Unit 5 Chemistry- Redox Equations

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  • Created on: 22-06-11 21:21

Redox and Oxidation

A loss of electrons is called oxidation. A gain of electrons is called reduction.

OIL RIG

Reduction and Oxidation happen simultaneously- hence the term "redox" reaction

An oxidising agent accepts electrons and gets reduced

A reducing agent donates electrons and gets oxidised

Oxidation States

1.All  atoms are treated as ions for this, even if they're covalently bonded

2. Uncombined elements have an oxidation state of 0

3. Elements just bonded to identical atoms, like O2 and H2, have an oxidation state of 0

4. The oxidation state of a simple monatomic ion, e.g. Na+, is the same as it's charge

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Oxidation states continued...

5. In compounds or compound ions, the overall oxidation state is just the ion charge:

e.g. SO4 2-

Overall oxidation state = -2, O oxidation state = -2 (total = -8) (within an ion, the most electronegative element has a negative oxidation state (equal to it's ionic charge). Other elements have more positive oxidation states). so state of S = +6

6. The sum of oxidation states for  neutral compound is 0. 

e.g. Fe2O3

Overall oxidation state = 0, oxidation state of O = -2 (total = -6) so state of Fe = +3

7. Combined oxygen is nearly always -2 (except in peroxides where it is -1, and fluorides OF2 where it's+2 and O2F2 where it's +1, and in O2, where it is 0). In H2O, oxidation state of O = -2, but in H2O2, oxidation state of H has to be +1 (as H can only lose 1 electron) so oxidation state of O = -1. 

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More on oxidation states:

8. Combined hydrogen is +1, except in metal hydrides where it is -1, and H2 where it's 0.

e.g. in HF, oxidation state of H = +1, but in NaH, oxidation state of H = -1.

9. The oxidation state of a ligand is equivalent to the charge on the ligand. 

e.g. Oxidation state of CN- = -1 and the oxidation state of NH3= 0

If you see roman numerals in a chemical name, it is the oxidation number of the atom or group immediately before it. 

e.g. Copper in copper (II) sulfate has a state of +2

and Manganese in manganate (VII) ion (MnO4-) has an oxidation number of +7

The oxidation state of an tom will increase by one for each electron lost.

The oxidation state will decrease by 1 for each electron gained. 

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Half Equations

Can show oxidation or reduction

An oxidaton half equation can be combined with a reduction half equation to make a full equation (remember to make sure it balances).

Zinc metal displaces silver ions from silver nitrate solution to form zinc nitrate and a deposit of silver metal. 

The zinc atoms each lose 2 electrons (oxidation) = Zn (s) ---> Zn 2+ (aq) + 2e-

The silver ions each gin 1 electron (reduction) = Ag+ (aq) + e- ---> Ag (s)

Two silver nitrate ions are needed to accept the two electrons releases by each zinc atom. So you need to double the silver half equation before the two half-equations can be combined = 2 Ag + (aq) + 2e- ---> 2Ag (s).

Now the number of electrons lost and gained balances, so half equations can be combined: (e-s aren't included in the full equation):

Zn (s) + 2Ag 2+ (aq) ---> Zn 2+ (aq) + 2Ag (s)

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Balancing half equations

Acidified manganate (VII) ions can be reduced to Mn2+ by Fe2+ ions. The Fe2+ ions are oxidised to Fe3+.

