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Developing Metals Flashcards

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1
Q

Give 2 metals found in their pure state (not compounds)

A

Gold, Silver, Copper

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2
Q

What are pure, gold and silver like- physically?

A

Soft

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3
Q

How can these metals be made less soft?

A

Mixing with other metals to form alloys

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4
Q

What oxidation states can Iron have?

A

+2 and +3

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5
Q

What oxidation states can copper have?

A

+1 and +2

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6
Q

When filling up electronic shells, what fills up first 4s or 3d?

A

4s

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7
Q

What’s Chromium’s electronic configuration?

A

1s2 2s2 2p6 3s2 3p6 3d5 4s1(doesn’t obey usual rule as 4s isn’t fully filled before 3d)

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8
Q

What’s Copper’s electronic configuration?

A

1s2 2s2 2p6 3s2 3p6 3d10 4s1 (also doesn’t obey usual rule)

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9
Q

When d-block atoms form ions, where are the electrons lost from first?

A

4s sub-shell

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10
Q

Ions in the d-block have their outermost electrons in which sub-shell?

A

3d

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11
Q

Give 2 examples of transition metals as catalysts

A

Hydrogenation of alkenes is catalysed by nickel or platinum.
Manufacture of ammonia by Haber process is catalysed by iron
Manufacture of sulfuric acid is catalysed by vanadium (V) oxide
Alloys of platinum and Rhodium used in car’s catalytic converters

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12
Q

What are the two types of catalysts?

A

Heterogeneous and Homogeneous

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13
Q

What’s heterogeneous catalysis?

A

When the catalyst is in a different phase/state to the reactants.

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14
Q

What’s a Homogeneous catalyst?

A

A catalyst in the same phase/state to the reactants.

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15
Q

Give some examples of coloured glass uses (2)

A
  • stained glass windows
  • Worktops
  • Doors
  • Vases
  • Glasses
  • Bottles
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16
Q

When you see coloured glass, why is it visible as a colour?

A

The colour which has the frequency which passes through the glass, the rest are absorbed.

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17
Q

In pottery, why do you see the colour you do?

A

The colour seen has a wavelength which is reflected off the surface and the rest of the colours are absorbed

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18
Q

What wavelength region is visible light in?

A

400nm- 700nm

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19
Q

What colour has a 400nm wavelength?

A

Violet

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20
Q

What colour has a 700nm wavelength?

A

Red

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21
Q

If frequency is high, wavelength is?

A

short

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22
Q

What colour is Fe2+ in solution?

A

Green

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23
Q

What colour is Fe3+ in solution?

A

yellow (rust colored)

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24
Q

What colour is Cu2+ in solution?

A

Blue

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25
Q

What’s disproportionation?

A

In a reaction if a species is both reduced and oxidised, it’s said to be disproportionate

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26
Q

What’s required for light to be absorbed by an atom?

A

The energy of the light matches the energy gap between two energy states in the atom.

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27
Q

When violet is absorbed by a solution, what colour is the solution visible as?

A

Yellow

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28
Q

What’s a redox reaction?

A

A reaction where electrons are transferred- reduction and oxidation occur

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29
Q

How can you reduce a redox reaction into more simple parts?

A

An oxidation half equation and a reduction half equation

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30
Q

How would you write the full redox reaction for Acified manganate (VII) being reduced to Mn2+ by Fe2+ ions?

A

The half equation for the iron being oxidised is: Fe 2+ –> Fe 3+ + e-
However the manganate being reduced is harder to give a half equation: MnO4 - –> Mn 2+.
First balance the oxygens, by adding a water to the right side of the equation: MnO4 - –> Mn 2+ + 4H2O
Now balance the Hydrogens by adding H+ ions to the left side: MnO4 - + 8H+ –> Mn2+ +4H20
Then you need to balance the charges using electrons: MnO4 - + 8H+ + 5e- –> Mn2+ + 4H2O

Now you have to make sure the electrons produced in the iron half equation matches that gained in the Manganate (VII). (x all components of the Fe half equation by 5)
Final redox equation:
MnO4 - (aq) + 8H+ (aq) + 5Fe 2+ (aq) –> Mn2+ (aq) + 5Fe 3+ (aq) + 4 H2O. –long ting

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31
Q

How would you use a titration to work out how much acid is needed to neutralise a base?

