OXIDATION NUMBERS If a redox equation is very complicated, it is sometimes hard to work out what has been reduced and what the reducing agent is. When this happens you can use a sort of ‘electronic book-keeping’, called oxidation numbers, to make the task easier. Oxidation numbers are assigned to atoms, in elements, compounds or ions, according to a set of rules. Once you have assigned them you can easily decide whether oxidation or reduction has taken place: When an element increases its oxidation number it has been oxidised. When an element decreases its oxidation number it has been reduced. If no species changes its oxidation number, the equation does not represent a redox reaction. Rules for assigning oxidation numbers. 1] The oxidation number of an atom in a free, uncombined element is zero. E.g. H2, O2, Mg, Na, S8 2] The oxidation number of an atom in a simple monatomic ion is equal to the charge on the ion. E.g. O2- O.N. of oxygen = -2, Mg2+ O.N. of magnesium = +2. 3] In a polyatomic ion the sum of the oxidation numbers of the atoms in that ion equals the charge on the ion. E.g. SO42- the O.Ns of sulfur and oxygen must add to –2. 4] In compounds, the sum of the oxidation numbers of all atoms in the compound must add to zero. 5] The oxidation number of oxygen in compounds is –2 except in peroxides, where it is –1. (The only peroxide you are likely to meet is hydrogen peroxide H2O2.) 6] The oxidation number of hydrogen in compounds is +1 except in the case of hydrides where it is –1. (A hydride is a compound containing a reactive metal and hydrogen e.g. lithium hydride, LiH). Example: Determine the oxidation number of each atom in the equation: 2Mg + O2 2MgO Both magnesium and oxygen on the left are elements, so Mg and O both have an oxidation number of 0 (rule 1). After reaction the product is ionic. Magnesium will have an oxidation number of +2, since it is in the form Mg2+ (rule 2). The oxidation number of oxygen will be –2 (rule 5). 0 0 2 2 2 Mg 2 O 2 2 Mg O Test Yourself: 1] Write oxidation numbers for every atom in the following: a. Hg b. HNO3 c. NH3 d. MnO2 3+ e. Cr f. H2O2 g. NO2 h. SO422] Which of the following are redox reactions? Identify the oxidised and reduced species in each case. a. CaO(s) + SiO2(s) CaSiO3(s) b. 2Rb(s) + Br2(l) c. 2HI(aq) + H2O2(aq) d. H2SO4(aq) + 2KOH(aq) 2RbBr(s) I2(s) + 2H2O(l) K2SO4(aq) + 2H2O(l) BALANCING REDOX EQUATIONS Write the formula for the reactant and its product Identify a reactant and its product Are atoms other than hydrogen & oxygen balanced ? No Balance atoms other than hydrogen and oxygen Yes Are the oxygen atoms balanced ? No Balance oxygen atoms by adding water molecules No Balance the hydrogen atoms by adding H + ions Repeat for other reactant & product Yes Are the hydronen atoms balanced ? Yes Are the charges balanced ? Yes Are the number of electrons in the two half equations equal ? Add two half equations to get net ionic equation. Yes Balance the charges by adding electrons to the side with the greatest positive charge. No No Balance electrons by multiplying by appropriate coefficients When the reaction is carried out in alkaline conditions we have to put in an extra step to the balancing of the equation. Once the H has been balanced by adding H+, add an equal amount of OH- ions to both sides of the equation, to convert all H+ to H2O molecules. (Then simplify by cancelling out water molecules.) Write a balanced equation for the following redox reactions: -- 1] Cr(OH)4 + Na2O2 2] MnO4 + HSO3 -- -- -- -- CrO4 + OH 2-- MnO2 + SO4 (Basic Conditions) (Basic Conditions) OXIDISING AGENTS You must know the species involved in each of these reactions, their colours and be able to write balanced half equations for each reaction. SUBSTANCE Hydrogen peroxide HALF-EQUATION H2O2 + 2H+ + 2e 2H2O colourless Halogen, X2 (Cl2, Br2 or I2) colourless 2X- X2 + 2e Cl2 (pale green) colourless Br2 (deep red but orange-red in solution) I2 (purplish black but brown in aqueous solution) Permanganate ion (acid solution) Permanganate ion (neutral solution) Permanganate ion (strongly alkaline solution) Dichromate ion (acid solution Nitric acid (concentrated) Nitric acid (dilute) MnO4 - + 8H+ + 5e Mn2+ + 4H2O purple colourless MnO4 - + 4H+ + 3e