lesson 1- introduction to redox reactions

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Lesson 1- Introduction to Redox Reactions

• Define oxidation and reduction in terms of oxygen loss/gain

• Define redox in terms of electron transfer.

• Construct redox half equations

• Assign Oxidation States

Unit 7- Redox Chemistry and Electrochemistry


What is a redox reaction?Oxidation is the addition of oxygen to a substance andreduction is the removal of oxygen from a substance.

lead oxide + carbon lead

carbon monoxide +

oxygen removedreduction

oxygen addedoxidation

Reduction and oxidation always take place together. Why is this type of reaction called a redox reaction?

redox = reduction and oxidation

Which substances are oxidized and reduced in this reaction?


Redox reactants – oxidized or reduced?


Redox and electronsMagnesium burns in oxygen to form magnesium oxide.

A redox reaction can also be explained in terms of the gain or loss of electrons.

magnesium + oxygen magnesium oxide

2Mg(s) O2(g) 2MgO(s)+

What happens to the atoms and electrons in this reaction?

It is obvious that the magnesium has been oxidized, but what has happened to the oxygen?


Oxidation and electron lossWhen magnesium burns in oxygen to form magnesium oxide, what happens to magnesium and its electrons?

Oxidation is the loss of electrons.

oxidized(electrons lost)

Mg Mg2+ O2-O+

The magnesium has been oxidized.

The Mg atom has lost 2 electrons to form a Mg2+ ion.


Reduction and electron gainWhen magnesium burns in oxygen to form magnesium oxide, what happens to oxygen and its electrons?

Reduction is the gain of electrons.

reduced(electrons gained)

Mg Mg2+ O2-O+

The oxygen has been reduced.

The O atom has gained 2 electrons to form a O2- ion.


Redox and OILRIGAn easy way to remember what happens to the electrons during oxidation and reduction is to think… OILRIG!


Using OILRIGWhat does OILRIG stand for in terms of redox reactions?

Oxidation

Is

Loss of electrons

Reduction

Is

Gain of electrons


What is a half-equation?

Redox reactions involve the transfer of electrons.

oxidation: Mg Mg2+ + 2e-

magnesium + oxygen magnesium oxide

2Mg (s) O2 (g) 2MgO (s)+

reduction: O2 + 4e- 2O2-

Equations written to show what happens to the electrons during oxidation and reduction are called half-equations.

What are the half-equations for the oxidation and reduction processes in this reaction?


What does each half-equation show?


Redox reactions – summary


1. Explain why the reaction between calcium and oxygen to form calcium oxide is an example of a redox reaction.

2. Define oxidation and reduction. In terms of oxygen and electrons.

3. Balance the following half equations

Br2(l) + _e¯ → 2Br¯(aq)

Ag+(aq) + _e¯ → Ag(s)

Fe3+(aq) + _e¯ → Fe2+(aq)

I2(s) + _e¯ → 2I¯(aq)

Cu +(aq) + _e¯ → Cu(s)

Cu2+(aq) + _e¯ → Cu(s)

Questions


Oxidation numbers• We can use oxidation numbers to decide whether a

something has been oxidised or reduced.

• Oxidation numbers show how many electrons are gained or lost by an element when atoms turn into ions.

• The oxidation numbers of the elements are zero. Oxidation numbers also distinguish between the compounds of elements such as iron which can exist in more than one oxidation state. In iron (II) chloride, the Roman number II shows that the iron is in oxidation state +2.


Oxidation State/NumberRules:

• The oxidation number of an element is zero.

• The oxidation number of an ion is equal to the charge of the ion.

• Hydrogen has an oxidation number of +1 (this doesn't apply to hydrides when hydrogen is -1).

• Oxygen has an oxidation number of -2 (except in peroxides when it is -1).

• Fluorine always has an oxidation number of -1

• It is then simply a matter of adding together all of the oxidation numbers of the elements in a compound and making sure that the total is = 0 or equal to the charge on a compound ion.


