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Chemistry A2 UNIT 5 > Period 3 and Redox Equilibria > Flashcards

Period 3 and Redox Equilibria Flashcards

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1
Q

What happens when sodium and magnesium with react water?

A

Sodium is in Group 1, magnesium in Group 2.

When they react, sodium loses one electron to form Na+ ion, while magnesium loses to electrons to form Mg2+.

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2
Q

Why is sodium more reactive than magnesium?

A

It takes less energy to remove one electron than it does to lose two. So more energy (usually heat) is needed for magnesium to react.

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3
Q

How does sodium react with water? What are the products, what does this mean?

A

Reacts vigorously with COLD water, forming a molten ball on the surface, fizzing and producing H2 gas.

2Na(s) + 2H2O (l) -> 2NaOH(aq) + H2(g)

Produces sodium hydroxide, so forms a strongly alkaline solution (pH 12-14).

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4
Q

How does magnesium react with water? What are the products, what does this mean?

A

Reacts very slowly with cold water. You can’t see any reaction, but it forms a weakly alkaline solution (pH 9-10).

Mg(s) + 2H2O(l) -> Mg(OH)2 (aq) +H2 (g)

The solution is only weakly alkaline because magnesium hydroxide is not very soluble in water, so relatively few hydroxide ions are produced.

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5
Q

What does magnesium react much faster with? (than water)

A

steam to form magnesium oxide.

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6
Q

How do period 3 elements react with oxygen?

A

Form oxides when they react with oxygen.
Usually oxidised to their highest oxidation states- the same as their group numbers.
Sulphur is an exception- it forms SO2.

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7
Q

What is needed to make SO3?

A

High temp. and catalyst.

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8
Q

Equation and reaction for sodium reacting with oxygen.

A

2Na(s) + 0.5O2(g) -> Na2O(s)
Vigorous reaction in air.
Yellow flame

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9
Q

Equation and reaction for magnesium reacting with oxygen.

A

Mg(s) + 0.5O2(g) -> MgO(s)
Vigorous reaction in air.
Brilliant white flame

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10
Q

Equation and reaction for aluminium reacting with oxygen.

A

2Al(s) +1.5O2(g) -> Al2O3(s)
Slow reaction in air.
No flame

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11
Q

Equation and reaction for silicon reacting with oxygen.

A

Si(s) +O2(g) -> SiO2(s)
Slow reaction in air
no flame

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12
Q

Equation and reaction for phosphorus reacting with oxygen.

A

P4(s) + 5O2 -> P4O10(s)
Spontaneously combusts
Brilliant white flame

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13
Q

Equation and reaction for sulfur reacting with oxygen.

A

S(s) + O2(g) -> SO2(g)
Burns steadily
Blue flame

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14
Q

Describe and explain the melting points of the metal oxides (Na2O, MgO, Al2O3).

A

All have high melting points because they form giant ionic lattices.
The strong forces of attraction between each ion mean it takes a lot of heat energy to break the bonds and melt them.

MgO has a higher melting point than Na2O because magnesium forms 2+ ions which attract O2- ions more strongly than the 1+ sodium ions in Na2O.

Al2O3 has a lower melting point than expected because the 3+ ions distort the oxygen’s electron cloud making the bonds partially covalent.

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15
Q

Describe and explain the melting points of the non-metal oxides. (SiO2, P4O10, SO2)

A

SiO2 has a higher melting point than the other non-metal oxides because it has a giant macromolecular structure. Strong covalent bonds hold the structure together so lots of energy is needed to break the bonds and the melting temperature is high.

P4O10 and SO2 are covalent molecules. They have relatively low melting points because they form simple molecular structures. The molecules are held together by weak intermolecular forces (dipole-dipole and vdw), which take little energy to overcome.

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16
Q

How do the ionic oxides of the metals Na and Mg react with water?

A

dissolve in water to form hydroxides.
The solutions are both alkaline, but sodium hydroxide is more soluble in water, so it forms a more alkaline solution than magnesium hydroxide.

Na2O(s) + H2O(l) -> 2NaOH(aq) pH 12-14
MgO(s) +H2O(l) -> Mg(OH)(aq) pH 9-10

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17
Q

How do the simple covalent oxides of the non-metal sulfur and phosphorus react with water?

A

They form acidic solution.
All of the acids are strong and so the pH of their solutions is about 0-2.

P4O10(s) + 6H2O (l) -> 4H3PO4(aq) phosphoric(V) acid
SO2(g) + H2) (l) -> H2SO3 (aq) sulfurous acid or sulfuric (IV) acid.
SO3(l) +H2O(l) -> H2SO4(aq) sulfuric (VI) oxide.

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18
Q

How does silicon dioxide react in water?

A

Insoluble in water because of giant covalent structure.

However it will react with bases to form salts so it’s classed as acidic.