MnO2 + 2H2O purple black MnO - + e MnO42- purple green Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O orange green NO3 - + 2H+ + e NO2 + H2O colourless brown NO3 - + 4H+ + 3e colourless NO + 2H2O colourless * (*Instantly reacts with oxygen in the air to form brown nitrogen dioxide gas) Fe3+ + e Iron (III) salt Copper (II) salt Iodate (or bromate) ion Fe2+ yellow pale green Cu2+ + 2e Cu blue pink / brown IO3 - + 6H+ + 6e I - + 3H2O colourless colourless REDUCING AGENTS SUBSTANCE Zinc Zn silvery Magnesium Mg silvery Iron Hydrogen Sulfur dioxide Iron (II) salt Sulfite ion Iodide ion colourless Mg2+ + 2e colourless Fe Fe2+ + 2e grey pale green H2 2H+ + 2e colourless colourless SO2 + 2H2O SO42- + 4H+ + 2e colourless colourless Fe2+ Fe3+ + e pale green yellow SO32- H2O SO42- + 2H+ + 2e colourless Thiosulfate ion HALF-EQUATION Zn2+ + 2e 2S2O32- colourless S4O62- + 2e colourless colourless 2I - I2 + 2e colourless brown ELECTROCHEMICAL CELLS All chemical reactions involve some sort of energy change. Normally this is an exchange of heat (thermal energy) between the system and its surroundings. If the reaction is exothermic the system is a source of thermal energy. In redox reactions (electron –transfer reactions) the energy transfer can either be thermal energy or electrical energy. e.g. the reaction between zinc metal and a copper salt. If zinc metal is added to a solution of copper ions, copper metal and zinc ions are formed and the solution becomes warm: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) H is negative Carried out in this way the reaction takes place at the surface of the zinc metal where electrons are transferred directly from zinc atoms to copper ions. If, however, the reactants are kept apart, the electrons can be led through a wire instead of being transferred directly. The energy appears as electrical energy. A device which does this is called an electrochemical cell. An early electrochemical cell was the Daniell cell. It was based on the reaction between zinc and copper ions and was once used to operate railway signals, telegraph relays and doorbells. In the inner compartment: Zn(s) Zn2+(aq) + 2e- oxidation In the outer compartment: Cu2+(aq) + 2eOverall reaction: Zn(s) + Cu2+(aq) Cu(s) reduction Cu(s) + Zn2+(aq) redox [The zinc electrode is coated with mercury to stop it reacting with the sulfuric acid.] The potential of the Daniell cell is 1.10 volts. In principle any oxidation-reduction reaction can be used to operate an electrochemical cell, but many have difficulties associated with them. Cells and Half Cells An electrochemical cell consists of two half-cells connected by a salt bridge. A half-cell, is an electrode and the couple it is in contact with, and is where the reaction takes place; oxidation in one, reduction in the other. Each half-cell (or electrode) consists of: (a) the oxidised and reduced forms of one of the reactants, one or both of them in water solution. In identifying a half-cell both forms, in the order oxidised form / reduced form, must be stated. E.g. Zn2+/ Zn, Cu2+/ Cu, Fe3+/ Fe2+ etc (b) A piece of metal through which the electrons pass to or from the external circuit. A voltage or potential difference is generated between the metal and the solution. The electrode potential of a metal, measured in volts, and is caused by the charge built up on the metal rod due to the dynamic equilibrium established between the metal atoms and the metal ions in solution. M+(aq) + e M(s) The position of the equilibrium, the charge and size of the electrode potential, depends on: The nature or reactivity of the metal – the electrode potential is more negative for more reactive metals as they have a greater tendency to lose electrons to form stable ions. The ionic concentration [M+] – increasing the concentration of [M+] will alter the equilibrium in favour of the formation of M(s) , so the electrode potential becomes more positive. The temperature – changes in temperature affect exothermic and endothermic reactions differently. It is impossible to measure the true potential difference of a metal in a solution of its ions because there is no complete circuit, so electrodes or half-cells must be connected to a second (reference) electrode to give a