Assigning Oxidation State- Examples

Cr2O3

KMnO4

HNO3


Assign oxidation numbers to each element in each of the following compounds:

1. KMnO4 2. NiO2

3. P4O6 4. Fe3O4 5. SF4 6. XeOF4 7. CO 8. Na2C2O4 9. As2O3 10. NaBiO3 11. Mg2P2O7

12. Hg2Cl2


Oxidation Numbers- Charged Species

Cr2O72-

KMnO4 -


Identifying Redox ReactionsExample 1:

Balanced Equations:

Zn + CuCl2 → ZnCl2 + Cu balanced equation

Zn + Cu2+ → Zn2+ + Cu ionic equation

Note: chloride is a spectator ion

Half Equations:

Zn → Zn2+ +2 e- oxidation

Cu2+ + 2e- → Cu reduction


Example 2: Identify redox processes

AgNO3 + NaCl → AgCl + NaNO3


Use oxidation numbers to identify whether the following processes, represented by equations,

involve oxidation (O), reduction (R), both or neither. Give

reasons for your decisions.

(a) F2 +2e → 2F-

(b) MnO2 + 2H2O → MnO4- + 4H+ + 3e

(c) H2C2O4 → 2CO2 +2H+ + 2e

(d) 2CrO42- + 2H+ → Cr2O7

2- + H2O

(e) C2H4 + Cl2 → C2H4Cl2

(f) Zn → Zn2+ + 2e

(g) Fe2+ + 2e → Fe


Example 3:

N2 + 3H2 → 2NH3

Oxidation is a increase in Oxidation Number

Reduction is a decrease in Oxidation Number


Task

Complete Worksheet 1:


Lesson 2- Electrolysis of molten ionic compounds

• Properties of ionic compounds

• Define the term electrolysis

• Describe the electrode products in the electrolysis of

-Molten lead(II) bromide



Starter• Say whether the following three reactions

are redox reactions or not, by showing the oxidation states of each of the elements involved:

a) Zn + 2HCl → ZnCl2 + H2

b) CuO + 2HCl → CuCl2 + H2O

c) MnO2 + 4 HCl → MnCl2 + Cl2 + 2 H2O

d) Cl2 + H2O → HCl + HOCl


What are ionic compounds?Ionic compounds are made up of positive metal ions and negative non-metal ions. What ions are in sodium chloride?

The positive and negative ions in an ionic compound attract each other strongly. It takes a lot of energy to separate them.

positive sodium ions

negative chloride ions

How does structure affect the properties of ionic compounds?


Do dissolved ionic compounds conduct?


Properties of ionic compounds


Does it conduct electricity?


Electrolysis

A process in which a chemical reaction is caused by the passage of an electric current.

A molten ionic compound can be split in to its elements. eg PbCl2 into Lead metal and Chlorine gas.


Electrode- The point at where the electric current enters and leaves a current a battery.

Electrolyte- an ionic compound which will conduct electricity when it is molten or dissolved in water; electrolytes will not conduct electricity when solid.


What is electrolysis?

An ionic compound conducts electricity when it is molten or in solution. The current causes the ionic compound to split up and form new substances.

Electrolysis has many uses, including:

purifying copper

plating metals with silver and gold

extracting reactive metals, such as aluminium

making chlorine, hydrogen and sodium hydroxide.

This process is called electrolysis, and means “splitting by electricity”.


heat

What happens during electrolysis?

In electrolysis, the substance that the current passes through and splits up is called the electrolyte.

Positive ions move to the negative electrode

(Cathode) and gain electrons.

This is reduction.

Negative ions move to the positive electrode

(Anode) and lose electrons.

This is oxidation.

The electrolyte contains positive and negative ions.

What happens to these ions during electrolysis?


Electrolysis of molten lead bromide


What redox processes occur at the electrodes during the electrolysis of molten lead bromide (PbBr2)?

Electrolysis of molten PbBr2 – redox equations

What is the overall equation for the electrolysis of molten lead bromide ?