Writing the equations:

1. Start with the ion being reduced: MnO4- (aq) ---> Mn2+ (aq)

2. Add 4 water molecules to balance the oxygen: MnO4- (aq) ---> Mn2+ (aq) + 4H2O (l)

3. Add H+ to balance hydrogen: MnO4- (aq) + 8H+ (aq) ---> Mn2+ (aq) + 4H2O (l)

4. Add e-s to balance charge: MnO4- + 8H+ + 5e- ---> Mn2+ + 4H2O (l)

5. Do the same for the other half equation: Fe 2+ (aq) ---> Fe 3+ (aq) + e-

6. Have to multiply the second equation by 5: 5Fe2+ (aq) ---> Fe3+ (aq) +5e-

7. Combine: MnO4-(aq) + 8H+(aq) + 5Fe2+(aq) --> Mn2+(aq) + 4H2O(l) + 5Fe3+(aq)

8.Check the charges balance

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Electrochemical Cells

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). The solutions are connected by a salt bridge made from filter paper soaked in KNO3 (aq). This allows ions to flow through and balance out the charges.There are always two reactions within an electrochemical cell- one's an oxidation and one's a reduction. So its a redox process.

Zinc/ copper electrochemical cell:

Zinc loses electrons more easily than copper. So in one half cell,  Zinc (from the zinc electrode) is oxidised to form Zn 2+ (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. 

So electrons flow through the wire from the most reactive metal to the least. A voltmeter in the external circuit shows the voltage between the two half-cells. This is the cell potential or e.m.f

You can also have cells involving solutions of two aqueous ions of the same element e.g. Fe2+ (aq) and Fe3+ (aq). The conversion of Fe2+ to 3+ and vice versa, happens on the surface of the electrode which is made of platinum which is an inert metal.

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Reactions at each electrode are reversible

Reactions that occur at each electrode in the zinc/ copper cell are:

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

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

The reaction can go in either direction. Which direction depends on how easily each metal loses electrons (i.e. how easily it's oxidised).

How easily a metal is oxidised is measured using electrode potentials. A metal thats easily oxidised has a very negative electrode potential, while one thats harder to oxidise has a less negative, or a positive electrode potential.

In the exam you will see tables, but the zinc half-cell has a more negative electrode potential, so zinc is oxidised (the reaction goes backwards), while the copper is reduced (the reaction goes forwards).

In drawing: the one with the more negative potential goes on the left, nd the oxides forms go in the middle: Zn (s) | Zn2+ (q) || Cu2+ (aq) | Cu (s)

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Electrode potentials

E* cell = (E* right hand side- E* left hand side). E* (supposed to be the circle thing) is the symbol for electrode potential.

The cell potential will always be a positive voltage, because the more negative E* value is being subtracted from the more positive value (right-left). 

Half-cell reactions are reversible, and just like any other reversible reactions, the equilibrium position is affected by changes in temperature, pressure + concentration. 

Changing the equilibrium position changes the cell potential. To get around this, standard conditions are used to measure electrode potentials- using these conditions means you always get the same value for the electrode potentil an you can compare values for different cells. 

Standard conditions are:

1. Any slution must have a concentration of 1.00 moldm-3

2. Temperature must be 298K (25oC)

3. The pressure must be 100kPa

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Standard Hydrogen Electrodes

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

BTW- the electrode is made of platinum because you can't has a gas electrode.

In the cell drawing, the platinum electrode is shown but it doesn't take part:

Pt | H2 (g) | H+ (aq) || Zn 2+ (aq) | Zn (s)

1. 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. 

The standard hydrogen electrode half-cell has an electrode potential of 0.00 V

The whole cell potential = E* right hand side - E* left hand side 

E* left hand side = 0.00 V, so the voltage reading will be equal to E* right hand side.

This reading could be positive or negative, depending on which way the electrons flow.

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Cells provide evidence that electrons are transfer

1. Redox reactions involve the transfer of electrons from one substance to another- or so the theory goes. 

2. You can use the theory to make a prediction, 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 reaction at the other. 

3. A current does flow (stick an ammeter in and measure). So the electrochemical cell provides evidence that electrons are transferred in redox reactions.

4. But an experiment doesn't prove a theory.

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Electrochemical Series

1. 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. 

e.g. Mg is more reactive than Zn, so it's more eager to form 2+ ions than Zn is. Mg has a more negative value so Mg would reduce Zn 2+ ( or Zn would oxidise Mg).