A

-Measure out a known volume of alkali using a pipette, in to a flask
- Add some indicator
-Firstly, do a rough titration to get an idea of around where the end point is by adding acid and swirling the flask
- Then do an accurate titration by running the acid within 2cm3 of the end point found in the rough titration, then add it drop by drop and record the exact volume of acid required to neutralise the base
-Repeat a few times to get an accurate mean finding
(you can also do these titrations the other way round- adding base to acid)

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32
Q

What’s the purpose of redox titrations?

A

To find out how much oxidising agent is required to react exactly with a quantity of reducing agent- or vice versa.

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33
Q

What type of redox titration do we have to know?

A

Manganate (VII) ions (MnO4 -) as the oxidising agent

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34
Q

How would you do this manganate (VII) titration?

A
  • measure out a quantity of reducing agent e.g. Fe 2+ ions using a pipette and put in a conical flask
    -Then add some dilute sulfuric acid to the flask (in excess) this is for ensuring there’s H+ ions to allow the reduction of manganate
  • Now start adding the aqueous manganate to the flask using the burette, swirling the flask as you go.
    -Stop when the colour just starts to change and record the volume of manganate used (this is a rough titration)
  • Then do more accurate ones to get 2 or more findings within 0.10cm3 of each other
    (you can also do the titration with the reducing and oxidising agent in the other position (e.g. manganate in the flask)
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35
Q

How do you know when the manganate reaction is completed?

A

There’s a colour change from the colourless reducing agent to pink (because the manganate (VII) ions are purple)- the exact moment it goes baby pink is when the reaction is done (end point)

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36
Q

How would you use the titration results to find the concentration of the reagent (reducing agent)? - In the manganate (VII) titration

A

Work out how many moles of MnO 4- you added to the flask (conc x vol)= moles
Then look at the balanced equation for the reaction to find how many moles of reducing agent react with every mole of MnO 4- to work out the moles of reducing agent that were in the flask.
Then work out the moles that would be in 1dm3- this is the concentration

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37
Q

27.5cm3 of 0.0200 mol dm-3 of auqeous potassium manganate (VII) reacted with 25cm3 of acified iron (II) sulfate solution. calculate the concentration of Fe 2+ ions in the solution.

A
  • Calculate moles of MnO 4- = 0.0275 x 0.0200= 5.5x10-4 moles
    -Look at the balanced equation -
    MnO4 - + 8H+ + 5Fe 2+ –> Mn 2+ + 4H2O +5Fe3+
    so five moles of Fe2+ react with one mole of MnO4 -, therefore x moles by 5- 5.5x10-4 x 5 = 2.75 x 10-3
  • Then work out the concentration by dividing moles by volume = 27.5x 10-3/ 0.025 =
    0.110 mol dm-3
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38
Q

Name two pieces of equipment you can use to accurately measure the volume of a solution in a titration

A
  • pipette

- burette

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39
Q

How to you determine when the end point of a titration has been reached?

A

When the colour change first occurs

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40
Q

Why don’t you need to add an indicator when doing a redox with manganate (VII) ions?

A

Because the manganate ions are coloured and the reagent is colourless, so you can see when a colour change occurs

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41
Q

What is meant by the term Coordination Number?

A

The number of dative covalent (coordinate) bonds formed between the central metal ion and the ligand.

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42
Q

If an ion has 6 ligands bonded to it coordinately, what shape and bond angle has it got?

A

Octahedral + 90o

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43
Q

If a metal ion has 4 coordinately bonded ligands what shapes and bond angles could it have

A
  • Tetrahedral + 109.5o

- Square planar + 90o

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44
Q

What is happening in a copper/zinc electrochemical cell

A

The Zn metal is in Zinc ion solution (Zn 2+) and the Cu metal is in Copper ion solution (Cu 2+)
Zinc loses electrons more easily than copper, so the zinc half cell is oxidised - Zn (s) –> Zn 2+ (aq) + 2e-
These electrons travel through the electrical circuit towards the Copper electrode which is reduced- Cu 2+ (aq) + 2e- –> Cu(s)
The voltmeter measures the potential difference between both cells

45
Q

How do you set up an electrochemical cell?

A
  • Get a strip of each metal you’re investigating (electrodes)
  • Clean the surfaces using emery paper and clean any oil off with propanone
  • Place each electrode in a separate beaker with aqueous ions of the metals e.g. Zn 2+ with zinc electrode
  • Create a salt bridge by soaking filter paper in a salt solution and drape between both solutions so either side is in the solution
  • connect the electrodes to a voltmeter using wires and crocodile clips. - you’ll get a voltage reading.
46
Q

What’s the purpose of the salt bridge in an electrochemical cell?

