Chapter 16 Redox Reactions p.1/7 16.1 Defining Oxidation and Reduction In terms of addition/removal of oxygen Consider the following reaction 1. Mg changed to MgO because Mg is oxidized by CuO. 2. CuO changed to Cu and is said to be reduced. 3. Oxidation is the addition of oxygen and reduction is the removal of oxygen. 4. Oxidation and reduction must occur together and this type of oxidation-reduction reaction is called REDOX reactions. (2) Another way to look at the reaction in terms of electron-transfer 1. Mg loses 2 electrons and transfers them to Cu 2+. 2. Cu 2+ accepts the 2 electrons and become Cu. 3. Mg undergoes oxidation as it loses electrons. 4. CuO undergoes reduction as Cu 2+ gains electrons. Summary: [Must remember items (1) to (5) by heart] (1) A redox reaction is a reaction involving transfer of electrons. (2) Oxidation is a process in which a substance loses electrons. (3) Reduction is a process in which a substance gains electrons. (4) An oxidizing agent is a substance which oxidizes others by accepting electrons. (5) A reducing agent is a substance which reduces others by donating electrons Q 16.1 Consider the displacement reaction: Fe(s) + Cu 2+ (aq) Fe 2+ (aq) + Cu(s) Explain why this is a redox reaction. Which is being oxidized? Why? Which is the oxidizing agent? Why?
16.2 Oxidation and Reduction in a simple chemical cell p.2/7 Changes at Electrodes Consider the following chemical cell (1) At the Mg electrode: Mg(s) Mg 2+ (aq) + 2e - (2) At the copper electrode: 2H + (aq) + 2e - H 2 When we add up (1) and (2), the overall equation is: Mg(s) + 2H + (aq) Mg 2+ (aq) + H 2 (g) Anode and Cathode An anode refers to the electrode at which oxidation takes place. In this case Mg is the anode. A cathode refers to the electrode at which reduction takes place. In this case Cu is the cathode. redcat means reduction occurs at the cathode. Ionic Half-equations and ionic equations Consider the half equations: (1) Zn(s) Zn 2+ (aq) + 2e - (2) Ag + + e - Ag(s) When (1) and (2) are added up, equation (2) must be multiply by a factor of 2 The overall equation is: Zn(s) + 2Ag + (aq) Zn 2+ (aq) + 2Ag(s) Half equations show only half of the overall reaction, either oxidation or reduction. When 2 half-equations combine together, an overall equation is obtained. Q.16.2 Write half-equations for the following changes. In each case, state whether it is an oxidation or a reduction. Fe 2+ Fe 3+ Cu 2+ Cu OH - O 2 + H 2 O (d) F2 F -
Q 16.3 Given the redox equation: Cl 2 (aq) + 2KI(aq) 2KCl(aq) + I 2 (aq) Rewrite it into an ionic equation. p.3/7 Split the ionic equation into two half-equations, one for oxidation, the other for reduction. 16.3 Simple Chemical Cells with Inert Electrode Setting up a Chemical Cell with Inert Electrode In this cell, two platinum electrodes are immersed in two solutions joined by a salt bridge. (1) At electrode P 2I - (aq) I 2 (aq) + 2e - and a yellow/brown colouration results. (2) At electrode Q Fe 3+ + e - Fe 2+ Adding half equations for (1) and (2) give the overall equation. (Equation 2 has to be multiplied by a factor of 2) 2I - (aq) + 2Fe 3+ (aq) I 2 (aq) + 2Fe 2+ (aq) Chemical Cells and Half Cells A chemical cell is made up of two half cells In a simple chemical cell, the metal couple are dipped into the same electrolyte and electrons flow through the external circuit. In a cell made by dipping two inert electrodes into two electrolytes, a salt bridge is used to connect the two half cells. Electrons flow through an external wire joining the two inert electrodes. Ions flow through the salt bridge and give electrical connection of the two half cells. 16.4 Oxidation Number The oxidation number concept The oxidation number of an element in a compound is the charge of an atom of the element would have if the atom existed as an ion.
