1 Chemistry Tutorial REDOX REACTIONS by Dr John G Wright Press PgDn for the next slide The Wright Stuff

2 Redox Reactions F This is the first of a set of tutorials on Redox Reactions. These are reactions where one reagent is oxidised and the other is reduced. Related topics are also examined. F This section covers oxidation and reduction.

3 Redox Reactions F Oxidation and Reduction(this tutorial, 22 slides) F Oxidation Numbers F Disproportionation F Balancing Half Equations F Electrochemical Cells F Problems The individual topics can be viewed as separate slide shows. If a slide show is missing from the web page, it means it is being updated and improved. (Or perhaps I just haven’t finished it yet!) The tutorials are divided into five main sections, each covering a separate topic concerning redox reactions, plus a problems section.

4 Oxidation and Reduction F This section looks at the meaning of the terms oxidation and reduction, both the old oxygen/hydrogen based definition, and the modern electron based definition. Examples illustrating each definition are given, and explained. F Students on introductory level courses may only need to read the section on the older definition, but those on advanced courses must study both sections, with the emphasis being placed on the modern electron-based definition.

5 Oxidation and Reduction F Oxidation - the gain of oxygen or the loss of hydrogen by an element or compound. F Reduction - the loss of oxygen or the gain of hydrogen by an element or compound. Although these definitions have now been replaced by more useful modern definitions, they are still used in introductory level courses, and are much easier to use when working with organic compounds. So we will start by looking at a few examples of the older definitions in action before moving on to the more important modern definitions which you must learn to use. In the past oxidation and reduction were defined in rather simple terms.

6 Oxidation Oxidation and Reduction  4Fe + 3O 2  2Fe 2 O 3 Here the iron has gained oxygen to form Fe 2 O 3, so the iron has been oxidised to form iron(III) oxide.  2CO + O 2  2CO 2 Here the carbon in carbon monoxide has gained oxygen, so the carbon has been oxidised to carbon dioxide.  2NO 2 + H 2 O 2  2HNO 3 Here the nitrogen in NO 2 has gained oxygen from the hydrogen peroxide (H 2 O 2 ), therefore the nitrogen has been oxidised. Notice that we don’t need to use O 2 as the source of oxygen in our reactions. The following examples show the older definition of oxidation being applied to reactions. First, gain of oxygen.

7 Oxidation and Reduction  CH 3 CH 2 CH 3 + 5O 2  3CO 2 + 4H 2 O  CH 3 CH 2 OH + 3O 2  2CO 2 + 3H 2 O  2H 2 S + 3O 2  2SO 2 + 2H 2 O  SiH 4 + 2O 2  SiO 2 + 2H 2 O In each of the above cases combustion occurs, elements on the left hand side of the equation gain oxygen to form the products, and therefore oxidation occurs. The combustion of anything is another example of oxidation where the old definition of gain of oxygen is easy to see. where the old definition of gain of oxygen is easy to see.

8 Oxidation and Reduction  CH 3 CH 2 OH + KMnO 4  CH 3 CHO + MnO 2 (This reaction occurs when excess alcohol is present.)  CH 3 CH 2 CH 2 OH + KMnO 4  CH 3 CH 2 CO 2 H + MnO 2 (This occurs when the KMnO 4 is present in excess.)  CH 3 CHOHCH 3 + KMnO 4  CH 3 COCH 3 + MnO 2 (Here the alcohol is present in excess.) In each case we can clearly see that the alcohol has lost some hydrogen when it reacts, therefore oxidation is occurring. (Note that in the second example gain of oxygen has also occurred.) Now let’s look at oxidation as loss of hydrogen. (All these equations are unbalanced for simplicity, and the hydrogen is lost as water.)

9 Oxidation and Reduction  You should be familiar with this reaction, which is the overall reaction occurring in the blast furnace. 2Fe 2 O 3 + 3C  4Fe + 3CO 2 F Clearly the Fe 2 O 3 has lost oxygen. Obviously this is reduction. In fact the name reduction comes from the early days of metallurgy, when chemists talked about “reducing a metal ore to the metal”. I.e. they reduced it from a complicated compound to the simple element.  CuS + O 2  Cu + SO 2 Here too a metal ore is reduced to the metal. It’s time to examine a few cases where reduction occurs. I.e. reactions where oxygen is lost or hydrogen is gained.

10 Oxidation and Reduction  CH 3 CH=CH 2 + H 2  CH 3 CH 2 CH 3  2CH 3 CH 2 CHO + H 2  2CH 3 CH 2 CH 2 OH  CH 3 CH 2 COCH 3 + H 2  CH 3 CH 2 CHOHCH 3 Not exactly difficult to spot reduction, is it? But what about the modern definition, the more important one? And why did chemists have to change such a relatively easy to understand idea as the older one? Finally, examples of reduction where hydrogen is gained. Many organic reactions are good examples of this type of reaction.

11 Oxidation and Reduction The older definitions are a little bit limited, in that only reactions which involve oxygen or hydrogen at some point can be considered to be oxidation or reduction. In many cases we can find groups of reactions where an element undergoes the same basic change in each of the reactions, but using the older definitions only one of them is classified as oxidation or reduction. As you should know, the majority of chemistry depends on the electronic configuration of atoms. So let’s look at what the electrons are up to in these reactions.

