Rosemary Waters and Jennifer Colley

Present

The Origin of an OIL RIG

Figure 1 - A Sales pitch on how to understand redox. Image taken from http://www.redoxmarketing.com/red-ox-marketing.html.

Table of Contents

IONIC EQUATIONS..............................................................................................................3Precipitate reactions......................................................................................................................3Metal and acid reactions...............................................................................................................3Metal and carbonate reactions......................................................................................................3Acid and Base reactions.................................................................................................................3Displacement Reactions................................................................................................................3

REDOX reactions and Galvanic cells....................................................................................4Oxidation number rules..............................................................................................................4Oxidant and Reductant...............................................................................................................5Using KOHES...............................................................................................................................6Using the reactivity series to predict redox reactions..................................................................7The simplest electric cell – Lemon Power....................................................................................9The galvanic cell........................................................................................................................10

Constructing a Galvanic cell.........................................................................................................11Electronic conductors...............................................................................................................12

I

IONIC EQUATIONSIonic Equations are short hand equations which can summarise a general reaction in very simple terms.

Precipitate reactionsIn precipitate reactions the ions which form the precipitate are the only ions that are named

e.g. Pb2+(aq) + 2Cl-

(aq) PbCl2(s)

Metal and acid reactionsThe metal which is used is important, but the acid used can be any acid, but it is important to recognise that a salt and hydrogen are produced.

Magnesium + hydrochloric acid magnesium chloride + hydrogen

Magnesium + nitric acid Magnesium nitrate + hydrogen

Magnesium + sulphuric acid Magnesium sulphate + hydrogen

Any acid can be represented by H+ because an acid by definition is a proton donor.

Mg(s) + 2H+(aq) Mg2+

(aq) + H2(g)

This ionic equation represents an acid reaction with a metal.

Metal and carbonate reactionsIn this reaction the common facts are carbonates (carbonate ion CO3

-) and hydrogen ions (H+) from the acids.

Calcium carbonate + hydrochloric acid calcium chloride + carbon dioxide + water

CO32-

(aq) + 2H+(aq) CO2(g) + H2O(l)

Acid and Base reactionsA neutralisation reaction is one in which the H+ (hydrogen ion) from an acid, combines with the OH- ions from the base to form water.

Potassium hydroxide + hydrochloric acid salt + water

OH- (aq) + H+(aq) H2O(l)

Displacement ReactionsIt is important to be able to understand the reactivity of metals, as one of the determining factors is that a more reactive metal will displace or push out) a less reactive metal from a solution of its salt.

e.g. magnesium is more reactive than copper.

Magnesium in copper sulphate solution pushes the copper out of solution and takes its place in the solution, e.g. magnesium + copper sulphate copper + magnesium sulphate

Mg(s) + Cu2+(aq) Cu(s) + Mg2+

(aq)

REDOX reactions and Galvanic cellsIn some reactions there are oxidation and reduction reactions. One cannot happen without the other.

One substance will be oxidised and the other one will be reduced. Hence the term REDOX, which means reduction-oxidation.

Oxidation reactions are represented by:

Gain of oxygen Loss of electrons Loss of hydrogen Increase in oxidation number

Reduction reactions are represented by:

Loss of oxygen Gain of electrons Gain of hydrogen Decrease in oxidation number

Oxidation number rules1. The oxidation number of an atom in its elemental form is zero. E.g. Mg(s) is zero and N2(g) is

zero.2. The oxidation number of a simple ion is its valence value. E.g.Al3+ ion has an oxidation

number of +3 and S2- (sulphide) has an oxidation number of -2.3. The oxidation number of hydrogen is +1 when it occurs with another non metal e.g. HCl, H2O

and NH4+.

However when it is a hydride with metals, it is -1.4. The oxidation number of oxygen in a compound is -2. Exceptions, peroxide compounds H2O2

when oxygen is -1.5. In neutral compounds, the sum of all the oxidation numbers is zero.

e.g. Sulfuric acid H2SO4

H2+ S? O4

2-

H = +2O= - 8Therefore sulphur in this instance is +6Many elements have a number of different oxidation values, depending on the compound in which they are formed.e.g. S S = 0 SO2 S = +4 SO3 S = +6 H2SO3 S = +4 H2SO4 S = +6

H2S2O7 S = +6

Oxidant and ReductantThe oxidant is the substance which does the oxidising and it is in itself REDUCED.

The reductant is the substance which does the reducing and in is in itself oxidised

For exampleMg + 2H+ Mg2+ + H2

You can break this reaction into two half reactions. Mg Mg2+ + 2eCan you apply the rules with oxidation numbers? Can you apply the rules with loss and gain of electrons?

2H+ + 2e H2

Can you apply the rules with oxidation numbers? Can you apply the rules with loss and gain of electrons?

OIL RIG is the acronym

Oxidation is losing electrons; reduction is gaining electrons.

(a) Which is oxidised?(b) Which is reduced?(c) Which is the oxidant?(d) Which is the reductant?

More complex redox reactions

REDOX reactions can become more complex. There are two important oxidising agents that you need to understand. The strongest oxidising agents (which in turn are reduced) are potassium dichromate (Cr2O7

2-) and potassium permanganate (MnO4-).

There is a process to be followed and it uses the acronym KOHES

When Cr2O72- is reduced it produces Cr3+ and MnO4

- produces Mn2+

Cr2O72- Cr3+

MnO42- Mn2+

It is important for you to understand that this is only a half equation where a substance is reduced. There will be another half equation for an oxidation reaction.

