Redox Reactions. the loss of electrons by another chemical, oxidation, so both are found together. You must remember

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1 Redox Reactions You are full of redox reactions. The redox reaction respiration releases the energy you need to live, and the food you eat ultimately comes from the redox reaction photosynthesis. Around you, mobile and laptop batteries work using redox reactions, as do fuel cells, bleaches, old fashioned wet photography, and metal corrosion, such as rusting iron. So what are redox reactions? The redox comes from two words: Reduction and Oxidation. In a reaction, oxidation is the loss of electrons from an atom or ion, and reduction is the gain of electrons by an atom or ion. In a redox reaction there is the gain of electrons by one chemical, reduction, and the loss of electrons by another chemical, oxidation, so both are found together. You must remember oxidation is the process of electron loss reduction is the process of electron gain

white solid sodium chloride is made The sodium has

2 Redox Reactions An example of redox When the metal sodium reacts with the green gas chlorine then the white solid sodium chloride is made The sodium has been oxidized and the chlorine has been reduced.

3 Redox Reactions Ancient reduction You may wonder why the word reduction is used to mean a gain of electrons. Thousands of years ago it was noticed that the metal made by smelting had less mass than the original ore. The ore was made smaller, reduced. The loss of oxygen made the mass less. More recently it was realised that the metal atoms were gaining electrons. Oxidizing agents and reducing agents Oxidizing agents oxidize other chemicals, so oxidizing agents are themselves reduced, so oxidizing agents gain electrons. Reducing agents reduce other chemicals, so they lose electrons. You must remember that oxidizing agents are electron acceptors reducing agents are electron donors Why oxidation? Oxidation used to mean just gain of oxygen, but it was realized, more importantly, that the other chemical was losing electrons to the oxygen. So it was decided that oxidation should have the broader meaning loss of electrons.

4 Redox Reactions White hot fire To allow trains to travel at high speeds the rails must be welded together so there are no gaps. This must be done in isolated places so molten iron is made using the Thermit reaction: Here the iron(iii) oxide, Fe 2 O 3, is reduced to iron, Fe, so the iron(iii) oxide is reduced, the iron(iii) oxide is the oxidizing agent The aluminium powder, Al, is oxidized to aluminium oxide, Al 2 O 3, so the aluminium is oxidized the aluminium is the reducing agent

5 Half Equations What are half-equations? Half equations show the gain or loss of electrons by one chemical. For example, this is the full equation for when sodium reacts with chlorine: Each sodium atom is losing an electron to a chlorine atom, so you could write this half equation: Each chlorine atom in a chlorine molecule gains an electron, so you could write this half equation: (You need to show two chloride ions because each chlorine molecule, Cl 2, contains two chlorine atoms.) The way sodium reacts does not depend on the other reactant. For example, if sodium reacted with bromine instead of chlorine the equation would be: Sodium is still gaining electrons in the same way; and the half equation for bromine becomes:

6 Half Equations Half-equations and redox Half-equations involve electron gain or loss, so they always are either oxidation or reduction. In this reaction an electron is lost, so this is oxidation: This time electrons are gained, so this is reduction; Why half? In a reaction, when one chemical loses electrons then another must gain them. Half-equations only show half the story, either the gain or the loss of electrons. Some ions are spectators Ions that take no part in the reaction are called spectator ions. When zinc is put into copper sulphate solution then the full equation is: Each zinc atom is oxidized: Each copper ion is reduced: Notice that the sulphate ions, SO 4 2, do not appear in the equations. They do not change during the reaction. They are dissolved in the water, so are just floating around. As they are said to only watch the reaction they are called spectator ions.

7 Oxidation States What is an oxidation state? An oxidation state is the number of electrons needed to be gained or lost to make a neutral atom. Using oxidation states is a way of working out how oxidized or reduced something is. It is similar to the charge on ions, except that it is also used for covalent compounds. Different oxidation states may have different Different oxidation states may have different colours. In the test tube are all the oxidation states of vanadium from pale yellow +5 to violet +2 at the bottom. Two oxidation states of manganese produced the colours at the top.