At the negative electrode:

Pb2+ + 2e- Pb (reduction)

lead bromide lead + bromine

PbBr2 (l) Pb (l) + Br2 (g)

At the positive electrode:

2Br- Br2 + 2e- (oxidation)


Electrolysis of molten PbBr2 – summary


Electrolysis problems 11) Define electrolysis. (1)

2) a) Explain monatomic and simple molecular substances do not conduct electricity. (2)

b) Explain why metals conduct electricity. (3)

c) Explain why ionic substances do not conduct electricity as solids but do when molten or dissolved.

(3)

3) When each of the following ionic substances is melted and electrolysed, layout your answers clearly:

• predict what would be produced at the positive electrode• give the half equation for the reaction at the positive electrode• predict what would be produced at the negative electrode• give the half equation for the reaction at the negative electrode

a) molten lead bromide


Electrolysis problems 11) Define electrolysis.

A) The use of an electrical current to split up the components of a molten ionic compound or the components within a solution (1)

2) a) Explain monatomic and simple molecular substances do not conduct electricity.

A) Molecular substances do not consist of ions or electrons that are able to move freely and conduct electricity. (2)

b) Explain why metals conduct electricity. (3)

A) Metals consist of free electrons which are able to move and transport charge.

c) Explain why ionic substances do not conduct electricity as solids but do when molten or dissolved.

A) The ions within a solid ionic substance are unable to move and transport electrical current, when molten or dissolved the ions are able to move and transport electrical charge


3) When each of the following ionic substances is melted and electrolysed, layout your answers clearly:

• predict what would be produced at the positive electrode• give the half equation for the reaction at the positive electrode• predict what would be produced at the negative electrode• give the half equation for the reaction at the negative electrode

a) molten lead bromide

At the negative electrode: Product: LeadHalf Equation: Pb2+ + 2e- Pb (reduction)

At the positive electrode: Product: Bromine GasHalf Equation: 2Br- Br2 + 2e- (oxidation)


Lesson 3- Electrolysis of Ions in Solution

• Define the term electrolysis

• Describe the electrode products in the electrolysis of

-Molten lead(II) bromide

-Concentrated hydrochloric acid

-Concentrated aqueous sodium chloride.

-Between inert electrodes (platinum or carbon).



Starter

Draw a rough sketch of the equipment used during electrolysis.

Include:

• Electrolyte

• Electrodes

• Anode

• Cathode


The electrolysis of HCl

Task

1.Can you construct balanced half equations for this process.

2.Combine the half equation to give the full equation of what is happening.

3.What are the products of this reaction.


At the negative electrode: Product: HydrogenHalf Equation: 2H+ + 2e- H2(reduction)

At the positive electrode: Product: ChlorineHalf Equation: 2Cl- -> Cl2 + 2e- (Oxidation)

Overall: 2HCl -> H2 + Cl2


Questions•Electrolysis: Splitting up with electricity.

1.This requires a liquid, or ________, which will conduct electricity.

2.Electrolytes are usually free _____ dissolved in water.

3.Electrolytes can also be ________ ionic substances.

4.The _________ is an electron pump, taking electrons away from the positive ______ to the negative _______. 5. Ions gain or lose electrons at the ______ and _____ atoms or molecules are formed.

6. The are ______ often made of _____ materials such as platinum and carbon

electrolyte

ions molten

electricity

anode

cathode

electrodes

neutral

electrodesinert


What would you expect the products of this reaction would be?

The electrolysis of

Concentrated aqueous sodium chloride.


Electrolysis of NaCl solution


Products of electrolysis of NaCl solution

The electrolysis of sodium chloride solution produces three very useful products:

Chlorine used for killing bacteria in water, for bleach and making plastics like PVC.

Hydrogen used for making margarine and fertilizers, and for rocket fuel.

Sodium hydroxide used in many chemical reactions, such as making soap, neutralizing acids and making paper.