Using KOHESK Balance the key element. Cr is the key element

Cr2O72- 2Cr3+

O Balance the oxygen on the LHS by adding enough water to the right hand side.

Cr2O72- 2Cr3+ + 7H2O

H Balance the hydrogen in the water on the RHS by adding enough H+ to the LHS

Cr2O72- + 14H+ 2Cr3+ + 7H2O

E Balance it electrically by adding electrons to one of the sides.

Cr2O72- + 14H+ + 6 2Cr3+ + 7H2O

(-2) (14+) (6+)

S Add the states

Cr2O72- (aq) + 14H+

(aq) + 6e 2Cr3+(aq) + 7H2O(l)

This is only one equation representing a reduction reaction.

Now for the oxidation reaction. We know this is an oxidation reaction as there is an increase in oxidation number.

Fe2+(aq) Fe3+

(aq) + e

You do not need to use the O and the H of KOHES as there is no oxygen present.

Then you need to add the simultaneous equations together and eliminate the electrons

Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O

6Fe2+ 6Fe3+ + 6e

Total equation:

Cr2O72-

(aq) + 6Fe2+(aq) + 14H+

(aq) 2Cr3+(aq) + 6Fe3+

(aq) + 7H2O(l)

Using the reactivity series to predict redox reactions.If a metal is more reactive than another it will displace (push out) that metal from a solution of its salt, e.g. if zinc is placed in a copper sulfate solution it will displace the copper from its solution and replace it.

Simply:

Zinc + copper sulfate copper + zinc sulfate

Zn(s) + CuSO4(aq) Cu(s) + ZnSO4(aq)

Hence there are two half reactions occurring.

Zinc zinc sulfate Copper sulfate copper

Simple ionic equations can represent this:

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

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

Zinc is oxidised, but at the same time it is also the REDUCING AGENT or the REDUCTANT

Copper sulfate is reduced to copper, however it then becomes the OXIDISING AGENT or the OXIDANT.

Then add the two half equations together. The electrons cancel each other out and in every case you must ensure that this happens.

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

(aq) + Cu(s)

You can always refer to the reactivity series of metals to determine whether or not a displacement reaction will occur.

This is a spontaneous reaction and can produce energy. The energy generated by such reactions can be harnessed and this is the basis for the working of batteries. One needs to make use of the moving electrons to produce electricity. Chemical energy can be harnessed to produce electricity.

Figure 2 - Reactivity of Metals. Image taken from Chemical Connections.

Figure 3 – Standard Electrode Oxidation and Reduction Potentials Values – A Colley Original

The simplest electric cell – Lemon PowerA very basic cell can be produced with two different metals in a lemon or any citrus fruit. The lemon becomes the electrolyte and this simple cell can be set up as illustrated.

Figure 4 - Lemon Power. Image take from http://www.kirksville.k12.mo.us/khs/teacher_web/alternative/electricity.html

The components then of this simple cell are:

Two metals of different reactivity A connecting wire An electrolyte

A small electric current will be produced and oxidation will occur at the zinc electrode and reduction will occur at the copper electrode.

Unfortunately this simple cell cannot be used to provide light on your bicycle in the dark.

And hence the simple electric cell can be modified to produce a much more sustainable galvanic cell.

The galvanic cellThe components of the galvanic cell are:

Two electrodes Two electrolytes A salt bridge A conducting wire

Most half cells contain a metal strip (the electrode) dipped into a solution of its ions. The wire connects the two electrodes provides a pathway for the electrons, which are able to move only along the electrodes and the wire. They cannot move in the solution.

The electrodes and the wire form the EXTERNAL CIRCUIT.

The two half cells are also connected by the SALT BRIDGE which contains another electrolyte. The purpose of the salt bridge is to provide ions for those lost or gained in the half cell. The movement of these ions within the electrolyte is called the INTERNAL CIRCUIT.

The electrochemical series is strictly valid only for the conditions under which it was determined:

A temperature of 25oC Pressure of 1 atmosphere 1M concentrations of solutions

In a galvanic cell, the stronger reductant is in the half cell with the negative electrode (anode).The stronger oxidant is in the half cell with the positive electrode (cathode).

Another way to predict the electrode reactions is to remember that the higher of the two half cell reactions in the electrochemical series goes forward and the lower one is reversed.The higher half reaction occurs at the positive electrode and the lower reaction occurs at the negative one.

In order to predict the voltage available from this cell, the formula that is used us as follows Eo = Eo - ER

Where Eo is the oxidant and ER is the reductant.The answer will then be in VOLTS.

Constructing a Galvanic cell

Figure 5 – Constructing a Galvanic cell. Image taken from Chemical Dimensions

An OIL RIG cat

This is just an amazing acronym which will help you in your pursuit of understanding.

There are two electrodes:

The anode (an) Oxidation occurs at the anode. Electrons are lost at the anode. Oxidation is losing (OIL) electrons.

Reduction is losing (RIG) electrons. Reduction occurs at the cathode (cat).

Electronic conductorsElectrodes are electronic conductors – they conduct current by the flow of electrons. They may be made of a metal or graphite.

Where a half cell contains a metal reductant, the metal serves as the electrode, for example a zinc electrode in the Zn2+/ Zn half cell. Where the reductant is not a metal the electrode is either graphite or a non reactive metal such as platinum, for example a graphite electrode in a Fe3+/ Fe2+ half cell or a platinum electrode in a Cl2 /Cl- half cell.

Figure 6 – Half-cells. Image taken from Chemical Dimensions

Further Examples

Can you determine the Eo values for the examples below?