12 Oxidation and Reduction  2Mg + O 2  2MgO  Mg + S  MgS  Mg + Br 2  MgBr 2 In each case the magnesium atom is losing two electrons to form an ion with a charge of two plus. Mg  Mg e - However, only the first reaction would be classified as oxidation. From the magnesium’s point of view, this is silly. It’s simply losing two electrons in every case. So the definition of oxidation and reduction needs to be improved. Examine the three reactions below. In each of them the metal is undergoing an identical change if we look at the electrons, but by the older definition only one of them is called oxidation.

13 Oxidation and Reduction F Oxidation - the loss of electrons by an atom. F Reduction - the gain of electrons by an atom. Now all three of the previous reactions can be called oxidation. And many more reactions which do not involve oxygen or hydrogen can now be easily classified. These are the definitions that you are expected to learn and use from now on, especially if you are on an advanced level course. (However note that for organic reactions the older definitions are still very useful, as it is sometimes hard to see which atoms have gained or lost electrons in many organic reactions.) So what are these more modern, more useful definitions?

14 Oxidation and Reduction F Oxidising agents cause oxidation to occur. An oxidising agent is the substance which accepts electrons in an oxidation- reduction reaction. (An oxidising agent is easily reduced.) F Reducing agents cause reduction to occur. A reducing agent is the substance which donates electrons in an oxidation-reduction reaction. (A reducing agent is easily oxidised.) There are a couple of other definitions that you have to learn when working with these definitions of oxidation and reduction.

15 Oxidation and Reduction  Here’s one from an earlier slide, plus a new one. 4Fe + 3O 2  Fe 2 O 3 2Fe + 3Cl 2  2FeCl 3  In both cases, the iron loses electrons Fe  Fe e - And iron also loses an electron in the next example.  2FeCl 2 + Cl 2  2FeCl 3 Fe 2+  Fe 3+ + e - Who said we had to start from an element? We can oxidise ions as well. F And who said the ion had to be positive? Have a careful look at the next example. Some examples of oxidation based on the modern definition.

16 Oxidation and Reduction  For example these half equations, where [O] represents an oxidising agent, show exactly this. 2I - + [O]  I 2 + 2e - I - + [O]  IO e - Notice that in the second example, although we start and finish with a negative ion, because electrons are lost in the process, it is oxidation. (A full explanation of half equations and oxidation numbers is given in later sections which will make these reactions easier to understand.) Some reagents can oxidise halide ions to the halide or to an oxo acid anion.

17 Oxidation and Reduction  4Fe + 3O 2  Fe 2 O 3 2Fe + 3Cl 2  2FeCl 3  We already know that the iron is oxidised, losing electrons. Fe  Fe e -  In the first reaction, the oxygen gains electrons, (this happens in many of the earlier reactions), in the second one the chlorine gains electrons, so both are being reduced. O 2 + 4e -  2O 2- Cl 2 + 2e -  2Cl - F Note that this means in all our examples both oxidation and reduction are occurring in the same reaction. Where do these electrons go? The other reagent uses them, it gains electrons, and so by our definition, it is reduced. Let’s examine some of our earlier reactions again.

18 Oxidation and Reduction  Mg + S  MgS Here the magnesium loses electrons, as we stated earlier. And the sulfur gains the electrons, and so the sulfur is reduced in this redox reaction. S + 2e -  S 2-  Mg + Br 2  MgBr 2 Again the Mg loses electrons and is oxidised, and this time it’s the bromine which gains electrons and gets reduced in this redox reaction. Br 2 + 2e -  2Br - Let’s look at a few more of our earlier examples of oxidation.

19 Oxidation and Reduction F The older definition was restrictive in that it only looked at one of the reagents to classify the reaction as either oxidation or reduction, but with the modern definition we can see that both reactions must occur at the same time. So chemists invented the term REDOX to describe these reactions where both oxidation and reduction are occurring. F It’s important to remember that as something loses electrons, something else must gain these electrons. We don’t have reactions where we end up with a bottle of electrons as one of the products! (Nor as the starting material.) We oxidise one reagent and reduce the other. We can’t have one without the other.

20 Oxidation and Reduction F Disproportionation is a special type of redox reaction, where the same element in the same compound is oxidised and reduced in the same reaction. F This means it forms two products, one of which is the product of the oxidation, and the other is the product of the reduction. To understand this particular type of redox reaction better, we must understand what is meant by the term oxidation number, sometimes called oxidation state. Therefore the next tutorial examines oxidation numbers, and ends with some examples of disproportionation.

21 The End F I hope you have enjoyed this fist tutorial. It was NOT produced using extremely expensive authoring software, but relatively inexpensive presentation software commonly supplied with office software. F I also hope you have increased your understanding of chemistry, learned something useful about chemistry and that it will increase your marks in examinations. F Now you are ready for the next section - Oxidation Numbers Bye for now, Dr John G Wright, The Wright Stuff

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