comparison of energy difference (voltage). One half-cell will tend to lose electrons and the other half-cell will tend to gain these electrons: this allows the relative reduction potential of each half-cell to be calculated. THE STANDARD HYDROGEN ELECTRODE The hydrogen electrode is used as a standard or reference electrode to measure the electrode potential of any half-cell. It is given the reference potential of 0.00 V. 2H+(aq) + 2e - H2(g) Eo = 0.00 V Eo is measured under standard conditions: All solutions are 1 mol L-1 all gas pressures are 1 atm or 100 kPa temperatures are 25oC (298 K) An inert electrode is used to carry the electrons to and from the hydrogen half-cell, as neither hydrogen gas nor ions can carry electrons. Platinum is the metal used because it is inert it catalyses the reaction that produces hydrogen + -2H (aq) + 2e H2(g) Once the standard hydrogen electrode is set up, and because it has a reference potential of 0.00 V , it can be used to determine the electrode potential of all other half-cells. Pure H2(g) at 298K and 1 atm Platinum electrode Acid solution containing 1.0 mol L-1 H+(aq), 298 K Holes in side of glass casing to allow Hydrogen gas to escape so pressure is Maintained at 1 atm. Measuring the Electrode Potential of a Half-Cell The half-cell whose Eo value is to be determined is set up with the hydrogen half-cell to make a complete cell. A salt bridge is used to connect the two half-cells and the potential difference of the cell is measured, using a voltmeter. Since the hydrogen half-cell is a reference half-cell with an electrode potential of zero, the reading on the voltmeter is solely due to the other half-cell. Half-Cell Conventions All half-cells consist of a species in its oxidised and reduced states and an electrode, which takes electric current to or away from the half-cell. Half-cells are written with the oxidised form of the species first. There are three types of half-cells: A metal rod and its aqueous solution A gas and its aqueous ions Two aqueous ions in solution together A Metal and its Aqueous Ions For metal/metal ion half-cells the metal itself acts as an electrode. The half-cell for a metal M(s) and its aqueous ions M+(aq) is written as: M+/M (oxidised form / reduced form). The slash (/) indicates phase boundary between reactants. For the zinc / zinc ions half-cell, written as Zn2+/Zn (oxidised form first): The half-reaction is Zn2+(aq) + 2e Zn(s) Eo = -0.76 V Zn2+(aq), the oxidised form of zinc can act as an oxidant Zn(s), the reduced form of zinc, can react as a reductant The solid zinc metal Zn(s) acts as the electrode. Eo values are given for the left to right direction in equations. Because E o is negative in the Zn2+| Zn half-cell, the reaction which usually occurs is: Zn(s) Zn2+(aq) + 2e – A Gas and its Aqueous Ions Gas / aqueous ion half-cells require an inert electrode to carry the current in or out of the half-cell, because gases cannot form conducting electrodes. Platinum Pt, or carbon C (in the form of graphite), are often used as electrodes. The half-cell for a gas X2 (g) and its aqueous ions X- (aq) is written as: X2 (g) , X –(aq)/ inert electrode. (The / is used to separate the conductor from the rest of the half cell.) For the chlorine/chloride half-cell, written as, Cl2 (g), Cl –(aq) / Pt (or C): The half-reaction is Cl2 (g) + 2e2Cl-(aq) Eo = +1.36 V Cl2(g) is the oxidant Cl-(aq) is the reductant Because Eo is positive in the Cl2, Cl-(aq) half-cell, the reaction which usually occurs is: Cl2(g) + 2e2Cl-(aq) Two Types of Aqueous Ions Ions can be oxidised and reduced to form other ions, so aqueous ion / aqueous ion half-cells exist. Ions, like gases cannot form conducting electrodes, so an inert electrode is used in these half-cells. The half-cell for the oxidised for of an aqueous ion (M2+(aq)) and the reduced form of the aqueous ion (M+(aq)) is written as: M2+(aq), M+(aq) / inert electrode. (The comma is used to separate the two reactants in the same solution and the / indicates a phase separation between the solution and the electrode.) For the iron (III) | iron (II) half-cell, written as Fe3+(aq), Fe2+(aq) / Pt (or C): The half-reaction is Fe3+(aq) + eFe2+(aq) 3+ Fe (aq) is the oxidant Fe2+(aq) is the reductant Eo = +0.77 V Because Eo is positive in the Fe3+(aq),Fe2+(aq) half-cell, the reaction that usually occurs is: Fe3+(aq) + eFe2+(aq) Electrode Potential Values The values of electrode potential provide information on: The strength of metals as reducing agents (high negative E o values) The strength of non-metals as oxidising agents (high positive Eo values) The prediction of the spontaneity of redox reactions. If the overall value of E o for a cell is positive, the reaction will occur spontaneously. Calculating the EMF or voltage of electrochemical (galvanic cells) Reducing Agents and Eo values Electrode potentials involving metals indicate: How readily hydrated metal ions in solution gain electrons (are reduced) to form metal ions. The more positive the Eo value, the more likely it is that the metal ion will form the metal. How readily metal atoms release electrons (are oxidised) to produce hydrated ions in solution. The most reactive metals (those that are the best reductants), are those that have the most negative electrode reduction potential values. Arranging the metals in order of their electrode potential (Eo) values produces the electrochemical series. Electrode potentials are good indicators of metal reactivities. The two exceptions are lithium and calcium, which are chemically less reactive than their Eo values would indicate. Electrochemical series Lithium Li Potassium K Calcium Ca Sodium Na Magnesium Mg Aluminium Al Zinc Zn Iron Fe Lead Pb Copper Cu Silver Ag Gold Au Metal ion/atom half-equation Li+ + e Li K+ + e K Ca2++ 2e Ca Na+ + e Na Mg2+ + 2e Mg Al3+ + 3e Al Zn2+ + 2e Zn Fe2+ + 2e Fe Pb2+ + 2e Pb Cu2+ 2e Cu Ag+ + e Ag Au+ + e Au Electrode potential Eo / V -3.02 -2.92 -2.87 -2.71 -2.30 -1.66 -0.76 -0.47 -0.13 +0.34 +0.81 +1.50 Metals with negative Eo values can reduce aqueous hydrogen ions to hydrogen gas. If the electrode potential is positive, the metal atoms do not reduce hydrogen ions and so they do not react with acids. Oxidising Agents and Eo values Electrode potentials involving non-metals indicate how readily ions or molecules accept electrons (are reduced) to form other species. The most reactive non-metals, (those that are the best oxidants), have the most positive electrode potentials values. The Eo values for the common oxidising agents give a good indication of their strengths as oxidants. Electrochemical series Fluorine Chlorine Oxygen Bromine Iodine Nitrogen Sulfur F Cl O Br I N S Non-metal reduction half-equation F2 + 2e 2FCl2 + 2e 2ClO2 + 4H+ + 4e 2H2O Br2 + 2e 2BrI2 + 2e 2IN2 + 6H+ + 6e 2NH3 S + 2H+ + 2e H2S Electrode potential Eo / V +2.87 +1.36 +1.23 +1.07 +0.54 +0.27 +0.14 SUMMARY OF REDOX THEORY Redox reactions involve competition for electrons. When a species is reduced it gains electrons; when a species is oxidised it loses electrons. When electricity flows through a molten salt, or through a solution of an electrolyte, the substance is split up in a chemical process called electrolysis. At the anode, anions (negatively charged ions) are oxidised. At the cathode, cations (positively charged ions) are reduced. The chemical changes at each electrode are summarised using half-equations. The overall cell reaction can be found by adding together the two half-equations. Electricity may be produced whenever two metals are immersed in conducting solutions, forming an electrochemical cell. Each half-cell reaction has its own tendency to attract electrons, which is measured by the electrode potential E of the half-cell. The half-cell against which all other electrode potentials are measured is the standard hydrogen half-cell or standard hydrogen electrode. The standard electrode potential Eo of a half-cell is the electrode potential of that half-cell relative to the standard hydrogen electrode under standard conditions. Eo is a measure of the oxidising or reducing power of the species in it. The standard conditions for electrochemical measurements are: all solutions have a concentration of 1 mol L-1 and all measurements are made at a temperature of 25oC and 101.3 kPa pressure. Standard electrode potentials can be used to predict the likelihood of a reaction proceeding spontaneously. If the Eo value for a cell is positive the reaction will be spontaneous. Cell