of using lithium batteries?

A
  • They are toxic if ingested.

* Rapid discharge of current can cause fire or even explosions.

158
Q

Describe how fuel cells work.

A

The cell uses external supplies of fuel and an oxidant. These external supplies need to be continuously supplied.

159
Q

Modern fuel cells are based on what type of fuels?

A
  • Hydrogen

* Hydrogen rich fuels, such as methanol

160
Q

What are the reactions that take place at the two electrons in an alkaline hydrogen fuel cell?

A

2H₂+4OH⁻→4H₂O+4e⁻

O₂+2H₂O+4e⁻→4OH→

161
Q

What are the disadvantages of fuel cells?

A

Hydrogen is a flammable gas with a low b.p.which makes it hard and dangerous to store and transport which also makes it expensive to buy.
Fuel cells have a limited lifetime and use toxic chemicals in their manufacture.

162
Q

What is the reason that some cells cannot be recharged?

A

The reaction of the cell is not reversible, a product is produced that either dissipates or cannot be converted back into the reactants.

163
Q

Why might the e.m.f. of a cell change after a period of time?

A

Concentrations of the ions change as the reagents are used up.

164
Q

How can the e.m.f. of a cell be kept constant?

A

Reagents are supplied constantly, so the concentrations of the ions are constant; E° remains constant.

165
Q

Define a transition element.

A

A d-block element that forms at least one stable ion with partially filled d subshell.

166
Q

Where are the transition metals located in the periodic table?

A

In the middle block from Ti to Cu.

167
Q

What are the 2 elements in the d-block that are not considered as transition metals?

A

Scandium and zinc.

168
Q

Why are scandium and zinc not considered as transition metals?

A

Scandium only forms Sc³⁺, where the d-orbitals are empty.

Zinc only forms Zn²⁺ where the d-orbitals are full.

169
Q

What is the noble gas configuration of chromium?

A

[Ar]4s¹3d⁵

170
Q

What is the noble gas configuration of copper?

A

[Ar]4s¹3d¹⁰

171
Q

Which electrons do transition metals lose first when forming ions?

A

4s

172
Q

What are some characteristic physical properties of transition metals?

A
  • Metallic
  • High density
  • High melting and boiling point
  • Shiny
  • Good conductors of heat and electricity
173
Q

What are the characteristic chemical properties of transition metals?

A
  • Variable oxidation states, so they can take part in many redox reactions
  • Coloured compounds/ions in solution
  • Good catalysts
  • Form complex ions
174
Q

In potassium manganate(VII) what is the oxidation number of Mn?

A

Mn=+7

175
Q

What are the 2 main ways in which transition metals act as an effective catalyst?

A
  • They provide a surface on which reaction can take place

* They change oxidation states to form intermediates required for pathways with lower activation energy.

176
Q

Give some examples of transition metal catalysts and the process/reactions they catalyse.

A
  • Iron-Haber process
  • Vanadium(V) oxide-contact process
  • Nickel-hydrogenation of alkenes
  • Manganese(IV)oxide-decomposition of hydrogen peroxide
  • Copper sulfate-hydrogen production
177
Q

What is a complex ion?

A

Transition metal ion bonded to one or more ligands by coordinate bonds.

178
Q

Define a ligand.

A

Molecule or an ion that can donate a pair of electrons to the metal to form a coordinate bond.

179
Q

What is a coordinate bond?

A

A bond in which one atom provides both electrons required for bonding, also known as dative bonding.

180
Q

What does the coordination number indicate?

A

The number of coordinate bonds formed between the metal ion and its ligands.

181
Q

What is a mono/unidentate ligand?

A

A ligand that forms one co-ordinate bond to the central metal ion (one lone pair to donate)

182
Q

What is a bidentate ligand?

A

A ligand that forms two coordinate bonds to the central metal ion (2 lone pairs to donate).

183
Q

What is a multidentate ligand?

A

A ligand that forms three or more coordinate bonds to the central metal ion.

184
Q

Give some examples of common monodentate ligands.

A

Cl⁻, H₂O, NH₃, CN⁻

185
Q

Name the most common bidentate ligand.

A

Ethane-1,2-diamine

186
Q

How many co-ordinate bonds can ethanedioate can it form to a transition metal ion?

A

2 co-ordinate bonds.

187
Q

How many co-ordinate bonds can it form to a transition metal ion?

A

2 co-ordinate bonds.

188
Q

What is ethane-1,2-diamine’s shortened name? How many co-ordinate bonds does it form?

A

The shortened bond is en, 2 co-ordinate bonds.

189
Q

How many coordinate bonds does EDTA⁴⁻ form?

A

Six.

190
Q

EDTA is a chelating agent, what does that mean?

A

EDTA decreases the concentration of metal ions in the solution by binding them and forming complex ions.