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19
Q

How does alumunium oxide react in water?

A

Insoluble in water, but it will react with acids and bases to form salts- act as an acid or a base (amphoteric).

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20
Q

What is the general equation for neutralising an acid with a base?

A

Acid + Base -> Salt + Water

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21
Q

Equations for reaction of sodium and magnesium oxides with acid.

A

Sodium and magnesium oxides are basic so will neutralise acids:

Na2O(s) + 2HCl(aq) -> 2NaCl(aq) +H2O(l)
MgO(s) + H2SO4(aq) -> MgSO4(aq) +H2O(l)

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22
Q

Equations for reaction of silicon, phosphorus and sulfur oxides with bases.

A

They are acidic so will neutralise bases:
SiO2(s) + 2NaOH(aq) -> Na2SiO3(aq) +H2O(l)
P4O10(s) + 12NaOH(aq) -> 4Na3PO4(aq) + H2O(l)
SO2(g) + 2NaOH(aq) -> Na2SO3(aq) + H2O(l)
SO3(g) + 2NaOH(aq) -> Na2SO4(aq) + H2O(l)

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23
Q

Aluminium oxides are amphoteric, what does this mean?

A

They can neutralise acids or bases.
Al2O3(s) + 3H2SO4(aq) -> Al2(SO4)3(aq) + 3H2O(l)
Al2O3(s) + 2NaOH(aq) +3H2O(l) -> 2NaAl(OH)4(aq)

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24
Q

What is a redox equation?

A

Reduction and oxidation happen simultaneously.

An oxidising agent gains electrons and is reduced, a reducing agent loses electrons and is oxidised.

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25
Q

OXIDATION

A

Loss of electron

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26
Q

REDUCTION

A

Gain of electrons

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27
Q

OXIDATION STATE

A

The oxidation state of an element tells you the total number of electrons it has donated or accepted.

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28
Q

What are the rules for assigning oxidation states?

A
  • Uncombined elements have an oxidation state of 0.
  • Elements just bonded to identical atoms have an oxidation state of 0.
  • The oxidation state of a simple monatomic ion is the same as its charge.
  • In compound ions, the overall oxidation state is the ion charge.
  • The sum of oxidation states for a neutral compound is 0.
  • Combined oxygen is always -2.
  • Combined hydrogen is always +1 (except in metal hydrides where it is -1 and H2 where it is 0).
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29
Q

How do you tell if an atom is oxidised or reduced in a reaction using oxidation states?

A

The oxidation state for an atom will increase by one for each electron lost. The oxidation state will decrease by one for each electron gained.

30
Q

If one of the ions in a half equation contains oxygen or hydrogen what might you have to add?

A

H2O and H+.

31
Q

How can electrochemical cells be made?

A

From two different metal dipped in salt solutions of their own ions connected by a wire (the external circuit).

32
Q

Describe an electrochemical cell made using copper and zinc.

A

A copper electrode is dipped in a solution of Cu2+ ions and a zinc electrode is dipped in a solution of Zn2+ ions. Zinc loses electrons more easily than copper.
So in the half cell on the left, zinc (from zinc electron) is oxidised to form Zn2+(aq) 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.

The solutions are connected by a salt bridge. This allows ions to flow through and balance out the charges- it completes the circuit.

33
Q

Example of a salt bridge.

A

strip of filter paper soaked in KNO3(aq)

34
Q

How do electrons flow through electrochemical cells?

A

Electrons flow through the wire from the most reactive metal to the least.

35
Q

What does the voltmeter show/

?

A

Shows the voltage between the two half cells- called cell potential or e.m.f or Ecell.

36
Q

Describe half cells involving solutions of two aqueous ions of the same element (Fe2+/Fe3+). Why is platinum used as the electrode?

A

You can make an electrochemical half-cell using solutions of Fe2+ and Fe3+ ions.
A platinum electrode is dipped into the solution .
Platinum is used as the electrode because it is inert and conducts electricity.
The conversion of Fe2+ to Fe3+, or vice versa happens on the surface of the electrode.
The direction of the conversion depends on the other half cell in the circuit.
If the other cell contains a metal that is less reactive than iron then Fe2+ will be oxidised to Fe3+ at the electrode.
But if the other cell contains a more reactive metal, Fe3+ will be reduced to Fe2+ at the electrode.

37
Q

What are the reactions at each electrode?

A

REVERSIBLE

38
Q

What are the half equations that occur at each electrode in the zinc/copper cell?

A

Zn2+(aq) +2e- Zn(s)

Cu2+(aq) +2e- Cu(s)

39
Q

What does the direction that the reversible reactions go in depend on?

A

How easily each metal loses electrons (how easily it is oxidised).

40
Q

How is how easily a metal is oxidised measured using?

A

Electrode potentials.
A metal that’s easily oxidised has a very negative electrode potential, while one that’s harder to oxidise has less negative electrode potential.