A

For salt ions to flow through to compete cell (circuit)

47
Q

Each half electrochemical cell can be represented with..

A

.. a half equation (showing reduction or oxidation)

48
Q

Are reactions at each electrode reversible?

A

YES - so write with forward and backward arrows

49
Q

How are electrochemical equations written?

A

The forward reaction always shows the reduction and backwards oxidation

50
Q

Every half cell has its own..

A

Electrode potential

51
Q

The half equation of the more positive standard electrode potential goes

A

Forwards (where reduction occurs because it’s more positive so attracts electrons Reduction Is Gain)

52
Q

The half equation of the more negative standard electrode potential goes

A

Backwards (where oxidation occurs as it’s more negative so loses electrons Oxidation Is Loss)

53
Q

In a electrochemical cell where Zn/Zn 2+ has an electrode potential of -0.76 volts and the Cu/Cu 2+ has an electrode potential of +0.34. What’s being reduced and what’s being oxidised

A

Copper- reduced

Zinc- oxidised

54
Q

Does oxidation happen at the anode or the cathode?

A

Cathode

55
Q

What does the term Standard Electrode Potential mean?

A

The measurement of voltage under standard conditions where the half cell is connected to a standard hydrogen electrode.

56
Q

What are the standard conditions used when measuring standard electrode potential?

A
  • all solutions at a concentration of 1.00 mol dm-3
  • Temperature of 298K (25oc)
  • Pressure of 100kPa
57
Q

What’s the equation for the reaction at the hydrogen electrode?

A

2H+ (aq) + 2e- –> H2 (g)

58
Q

When drawing electrochemical cells, which goes on the left, and which goes on the right?

A

Oxidation happening shown on left
Reduction happening shown on right
e.g. zinc on left, copper on right

59
Q

When drawing the electrochemical cells when standard hydrogen is being used, what side is hydrogen drawn on?

A

ALWAYS LEFT

60
Q

What potential does the standard Hydrogen half cell have?

A

0v

61
Q

How do you calculate the cell potential (E cell) from standard electrode potentials?

A

take the more negative away from the more positive

62
Q

Whats the cell potential when Mg/Mg 2+ = -2.37v and Br/Br- = +1.09

A

1.09-(-2.37)= +3.46 v

63
Q

Is the cell potential always positive or always negative?

A

Always positive

64
Q

Do more reactive metals form more or less positive electrodes? +why?

A

less- they want to lose electrons to form positive ions

65
Q

Do more reactive non-metals form more or less positive electrodes? + why?

A

More- They want to gain electrons to form negative ions.

66
Q

The more negative the electrode potential for a half equation, the

A

Less likely the reaction is to happen going in this direction (reduction)

67
Q

The more positive the electrode potential for a half equation, the

A

More likely the reaction is to happen going in this direction (reduction)

68
Q

What’s a reducing agent?

A

Something which is losing electrons, and giving them to another species and be oxidised itself.

69
Q

What’s an oxidising agent?

A

Something which is gaining electrons- from another species, and being reduced itself.

70
Q

How do you work out if a reaction is feasible using electrode potentials?

A

Write out both the half equations (going in direction of reduction) with the more negative on top and more positive below.
From this draw anticlockwise arrows going round the equations.
Then balance out equations based on electrons
Then use the anti clockwise arrows to write full equation (without electrons) with chemicals at the base of arrows being the reactants and those at the head being the products. This is the only possible way the reaction can occur (feasible) anything else isn’t feasible.

71
Q

What, chemically, is happening during rusting?- explain and give equations

A
  • there’s two half equations:
    Fe 2+ (aq) + 2e- Fe (s) - more negative electrode potential, therefore going left
    2H2o (l) + O2 (g) + 4e- 4OH- - more positive, therefore going right
  • after balancing, the full equation is:
    2H20 (l) + O2 (g) + 2Fe (s) —> 2Fe 2+ (aq) +4OH -(aq)
    -Then the Fe2+ and the OH- ions form Iron (II) hydroxide precipitate -
    Fe 2+(aq) + 2OH- (aq) –> Fe(OH)2 (s)
    This is further oxidised by oxygen to Iron (III) hydroxide:
    2H2O (l) + O2 (g) + 4Fe(OH)2 (s) –> 4Fe (OH)3
    This gradually turns to hydrated iron (III) oxide- rust
72
Q

Briefly, what are the two ways to prevent rust?

A

Barrier e.g paint
Or
The sacrificial method

73
Q

What does using a barrier do to prevent iron from rusting?