191
Q

What ion is usually formed when a transition metal compound is dissolved in water? What shape is it?

A

An aqua ion, 6H₂O ligands around the central metal ion. An octahedral complex ion is formed.

192
Q

If a transition metal ion has 2 ligands, what shape is it usually?

A

Linear.

193
Q

If a transition metal ion has 4 ligands, what shape is it usually?

A

Tetrahedral.

194
Q

Name an exception to the general rule that an ion with 4 ligands is generally tetrahedral. What shape is it?

A

Platin is square planar and forms cisplatin.

195
Q

What shape is a complex ion if it has 6 ligands?

A

Octahedral.

196
Q

How can complex ions display E-Z or cis-trans isomerism? What shapes of ions does this apply to?

A

Ligands differ in the way in which they are arranged in space.
2 ligands of the same type can be on the same side of the metal ion (next to each other), which forms the E or cis isomer.
2 ligands of the same type can be on opposite sides of the metal ion (not next to each other), which forms the Z or trans isomer.
This applies to square planar and octahedral complex ions.

197
Q

What conditions are needed for a complex ion to display optical isomerism?

A

It usually applies to octahedral molecules with 2 or more bidentate ligands, so that the mirror images are non-superimposable.

198
Q

What is cis-platin used for? Describe its mechanism of action.

A

It is used as an anti cancer drug.

It binds to the DNA of fast growing cancer cells and prevents cell division, hence reducing the growth of cancer.

199
Q

Which metal ion is present in a haem group?

A

Fe²⁺

200
Q

What is the coordination number of a haem group?

A

6

201
Q

What are the ligands in a haem group?

A

4 nitrogen forming the porphyrin ring.
1 oxygen
1 globin

202
Q

Why is CO toxic?

A

CO is also able to co-ordinately bond to Fe²⁺, and bonds more strongly Fe²⁺ than O₂.
This stops O₂ from bonding to haemoglobin, so O₂ cannot be transported around the body.

203
Q

What colour is [Cu(H₂O)₆]²⁺ solution?

A

Pale blue.

204
Q

What colour is [Cu(NH₃)₄(H₂O)₂]²⁺ solution?

A

Deep blue.

205
Q

What colour is [CuCl₄]²⁻ solution?

A

Yellow.

206
Q

What colour is [Cr(H₂O)₆]³⁺ solution?

A

Dark green.

207
Q

What is a precipitation reaction?

A

A reaction where soluble ions in separate solutions are mixed to form an insoluble compound.

208
Q

What colour is Fe²⁺’s aqua ion?

A

Green.

209
Q

What colour is Fe³⁺’s aqua ion?

A

Pale brown.

210
Q

What colour is Cr³⁺’s aqua ion?

A

Violet.

211
Q

What colour is Mn²⁺’s aqua ion?

A

Pink.

212
Q

When Cu²⁺ reacts with NH₃ or OH⁻, what colour is the precipitate?

A

Blue.

213
Q

When Fe²⁺ reacts with NH₃ or OH⁻, what colour is the precipitate?

A

Green.

214
Q

When Mn²⁺ reacts with NH₃ or OH⁻, what colour is the precipitate?

A

Brown.

215
Q

When Cr³⁺ reacts with NH₃ or OH⁻, what colour is the precipitate?

A

Green.

216
Q

When Fe³⁺ reacts with NH₃ or OH⁻, what colour is the precipitate?

A

Brown.

217
Q

What colour is MnO₄⁻?

A

Purple.

218
Q

Write a half equation for the reduction of MnO₄⁻ to Mn²⁺.

A

MnO₄⁻+8H⁺+5e⁻→Mn²⁺+4H₂O

219
Q

Why are redox titrations with titrations with transition metal compounds said to be self-indicating?

A

They usually involve a colour change as the metal is changing oxidation state; though sometimes an indicator is still needed/useful.

220
Q

What colour is Cr₂O₇²⁻?

A

Orange.

221
Q

What colour is Cr³⁺?

A

Green.

222
Q

Write a half equation for the reduction of Cr₂O₇²⁻ to Cr³⁺

A

Cr₂O₇²⁻+14H⁺+6e⁻→2Cr³⁺+7H₂O

223
Q

Describe how to test for metal ions.

A
  • Fill half of the test tube with the sample.
  • Add aqueous ammonia or sodium hydroxide drop by drop.
  • Colour of precipitate indicates the ion present.
224
Q

Write the method used to test for ammonium ions.

A
  • Fill half the test tube with the sample.
  • Add sodium hydroxide and warm it gently.
  • For a positive result, a smelly gas is produced; damp red litmus paper turns blue.
225
Q

Why does lime water turn milky in the presence of carbon dioxide?

A

When carbon dioxide is bubbled in calcium hydroxide (lime water), calcium carbonate precipitate forms.

A-Level Chemistry (6 decks)

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