41
Q

How do you draw the shorthand way of representing electrochemical cells? (Zn/Cu)

A

Zn(s) | Zn2+ || Cu2+(aq) | Cu(s)

42
Q

What are the conventions when drawing an electrochemical cell?

A
  1. The half cell with the more negative potential goes on the left.
  2. The oxidised forms go in the centre of the cell diagram.
43
Q

How do you write half equations from a cell diagram?

A

The positive electrode is on the right and the negative electrode is on the left.
reduction always happens at the positive electrode, so the half equation written forward.
e.g. Pb2+ +2e- Pb
Oxidation always happens at the negative electrode, so the half equation is written backwards.
e.g. Al Al3+ + 3e-

44
Q

How do you calculate standard electrical potential?

A

Ecell= Erhs - Elhs
Cell potential will always be a positive voltage, because the more negative E value is being subtracted from the more positive E value.

45
Q

Why do conditions effect the electrode potential?

A

Half cell reactions are reversible. So just like any other reversible reaction, the equilibrium position is affected by changes in temperature, pressure and concentration.
Changing the equilibrium position changes the cell potential.

46
Q

How do you get around conditions changing the electrode potential?

A

Standard conditions are used- 298K, 100kPa and 1.00M solution of ions.

47
Q

What does using standard conditions mean?

A

You always get the same value for the electrode potential and you can compare values for different cells.

48
Q

What is the hydrogen electrode used for?

A

To measure the electrode potential of a half cell against.

49
Q

Describe the hydrogen electrode.

A

Hydrogen gas is bubbled into a solution of aqueous H+ ions. A platinum electrode is used as a platform for the oxidation/reduction reactions.

Pt | H2(g) | H+(aq) || Zn2+(aq) | Zn(s)

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

50
Q

STANDARD ELECTRODE POTENTIAL

A

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

51
Q

Why can standard hydrogen electrodes be used to calculate standard electrode potentials?

A

Standard hydrogen electrode half-cell has an electrode potentials of 0.00V.

52
Q

How do cell provide evidence that electrons are transferred?

A
  1. Redox reactions involve the transfer of electrons from one substance to another but you can’t see the electrons move. THEORY
  2. What you can do is use the theory to make a prediction, and then test the prediction with an experiment. E.g. you could predict that a current will flow between the electrodes of an electrochemical cell if an oxidation reaction happens at one electrode and a reduction at the other electrode.
  3. And a current does flow. So the electrochemical cell provides evidence that electrons are transferred in redox reactions.
  4. EXPERIMENT DOE SNOT PROVE THEORY.
53
Q

What is an electrochemical series?

A

Long list of electrode potentials for different electrochemical half cells.
The electrode potentials are all written in order from most negative to positive. Half equations are always written as reduction reactions.

54
Q

The more reactive a metal is…

A

The more it wants to lose electrons to form a positive ion.

More reactive metals have more negative electrode potentials. (easily oxidised).

55
Q

The more reactive a non-metal is…

A

The more it wants to gain electrons to form a negative ion. More reactive non-metals have more positive standard electrode potentials, (more easily reduced).

56
Q

What happens when two half equations are put together in an electrochemical cell?

A

The one with the more -ve electrode potential goes in the direction of oxidation (backwards) and the one with the more +ve electrode potential goes in the direction of reduction (forwards).

57
Q

How do you use the anti-clockwise rule to predict whether a redox reaction will happen and to show which direction it will go in?

A
  1. Find the two half equations for the redox reaction, and write them both out as reduction reactions.
  2. Use an electrochemical series to work out which half-equation has the more negative electrode potential.
  3. Put half-equation with more negative electrode potential on top of the other one.
  4. Draw on two anti-clockwise arrows- one going from the products of the top equation to the reactants of the top equation and one going from the reactants of the bottom equation to the products of the bottom equation.
  5. If you mix the substances at the non-pointy ends of the arrows a redox reaction will occur- the arrows show the direction that the half equations will go. But if you use any other combination of reactants there will be no reaction.
58
Q

What do non-rechargeable cells use?

A

Irreversible reactions

59
Q

What is a common type of non-rechargeable cell?

A

A dry cell alkaline battery

60
Q

Where are dry cell alkaline batteries usually found?

A

TV remote controls, torch, smoke alarms

Useful gadgets that don’t use a lot of power or are only used for short periods of time.

61
Q

What is an example of a dry cell alkaline battery? What are the half equations? How is the cell drawn?

A

Zinc-carbon dry cell batteries have a zinc anode and a mixture of manganese dioxide and carbon for a cathode.
In between the electrodes is a paste of ammonium chloride, which acts as an electrolyte.