A

it keeps out water and/or oxygen

74
Q

What’s the sacrificial method for preventing rusting of iron?

A

By coating the iron with a more reactive metal, this means the more reactive metal reacts with oxygen and water instead of the iron

75
Q

What two substances are required for rusting?

A

Water and Oxygen

76
Q

Which periodic table block are the transition metals found?

A

d

77
Q

Define a transition metal

A

a d-block element that can form at least one stable ion with an incomplete d-sub-shell

78
Q

Why aren’t Scandium and Zinc transition metals?

A

Scandium only forms one ion (Sc 3+) and its d-shell is empty

Zinc forms one ion (Zn 2+) and its d-shell is full

79
Q

What’s the electronic configuration of Iron?

A

1s2 2s2 2p6 3s2 3p6 4s2 3d6

80
Q

Define ligand

A

An atom/ion/molecule which donates a pair of electrons to a central transition metal ion to form a coordinate (dative-covalent) bond.

81
Q

Define complex/ complex ion

A

A central metal atom or ion surrounded by coordinately bonded ligands.

82
Q

What is ligand substitution?

A

When one ligand is swapped for another

83
Q

What are the 2 complexes of iron you should know?

A

[Fe(H2O)6] 2+

[Fe(H2O)6] 3+

84
Q

What are the 3 complexes of Copper you should know?

A

[Cu(H2O)6] 2+
[Cu(NH3)4] 2+
[CuCl4] 2-

85
Q

What’s a bidentate ligand?

A

A ligand with two lone pairs- can form two coordinate bonds

86
Q

What’s a monodentate/ Unidentate ligand?

A

A ligand with one lone pair, which can form one coordinate bond

87
Q

What’s a polydentate ligand?

A

A ligand with more than two lone pairs, can form as many coordinate bonds as its lone pairs/ structure allows.

88
Q

What’s the structure of Ethanedioate?

A

2 carbons bonded to each other each carbon as a double bonded oxygen and a single bonded oxygen- which has a lone pair.

89
Q

What sort of ligand is Ethanedioate?

A

Bidentate, because it’s got two lone pairs

90
Q

What does Iron (II) hydroxide look like?

A

Dark green precipitate

91
Q

What’s the formula of Iron (II) Hydroxide?

A

Fe(OH)2

92
Q

What’s the formula for Aqueous Iron (II) ions?

A

[Fe(H2O)6] 2+

93
Q

What’s the appearance of Aqueous iron (III) ions?

A

Pale- green solution

94
Q

What’s the formula of Iron (III) Hydroxide?

A

Fe(OH)3

95
Q

What’s the appearance of Iron (III) Hydroxide?

A

Red/orange precipitate

96
Q

What’s the formula of Copper (II) hydroxide?

A

Cu(OH)2

97
Q

What’s the appearance of Copper (II) hydroxide?

A

blue precipitate

98
Q

What’s the formula of aqueous copper (II) ions?

A

[Cu(H2O)6] 2+

99
Q

What’s the appearance of aqueous copper (II) ions?

A

Pale- blue solution

100
Q

What’s the formula of Copper (II) ammonia complex?

A

[Cu(NH3)4(H2o2)] 2+

101
Q

What’s the appearance of Copper (II) ammonia complex?

A

dark blue solution

102
Q

Does iron (II) and iron (III) form complexes with ammonia?

A

nope

103
Q

Why are transition metals such good catalysts?

A

They have varying oxidation states by gaining or losing d-orbital electrons, so can transfer electrons to speed up a reaction

104
Q

What’s good about using transition metals as catalysts?

A
  • good for industry

- good for environment as reduced energy required

105
Q

What’s bad about using transition metals as catalysts?

A

They can pose health risks as many of them and their compounds are toxic.

106
Q

Describe what happens when a transition metal acts a heterogeneous catalyst

A
  • The reactant molecules are attracted to the surface of the catalyst and adsorb on to it
  • The surface of the catalyst activates the molecules so they can react for easily
  • The products desorb off the surface, leaving space for new reactants to take their place.
107
Q

What two properties are required from heterogeneous catalysts?

A
  • Attract reactant strongly so they can hold to surface while reacting
  • Not attract the product molecules strongly, so they desorb and don’t poison the catalyst.
108
Q

How do homogeneous catalysts work?

A

By allowing a different reaction pathway at a lower activation enthalpy and providing an intermediate. The catalyst is always reformed to react again.

109
Q

What’s the maximum number of ethanedioate ligands that can complex with a metal ion?

A

3

chemistry (4 decks)

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