Chlorine is expected as a product of this process but hydrogen and sodium hydroxide are surprising products.

What happens at the electrodes to form these products?


How does the chlorine form?In the electrolysis of NaCl solution, the negative chloride ions (Cl-) are attracted to the positive electrode.

Are the Cl- ions oxidized or reduced?

How many electrons are lost by each Cl- ion?

How many Cl- ions join to make a Cl2 molecule?

What is the half-equation for this redox process?2Cl- Cl2 + 2e- (oxidation)

oxidized

one

two

Here, the Cl- ions lose electrons to make chlorine atoms, which then form chlorine molecules (Cl2).


Why is sodium not formed? In the electrolysis of sodium chloride solution, the Na+ ions might be expected to form sodium at the negative electrode.

For all ionic compounds containing a metal that is more reactive than hydrogen, electrolysis of a solution of the compound will produce hydrogen rather than the metal.

At the negative electrode, the H+ ions compete with the Na+ ions. The H+ ions gain electrons; the Na+ ions stay in solution.

Instead, hydrogen gas is produced here.

This is because sodium chloride solution also contains H+ ions from some of the water: H2O (l) H+ (aq) + OH- (aq).


In the electrolysis of NaCl solution, the positive hydrogen ions

(H+) are attracted to the negative electrode.

Here, the H+ ions gain electrons to make

hydrogen atoms, which then form hydrogen molecules (H2).

How does the hydrogen form?

Are the H+ ions oxidized or reduced?

How many electrons are gained by each H+ ion?

How many H+ ions join to make a H2 molecule?

What is the half-equation for this redox process?2H+ + 2e- H2

(reduction)

reduced

one

two


How does the sodium hydroxide form?Sodium chloride solution has four types of ions:

What is the overall equation for the electrolysis of a sodium chloride solution?

The Cl- ions form chlorine at the positive electrode and the H+ ions form hydrogen at the negative electrode. So, what’s left?

Na+ and OH- ions are left behind and so a solution of sodium hydroxide (NaOH) is formed.

2NaCl (aq) + 2H2O (l) H2 (g) + Cl2 (g) + 2NaOH (aq)

Na+ and Cl- ions from the sodium chloride

H+ and OH- ions from the water.


Electrolysis of NaCl solution


Hydrogen or metal?

Complete the table for these compounds.

incr

easi

ng

rea

ctiv

ity

potassiumsodiumcalcium

magnesiumaluminium

zinciron

copper

gold

lead

silver

(carbon)

(hydrogen)

platinum

If an ionic compound contains a metal that is more reactive than hydrogen, electrolysis of a solution of the compound will produce hydrogen, not the metal.

hydrogen

copper

hydrogen

silver

hydrogen

potassium chloride

copper sulphate

sodium bromide

silver nitrate

zinc chloride

Ionic compound Product at the negative electrode


Lesson 4- Electroplating

• Describe the electroplating of metals• Name uses of electoplating



Electoplating

• The metal plating process (also called electroplating)uses electrolysis of a solution containing ions of the plating metal.

• The anode is made from the pure plating metal.The metal object which needs plating is used as the cathode.

• Most metals can be plated.Common plating metals are gold, nickel and silveras well as chromium and zinc referred to above.


Silver plating could be done in the cell below.

When electricity is passed through the cell silver is dissolved at the anode by oxidation.Ag+ ions go into the silver nitrate solution.

Silver is deposited onto the surface of the object by reduction at the cathode.

As silver ions move from the anode to the cathodethe anode gets smaller as the object becomes silver plated. This is a redox reaction.

Ag(s) →Ag+(aq) + e-

Ag+(aq) + e-→ Ag(s)


Electroplating

• You can use electrolysis to coat one metal with another. This is called ____________. Electroplating is used a great deal in industry, for example; chrome-plating car bumpers.

• If you wanted to coat a nickel vase with silver, you would set the vase as the _______ and the silver as the _________

• At the _______ : Silver _______ forming silver ions. • • At the ________: Silver ions receive ________ and form a layer of

silver on the vase.