8 How to Work out Oxidation States Use these rules to calculate oxidation state: Elements always have an oxidation state of zero. In a compound, the sum of the oxidation numbers equals zero. In an ion, the sum of the oxidation numbers equals the charge. In a compound Group 1 atoms always have a +1 oxidation state, e.g. Na is +1 in NaCl. Group 2 atoms always have a +2 oxidation state, e.g. Mg is +2 in MgCl 2. Group 3 atoms always have a +3 oxidation state, e.g. Al is +3 in AlCl 3. Fluorine always has a 1 oxidation state, e.g. F is 1 in KF. Oxygen has a 2 oxidation state, unless it is in a peroxide compound, such as H 2 O 2, when O is 1, or with fluorine (as F is more electronegative than O); e.g. O is 2 in MgO, but is 1 in Na 2 O 2, and +2 in OF 2 Chlorine has a 1 oxidation state, unless it is with F or O (as they are more electronegative than Cl), e.g. Cl is 1 in NaCl, but +1 in Cl 2 O, and +3 in ClF 3. Hydrogen is +1 except in metal hydrides where it has an oxidation state of 1, e.g. H is +1 in HCl, +1 in H 2 O, but 1 in NaH.

9 Oxidation State Examples Here are some examples of common compounds with all the oxidation numbers: Sodium Chloride (common salt), NaCl, Na = +1, Cl = 1 Sodium Carbonate (washing soda), Na 2 CO 3, Na = +1, C = +4, O = 2 Calcium Fluoride (fluorspar), CaF 2, Ca = +2, F = 1 Calcium Hydroxide (lime water), Ca(OH) 2, Ca = +2, O = 2, H = +1 Potassium Nitrate (saltpetre), KNO 3, K = +1, N = +5, O = 2 Iron(III) Oxide (haematite), Fe 2 O 3, Fe = +3, O = 2 Copper(II) Sulfate, CuSO 4, Cu = +2, S = +6, O = 2

10 Compound Names Old and new names Compounds used to be named differently. At one time each writer would have their own names for compounds. It was very confusing, so internationally chemists agreed standard names. Later it was thought that the words used were difficult or confusing, so internationally it was agreed to use numbers. Here are some examples; KNO 3 was called common nitre or saltpetre, then potassium nitrate, but now is called potassium nitrate(v), because the N has an oxidation state of +5. KNO 2 was called potassium nitrite, but now is called potassium nitrate(iii), because the N has an oxidation state of +3. As further examples here are some chlorine compounds; Some of the old names are still used. For example, KMnO 4 should be called potassium manganate(vii), but is was known as potassium permanganate.

How To Work Out Oxidation States You need to be able to calculate oxidation states in various situations. These worked examples will help you when you meet more difficult questions.

11 How To Work Out Oxidation States You need to be able to calculate oxidation states in various situations. These worked examples will help you when you meet more difficult questions. Step 1 Write down the formula. Step 2 For the oxidation states known, write the oxidation states above the symbol. Remember an oxidation state is for one atom. Step 3 For the oxidation states known, write the sum of the oxidation states below the symbol. Step 4 Work out the oxidation state of the unknown. For a compound, the sum of the oxidation states must equal zero. For an ion, the sum of the oxidation states must equal the charge. If there is more than one atom of the element, then its number is the sum of the oxidation states. These are the colours of the oxidation states of the radioactive element plutonium which vary from Pu(III) to Pu(VII). Plutonium could be used in nuclear power stations or to make nuclear bombs. Understanding the oxidation states of plutonium will help to clear up the waste from the Cold War.

13 How To Work Out Oxidation States Combining half-equations When you are given two half-equations, sometimes you will need to join them to make one full equation. The aim here is make sure that the same number of electrons that are donated by one half-equation are accepted by the other. How to combine two half-equations The easiest way to combine two half-equations is to work in steps. These worked examples will help you when you meet more difficult questions. Refer to the examples while reading these steps; Step 1 Write out the two half-equations. Step 2 Note the number of electrons each half-equation gains or loses. So that both equations involve the same number of electrons, you may have to multiply up one or both equations. Step 3 Multiply up the reactants and products. Step 4 Write all the reactants together, and the products together. Step 5 There should be the same number of electrons on both sides of this equation. Cancel them. What is left is the full balanced equation.

14 How To Work Out Oxidation States

15 How To Work Out Oxidation States

16 How To Work Out Oxidation States The cell in the photograph uses zinc and silver oxide to store the energy. The half-equations are: Zn(s) + 2OH (aq) Zn(OH) 2 (s) + 2e Ag 2 O(s) + H 2 O(l) + 2e 2Ag(s) + 2OH (aq) The overall discharge equation is: Zn(s) + Ag 2 O(s) + H 2 O(l) Zn(OH) 2 (s) + 2Ag(s) The cell has a high energy density, but is very expensive.

17 How To Work Out Oxidation States

1.7 REDOX. Convert these to ionic and half equations and you can see clearly how the electrons are transferred:

1.7 REDOX. Convert these to ionic and half equations and you can see clearly how the electrons are transferred: 1.7 REDOX Oxidation and Reduction: Oxidation and reduction reactions can be identified by looking at the reaction in terms of electron transfer: Our understanding of oxidation and reduction was limited