Topic 9: Redox Processes

• Addition of oxygen • Loss of hydrogen • Loss of electrons

• Loss of oxygen • Addition of hydrogen • Gain of electrons

• The oxidation number of an atom is the charge that would exist on an individual atom if the bonding were completely ionic • It is like the electronic ‘status’ of an element • The oxidation number is the number of electrons which must be added or removed to become neutral • The oxidation number is always written with the charge before the number • In molecules or compounds, the sum of the oxidation numbers on the atoms is zero

• used to tell if oxidation or reduction has taken place • helps to work out what has been oxidised and/or reduced • used to construct half equations and balance redox equations

• Many atoms, such as S, N and Cl, can exist in a variety of oxidation states • The oxidation number of these atoms can be calculated by assuming that the oxidation number of the other atom is fixed

• The oxidation number of any uncombined element is 0 • Many atoms or ions have fixed oxidation numbers in compounds (group 1 elements are always +1, fluorine is always -1, oxygen is -2, hydrogen is +1) - there are a few exceptions • The oxidation number of an element in a mono-atomic ion is always the same as the charge (Cl- is -1) • The sum of the oxidation number in a compound is 0 • The sum of the oxidation number in an ion is equal to the charge on the ion • The more electromagnetic element is always given the negative oxidation number

• An oxidising agent is a substance that oxidises another atom or ion by causing it to lose electrons • An oxidising agent itself gets reduced – gains electrons • Therefore, the oxidation number of the oxidising agent decreases • Some substances can act both as oxidising and reducing agents

• A reducing agent is a substance that reduces another atom or ion by causing it to gain electrons • A reducing agent itself gets oxidised – loses/donates electrons • Therefore, the oxidation number of the reducing agent increases • Some substances can act both as oxidising and reducing agents

For a reaction to be recognised as a redox reaction, there must be both an oxidising and reducing agent

• Transition metals are characterized by having variable oxidation numbers. • Oxidation numbers can be used in the names of compounds to indicate which oxidation number a particular element in the compound is in • Where the element has a variable oxidation number, the number is written afterwards in Roman numerals. • This is called the STOCK NOTATION • For example, iron can be both +2 and +3 so Roman numerals are used to distinguish between them Fe2+ in FeO can be written as iron(II) oxide Fe3+ in Fe2O3 can be written as iron(III) oxide

• Write the unbalanced equation and identify the atoms which change in ox. no. • Deduce the oxidation number changes • Balance the oxidation number changes • Balance the charges • Balance the atoms

• Metals can be ranked in order of reactivity • Metals higher in reactivity can displace less reactive metals from their compounds in solutions or from their oxides • The more reactive metal acts as a reducing agent • This allows metals to be ranked from most reactive (strongest reducing agents) to least reactive • The more reactive a metal is the better it is at pushing electrons onto less reactive metal ions. • The more reactive metal undergoes oxidation • The less reactive metal ion undergoes reduction

• In a titration, the concentration of a solution is determined by titrating with a solution of known concentration. • In redox titrations, an oxidizing agent is titrated against a reducing agent • Electrons are transferred from one species to the other • Indicators are sometimes used to show the endpoint of the titration. However, most transition metal ions naturally change colour when changing oxidation state

• manganate(VII) titrations • iodine-thiosulfate titrations

• A redox reaction occurs between acidified manganate (VII) ions and iron (II) ions: MnO4– (aq) + 8H+ (aq) + 5Fe2+ (aq) → Mn2+ (aq) + 5Fe3+ (aq) + 4H2O (l) • This reaction needs no indicator as the manganate (VII) is a strong purple colour which disappears at the end point, so the titration is self-indicating • This reaction is often used for the analysis of iron for example in iron tablets (health supplement)

• A redox reaction occurs between iodine and thiosulfate ions: 2S2O32– (aq) + I2 (aq) → 2I–(aq) + S4O62– (aq) • The light brown/yellow colour of the iodine turns paler as it is converted to colourless iodide ions • When the solution is a straw colour, starch is added to clarify the end point • The solution turns blue/black until all the iodine reacts, at which point the colour disappears. • This titration can be used to determine the concentration of an oxidizing agent, which oxidizes iodide ions to iodine molecules • The amount of iodine is determined from titration against a known quantity of sodium thiosulfate solution • This reaction can be used for the analysis of chlorine in bleach

• The Winkler Method is a technique used to measure dissolved oxygen in freshwater systems • Dissolved oxygen is used as an indicator of the health of a water body, where higher dissolved oxygen concentrations correlate with high productivity and little pollution

• The biological oxygen demand (BOD) is the amount of oxygen used to decompose the organic matter in a sample of water over a specified time period, usually 5 days, at a specified temperature • A high BOD indicates a greater quantity of organic waste in the water, which means a lower level of dissolved oxygen

• Voltaic (or Galvanic) cells generate electricity from spontaneous redox reactions • Instead of electrons being transferred directly from the zinc to the copper ions a cell is built which separates the two redox processes • Each part of the cell is called a half cell • If a rod of metal is dipped into a solution of its own ions, an equilibrium is set up • This is a half cell and the strip of metal is an electrode • The position of the equilibrium determines the potential difference between the metal strip and the solution of metal. This is known as an electrode potential. • Note that a chemical reaction is not taking place - there is simply a potential difference between the rod and the solution

• the nature of the ions in solution • the concentration of the ions in solution • the type of electrode used • the temperature

• If two different electrodes are connected, the potential difference between the two electrodes will cause a current to flow between them. Thus an electromotive force (emf) is established and the system can generate electrical energy • A typical electrochemical cell can be made by combining a zinc electrode in a solution of zinc sulphate with a copper electrode in a solution of copper sulphate • The circuit must be completed by allowing ions to flow from one solution to the other • This is achieved by means of a salt bridge - often a piece of filter paper saturated with a solution of an inert electrolyte such as KNO3(aq) • The e.m.f can be measured using a voltmeter • Voltmeters have a high resistance so that they do not divert much current from the main circuit • The combination of two electrodes in this way is known as an electrochemical or voltaic cell, and can be used to generate electricity

• positive electrode • most favours reduction • Reduction takes place • cations flow to the positive half cell since there is a surplus positive charge from the metals becoming cations in the negative half cell

• negative electrode • oxidation takes place • anions flow to the negative half cell to replace the negative charge of the electrons

• 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. • This process is called electrolysis, a word which comes from Greek and means “splitting by electricity”