Electroplating

You can use electrolysis to coat one metal with another. This is called electroplating. Electroplating is used a great deal in industry, for example; chrome-plating car bumpers.

If you wanted to coat a nickel vase with silver, you would set the vase as the cathode and the silver as the anode.

At the anode: Silver dissolves forming silver ions. At the cathode: Silver ions receive electrons and form a

layer of silver on the vase.


2)A steel fork was

placed in a solution containing silver ions, and connected in a circuit as shown. After a few minutes the fork was coated with a layer of silver.

1) Balance the following half equations.a) Br– e– Br2 c) Al3+ e– Al

b) K+ e– K d) O2– e– O2

solution containingsilver ions, Ag+(aq)

steel forkpure silver

a) Was the fork the positive or negative electrode?b) Write an equation for the reaction that takes place at the positive electrode.c) Write an equation for the reaction that takes place at the negative electrode.d) Steel is quite often plated by other metals in this way. Why is steel often

electroplated?


Lesson 5- The extraction of Aluminium

• Describe, in outline, the manufacture of aluminium from pure aluminium oxide in molten cryolite,



Why and how is aluminium extracted?

Aluminium is one of the most useful metals in the world.

Aluminium ore (bauxite) has a very high melting point (2050 °C).

Electrolysis is used to extract aluminium from its ore. Why is it not possible to extract aluminium by heating its ore with carbon?

For electrolysis, the ore is dissolved in a compound called cryolite (Na3AlF6), which lowers the melting point to 700 °C. Why is this important economically?


potassiumsodiumcalcium

magnesiumaluminium

zinciron

copper

gold

lead

silver

(carbon)

(hydrogen)

platinum


Extracting aluminium


What redox processes occur at the electrodes during the electrolysis of aluminium oxide (Al2O3)?

Extracting aluminium – redox equations

What is the overall equation for the extraction of aluminium by electrolysis?


Al3+ + 3e- Al (reduction)

aluminium oxide aluminium + oxygen

2 Al2O3 (l) 4 Al (l) + 3 O2 (g)


2O2- O2 + 4e- (oxidation)


Extracting aluminium – summary


Purifying copper using electrolysis


Lesson 6: Purification of Copper

-Describe the principle of electrolysis.

-Understand the method used to refine copper and the reasoning behind it.

- Name uses of copper

- Be familiar with and able to use in sentences the vocabulary for this topic


Potassium permanganate

crystals

Filter paper soaked in salt

solution


What’s going to

happen in this

experiment?

The purple potassium

permanganate is going to spread over the paper because of diffusion.

The purple colour is going to

move to the negative side

only.

The purple colour is going to

move to the positive side faster than

the negative side.

I think it’s an example

of electrolysis.


Potassium permanganate

crystals

Filter paper soaked in salt

solution


Notes

• Sketch a copy of the diagram.

• Describe the movement of colour.

• Explain the movement of colour.


Uses of electrolysis

What uses of electrolysis can you think of?


Copper is not very reactive and can occur native but it is rare to find pure copper. Usually, it is found combined with other elements, such as in the ore malachite.

The copper extracted from compounds by reduction with carbon is impure (blistered copper). Electrolysis can actually be used at this stage to remove the impurities and obtain pure copper.

How is copper purified?


Uses of Copper• Copper is an excellent conductor

and does not corrode quickly. These properties make it a good material for wiring and plumbing.

• Only pure copper can be used for electric wires. Even a very low level of impurities will reduce copper’s conductivity.


Purifying copper using electrolysis


Complete the worksheet !!

Read the questions carefully...


Purifying copper – redox equations

What happens at the electrodes during the purification of copper by electrolysis ?


Cu2+ + 2e- Cu (reduction)

This process is carried out on a huge scale in industry and the copper formed on the negative electrodes is 99.99% pure.


Cu Cu2+ + 2e- (oxidation)

The precious metals recovered from the impurities are also sold off and help to make this industrial process profitable.