p.4/7 Rules for assigning Oxidation Numbers (O.N.) (1) The oxidation number of a free element is zero. (2) The oxidation number of an element in a simple ion is equal to the ionic charge. e.g. in NaCl, the O.N. of Na = +1 and O.N. of Cl = -1. (3) The oxidation numbers of some elements in their compounds are fixed: All Group I elements, e.g. Na +1 All Group II elements, e.g. Mg +2 Hydrogen in most of its compounds +1 (d) Fluorine (the most electronegative element) -1 (e) Chlorine, bromine and iodine in most of their compounds -1 (f) Oxygen in most of its compounds, e.g. MgO, KOH, H 2 O -2 (4) The algebraic sum of oxidation numbers of all elements in a compound is zero. e.g. in MgO, the sum = (+2) + (-2) = 0 The algebraic sum of oxidation numbers of all atoms in a polyatomic ion is equal to the ionic charge. e.g. in OH -, -2 + (+1) = -1, which is the charge of OH -. (5) In a compound or polyatomic ion, the oxidation number of a certain element can be calculated. e.g. in MnO 4 -, the overall charge = -1 and let a be the O.N. of Mn, then 1 = a + 4(-2); a = +7. (6) The oxidation number of an element in one compound may differ from that in other compounds. e.g. in H 2 SO 4, the O.N. of S = +6, in H 2 SO 3, the O.N. of S = +4 Q 16.4 Find the oxidation numbers of the underlined elements in the following substances or ions: S 8 Mg 3 N 2 Na 2 SO 4 (d) Na 2 SO 3 (e) K 2 CrO 4 (f) NH 4 Cl (g) HCO 3 - (h) Cr 2 O 7 2- (i) NF 3 Oxidation Number Charts The oxidation number chart is used to show the different oxidation numbers of an element in different compounds. Refer to Book 1B p.23 Table 16.2 16.5 The Stock System of Naming Compounds The stock system of naming compounds is based on the oxidation number concept. Naming of cations Some metals (usually transition metals) form cations with more than one oxidation number, e.g. Fe 3+ is named as iron(iii) and Fe 2+ is named as iron(ii) Naming of polyatomic anions The names of some common polyatomic ions are given below. (1) SO 4 2- is named as sulphate(vi) (2) SO 3 2- is named as sulphate(iv)
(3) NO 3 - is named as nitrate(v) (4) NO 2 - is named as nitrate(iii) (5) MnO 4 - is named as manganate(vii) (6) Cr 2 O 7 2- is named as dichromate(vi) p.5/7 Q 16.5 Name the following compounds Cu(NO 3 ) 2 Pb(OH) 2 Cr 2 (SO 4 ) 3 (d) KHSO 3 (e) Cu 2 S (f) FeSO 4.7H 2 O (g) (NH 4 ) 2 Cr 2 O 7 (h) BaSO 3 (i) Ca(HCO 3 ) 2 (j) (CH 3 COO) 2 Mg (k) FeCl 2 (l) KClO 2 16.6 Defining Redox Reactions in terms of Oxidation Number Let us look at the redox reaction between magnesium and copper(ii) oxide again. (1) Oxidation is a process in which the oxidation number of an element in a substance increases. (2) Reduction is a process in which the oxidation number of an element in a substance decreases. (3) A redox reaction is a reaction in which the reacting substances undergo changes in oxidation number. Q 16.6 Fill in the following table: Defined in terms of Oxidation Reduction oxygen electron oxidation number Q 16.7 Study the following reactions. CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) Pb(NO 3 ) 2 (aq) + 2NaBr(aq) PbBr 2 (s) + 2NaNO3(aq) 6FeSO 4 (aq) + 3Cl 2 (g) 2Fe 2 (SO 4 ) 3 (aq) + 2FeCl 3 (aq)
In each case, state (i) whether the reaction is redox or not (ii) the formula of the oxidizing agent (if applicable) (iii) the formula of the reducing agent (if applicable) (iv) the name of the element which is oxidized (if applicable) p.6/7 Q 16.8 Comment on the following sentence which consists of two statements joined by the word because. Ammonia is a reducing agent because an ammonia molecule has a lone pair of electrons. 16.7 Strong Oxidizing agents and Reducing agents Refer to Book 1B p.28 & 29 for tables of strong oxidizing and reducing agents. A large number of redox reactions can result from various combinations of oxidizing agents and reducing agents. 16.8 Balancing Redox Equations The following is a list of half equations for three common oxidising agents. (1) 8H + (aq) + MnO 4 - (aq) + 5e - Mn 2+ (aq) + 4H 2 O(l) (2) 14H + (aq) + Cr 2 O 7 2- (aq)+ 6e - 2Cr 3+ (aq) + 7H 2 O(l) (3) 2H + (aq) + NO 3 - (aq) + e - H 2 O(l) + NO 2 (g) The following is a list of half equations for three common reducing agents. (4) Fe 2+ (aq) Fe 3+ (aq) + e - (5) SO 2 (g) + 2H 2 O(l) 4H + (aq) + SO 4 2- (aq) + 2e - (6) 2I - (aq) I 2 (aq) + 2e - Activity Write balanced ionic equations for the following reactions. A solution of acidified potassium manganate(vii) with an aqueous solution of sulphur dioxide.
p.7/7 A solution of acidified potassium dichromate(vi) with a solution of potassium iodide. concentrated nitric acid added to a solution of iron(ii) nitrate. 16.11 Tests for Oxidizing Agents and Reducing Agents Test for Oxidizing Agents (making use of a reducing agent) A piece of wet starch-iodide paper is dipped into the oxidizing agent. In the presence of an oxidizing agent, the following reactions occur: (i) 2I - (aq) I 2 (aq) + 2e - (ii) I 2 (aq) + starch deep-blue complex Test for Reducing agent (making use of an oxidizing agent) A piece of filter paper moistened with acidified potassium permanganate is used to test for reducing agent. In contact with a reducing agent, the following reaction occurs: MnO 4 - (aq) + 8H + (aq) + 5e - Mn 2+ (aq) + 4H 2 O(l) There will be a change of colour from purple to colourless, indicating the presence of a reducing agent. Q 16.13 Which of the following will decolorize acidified potassium permanganate solution? FeSO 4 (aq), ZnSO 4 (aq), SO 2 (g) (Hint: which of the above is a reducing agent?) Which of the following will turn wet starch-iodide paper dark blue? Br 2 (aq), Fe 2 (SO 4 ) 3 (aq), HCl(aq), I 2 (aq) (Hint: which of the above is an oxidizing agent?) Refer to a summary on p.36 & 37 of Book 1B.