2. 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.

e.g. Cl is more reactive than Br, so it's more eager to form a negative ion. Cl is more positive, so Cl would oxidise Br, or Br would reduce Cl). 

More negative electrode potentials mean: The right-hand substances are more easily oxidised (they lose electrons more easily).The left-hand substances are more stable.

More positive electrode potentials mean: The left-hand substances are more easily reduced (they gain electrons more easily). The right-hand substances are more stable.

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Anti-clockwise rule

This predicts whether a reaction will occur and which direction it will go in. 

e.g. will zinc react with aqueous copper ions?

1. Write the two half-equations down, putting the one with the more negative standard electrode potential on top

2. In each equation, the oxidised state of the substance goes on the left and the reduced on the right. 

3. Then you draw some anti-clockwise arrows, these give you the direction of each half equation:

Zn 2+ (aq) + 2e- <--> Zn (s) as the top and Cu 2+ (aq) + 2 e- <--> Cu (s) as the bottom

Equation is: Zn (s) + Cu 2+ (aq) <--> Zn 2+ (aq) + Cu (s) so they do react

The cell potential is E* bottom - E* top

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More on cells

Electrochemical cells are used as batteries.

Some types of cell are rechargeable while others can only be used until they run out.

A common type of non-rechargeable cell is a dry cell alkaline battery. (Useful for things that don't use a lot of power or only used for short periods of time).

Non-rechargeable cells use irreversible reactions. It is not practical to reverse things 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. e.g. in a zinc-carbon dry cell the zinc node forms the casing of the battery, and so becomes thinner as the zinc is oxidised. 

Another reason they cannot be charged is because the ammonium ions (which act as an electrolyte) produce hydrogen gas, which escapes from the battery. Without the hydrogen, the ammonium ions cannot be reformed by reversing the reactions.

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Rechargeable cells

They use reversible reactions. Examples are lead-acid cells, nickel-cadmium or NiCad and the L ion or lithium ion. 

To recharge,  current is supplied to force the electrons to flow in the opposite direction around the circuit and reverse the reactions. This is possible because one of the substances in a rechargeable battery escape or are used up. 

Non-rechargeable batteries have advantaged are disadvantages:

1. Cost: They are cheaper than rechargeable, however they have to be replaced.

2. Lifetime: Will work for longer than rechargeable, but they have to be disposed of

3. Power: Can't supply as much power, no use in devices that require a lot e.g. mobiles

4. Use of resources and waste: Use more resources and create more waste. Both types of batteries can be recycled, but many are thrown into land-fill sites

5. Toxicity: Less likely to contain toxic metals- lead and cadmium. (Can contain mercury). Less hazardous if they leak out and pollute water sources. 

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Fuel cells

In most cells the chemicals that generate electricity are contained in the electrodes and the electrolyte that form the cell.

In a fuel cel, the chemicals are stored separately out side the cell and are fed in when electricity is required.

E.g. The hydrogen-oxygen fuel cell (used to power electric vehicles)

Hydrogen and oxygen gases are fed into two separate platinum- containing electrodes. 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: 2H2 ---> 4H+ + 4e-

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: O2 + 4H+ + 4e- ---> 2H2O

Overall: H2 and O2 react to make water: 2H2 + O2 --> 2H2O

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Advantages

Fuel cells don't need electrical recharging.

As long as hydrogen and oxygen are supplied, the cell will continue to produce electricity. 

Another benefit is that the only waste product is water- no toxic chemicals or CO2 emissions to deal with.

The downside is that you need energy to produce a supply of hydrogen and oxygen.

They can be produced from the electrolysis of water - i.e. by reusing the waste product from the fuel cell

But this requires electricity, which is normally generated by burning fossil fuels, so the process is not carbon neutral.

Also hydrogen is highly flammable so it needs to be handled carefully when stored or transported. 

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Comments

emilita.nxn

These are amazing! 

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