Zn(s) -> Zn2+(aq) +2e- E= -0.76V
2MnO(s) +2NH4+(aq) +2e- -> Mn2O3(s) + 2NH3(aq) + H2O(l) E= +0.75V

Zn(s) | Zn2+(aq) || MnO2(s) | Mn2O3(s)

Ecell = +0.75 - (-0.76) = +1.51V

62
Q

Why do the half cell equations in a dry cell alkaline battery not have reversible arrows?

A

Not practical to reverse them in a battery. They can be made to run backwards under the right conditions, but trying to do this in a battery can make it leak or explode.
This is because the zinc anode forms the casing of the battery, so becomes thinner as zinc is oxidised.

63
Q

What is another reason why dry cell alkaline batteries cannot be recharged?

A

Ammonium ions would produce hydrogen gas, which would escape from the battery. Without hydrogen , the ammonium ions couldn’t be reformed by reversing the reaction.

64
Q

What do rechargeable cells use?

A

Reversible reactions

65
Q

Where are rechargeable batteries found?

A

Mobile phones, laptops, cars, etc.

66
Q

What is an example of a rechargeable cell? Half equations/cell diagram.

A

LEAD-ACID CELLS (used in car batteries).
Normally consist of 6 cells connected in a series.
Each cell is made up of a lead (IV) dioxide anode and a lead cathode immersed in sulphuric acid electrolyte.
Both electrodes end up coated in lead (II) sulfate.

The half equations are:
Pb(s) + SO42-(aq) PbSO4(s) +2e- E= -0.36V
PbO2(s) + SO4(aq) + 4H+ +2e- PbSO4(s) +2H2O(l)
E= +1.69V

Pb(s) | PbSO4(s) || PbSO4(s) | PbO2(s)

+1.69 - (-0.36) =+2.05

67
Q

What are two other types of rechargeable battery?

A

NiCad (nickel-cadmium)
and L ion (lithium ion),
To recharge these batteries a current it supplied to force electrons to flow in the opposite direction around the circuit and reverse the reactions. This is possible because none of the substances in a rechargeable battery escape or are used up.

68
Q

What are the pros and cons of non-rechargeable cells?

A

COST: non-rechargeable batteries are cheaper than rechargeable to buy. However, non-rechargeable batteries have to be replaced every time they run out, so rechargeables are cheaper in the long run.

LIFETIME: A non-rechargeable battery will usually work for longer than a rechargeable battery. But once a rechargeable battery has run out, you can recharge it and use it again whereas non-rechargeables have to be disposed of.

POWER: Non-rechargeable batteries can;t supply as much power as rechargeables, so are no use in devices that use a lot of power- like a mobile phone or a laptop.

USE OF RESOURCES AND WASTE: More non-rechargeable batteries are produced because they can only be used once, which uses more resources and means they create more waste than rechargeables. Both types of battery can be recycled and the metals in them recovered to use again, but often we just chuck them in the bin and they end up in landfill.

TOXICITY: non-rechargeable batteries are less likely to contain the toxic metals lead and cadmium (although they may contain mercury), so they’re less hazardous in landfill if the contents leak out and pollute water sources.

69
Q

How are fuel cells different to most cells?

A

In most cells the chemicals that generate electricity are contained in the electrodes and the electrolyte that form the cell. In a fuel cell the chemicals are stored separately outside the cell and fed in when electricity is required.

70
Q

What is an example of a fuel cell?

A

The hydrogen-oxygen fuel cell, which can be used to power electric vehicles.

Hydrogen and oxygen gases are fed into two separate platinum containing electrodes. These electrodes are usually made by coating a porous ceramic material with a thin layer of platinum, rather than using solid platinum rods. This is cheaper and it provides a larger surface area so the reactions go faster.

The electrodes are separated by an ion-exchange membrane that allows protons (H+ ions) to pass through it, but stops electrons going through it.

Hydrogen is fed to the negative electrode. The reaction that occurs is:
H2 -> 2H+ + 2e-

The electrons flow from the negative electrode through an external circuit to the positive electrode. The H+ ions pass through the ion-exchange membrane towards the positive electrode. Oxygen is fed to the positive electrode. The reaction here is:
O2 + 4H+ +4e- -> 2H2O

The overall effect is that H2 and O2 react to make water :
2H2 + O2 -> 2H2O

71
Q

What are the pros and cons of fuel cells?

A

PROS:

  • Don’t need electrical charging. As long as hydrogen and oxygen are supplied, the cell will continue to produce electricity.
  • Only waste product is water, no nasty toxic chemicals to dispose of and no CO2 emissions from the cell itself.

CONS:
-Need energy to produce a supply of hydrogen and oxygen. They can be produced from the electrolysis of water (using waster product) but this requires electricity, and this electricity is normally generated by burning fossil fuels.
Process isn’t usually carbon neutral.
-Hydrogen is highly flammable so needs to be handled carefully when stored or transported.

Chemistry A2 UNIT 5 (5 decks)

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