Labelling electrolysis


Plenary

• Taboo


Lesson 7: Electrochemical Cells

• Describe the production of electrical energy from simple cells, i.e. two electrodes in an electrolyte.

• State the use of batteries as a convenient, portable energy source


Lesson 8: Past Paper questions

•Complete past paper questions


Lesson 9: Poster Presentations

•Produced a poster on the given subject in as much detail as possible.

•A good poster will be clear, well structure, NEAT, eye catching and contain accurate scientific information

•You will be expected to give a short 3 minute presentation explaining the scientific concepts in the poster.

•You may be asked questions about this


PAL 1 Groups

Group Title Members

1 Electrolysis of Sodium Chloride and its products Alex, Steven, Gladys, Shirley

2 Extraction of Aluminium Jerry, Alexander, Frank

3 General Electrolysis and Electroplating Greg, Kitty, Alice, Alen

4 General Electrolysis and Purification of Copper Alex. H, Fiona, Vivian

5 Extraction of Iron in the blast furnace Leondro, Zoen, Kibi

6 Extraction of Zinc Moey, Arthur, Lesley

7 Electrolysis and its uses Jane, Helen ,Ennl


PAL 2 Groups

Group Title Members

1 Electrolysis of Sodium Chloride and its products Cenery, Alice, Emily,

2 Extraction of Aluminium Joyce, Grace, Peter

3 General Electolysis and Electroplating Stella, Hacken, Daisy

4 General Electrolysis and Purification of Copper Lee, Maggie, Cindy

5 Extraction of Iron in the blast furnace Frankic, Jeffrey, Johnny

6 Extraction of Zinc from Zinc Gilbert, Kizi ,Angela

7 Electrolysis and its uses John, Cheney,Simon


Purifying copper – redox equationsWhat happens at the electrodes during the purification of copper by electrolysis ?


Cu2+ + 2e- Cu (reduction)

This process is carried out on a huge scale in industry and the copper formed on the negative electrodes is 99.99% pure.


Cu Cu2+ + 2e- (oxidation)

The precious metals recovered from the impurities are also sold off and help to make this industrial process profitable.


Purifying copper – true or false?


Electrolysis of dilute sulfuric acidElectrolysis can be used to split water (H2O)

into its elements, hydrogen and oxygen.

Which product will form at each electrode?

The conductivity of water can be improved by adding dilute sulfuric acid. This releases more ions so that more current flows during electrolysis, which creates hydrogen and oxygen.

This is how hydrogen for fuel cells can be made and how oxygen can be produced from water on spacecraft.

Water is a covalent compound and so is a poor conductor of electricity. However, it does contain a few free H+ and OH- ions:

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


Electrolysis of dilute sulfuric acid


What happens at the electrodes during the electrolysis of dilute sulfuric acid?

Electrolysis of dilute H2SO4 – redox equations

What is the overall equation for the electrolysis of dilute sulfuric acid?


2H+ + 2e- H2 (reduction)

2H2O (l) 2H2 (g) + O2 (g)

Twice as much hydrogen forms as oxygen. Why is this?


4OH- 2H2O + O2 + 4e- (oxidation)

In water, there are 2 hydrogen atoms for every oxygen atom, so the ratio by volume, of H2 to O2, is 2:1.


What are the products of electrolysis?


Glossaryelectrode – A solid conductor of electricity, which is used to make electrical contact with an electrolyte.

electrolysis – The process which uses electricity to split up compounds.

electrolyte – A substance which conducts electricity and can be split up by a current when molten or in solution.

ions – Charged particles formed when atoms lose or gain electrons.

oxidation – A type of reaction involving the gain of oxygen or the loss of electrons.

redox – A type of reaction in which oxidation and reduction take place at the same time.

reduction – A type of reaction involving the loss of oxygen or the gain of electrons.


Anagrams


Oxidized or reduced?


What is involved in electrolysis?

lesson 1- introduction to redox reactions

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