voltage- the potential of the right-hand electrode with respect to the lefthand electrode Half-cell notation E.g. zinc-copper cell Zn(s)/ Zn2+(aq) Cu2+(aq) /Cu(s) Eo = +1.10V Left-hand electrode Right-hand electrode Electrolyte in contact with left – hand electrode Salt bridge Electrolyte in contact with right – hand electrode In this reaction the cell voltage is positive i.e. the right-hand electrode is positive with respect to the left-hand electrode. Electrons flow from the zinc to the copper. By convention the cell notation always refers to the reaction taking place from left to right. Eocell = ERHE - ELHE The salt bridge: Without a salt bridge the half-cell containing the zinc would slowly become positively charged as the electrons left it. The copper half-cell would slowly become negatively charged. With the salt bridge present ions are able to move in and out of the solutions keeping both half-cells electrically neutral. The salt used in the salt bridge is chosen so that it does not react with the ions in either half-cell. Oxidising and Reducing agents: The standard electrode potential of a half-cell is a measure of the oxidising or reducing power of the species in it – i.e. their ability to compete for electrons. In general the stronger an oxidising agent, the more positive its electrode potential. A strong reducing agent has a large negative electrode potential. When looking at a table of standard electrode potentials for a number of half reactions, the substances to the left of the double arrows are oxidising agents, becoming reduced when they react in the direction left to right. The best oxidising agent have a large positive Eo value and are on the left of the equations. Substances to the right of the double arrows are reducing agents, becoming oxidised when they react in the direction right to left. The best reducing agents are on the right of the equations and have a large negative Eo value. Ag+(aq) + e - Ag(s) E.g. Oxidising agent Reduced during left to right reaction Also known as oxidant The more positive the Eo value, the more powerful the oxidising agent Reducing agent Oxidised during right to left reaction Also known as reductant The more negative the Eo value, the more powerful the reducing agent Half-Cell conventions: In a cell diagram the following conventions are used: A single / represents a phase boundary between reactants e.g. between a solid and a solution A double vertical line represents a boundary between two solutions (usually a salt bridge) A comma is used to separate two reactants in the same solution, or a gas from an aqueous ion. The conductors are written on the extreme left and right of the cell diagram. When there is no metal in a redox couple, an electrical connection is made using an inert electrode such as platinum or carbon. The half-cell is then written as: e.g. Fe3+(aq),Fe2+(aq) / Pt(s). The least oxidised species present is written next to the electrode. The hydrogen electrode is written as: 2H+(aq) , H2(g) / Pt(s). Uses of Electrode Potentials: Calculation of predicted cell emfs (voltages) Prediction of preferred direction of redox reactions Comparison of relative strengths of oxidising and reducing agents ELECTROLYSIS CHEMICAL REACTIONS DRIVEN BY ELECTRICITY You have already seen that when a spontaneous redox reaction occurs, an electric current is produced. If we use electric current from a battery or power supply and pass it through a conducting solution redox reactions will be produced. This process is called electrolysis. During electrolysis electrical energy is converted to chemical energy. The reactions that occur in electrolytic cells are essentially the opposite to those occurring in electrochemical cells. Electrolytic Cell Electrical Energy Chemical Energy Electrochemical Cell Reactions in electrolytic cells would not normally happen without the application of electrical energy and are called non-spontaneous reactions. Chemicals formed by electrolysis are often difficult to obtain by other means and many useful materials are manufactured by this process. Important applications of electrolysis include: plating a thin film of metal on surfaces of other metals to improve appearance or prevent corrosion. (Electroplating) extraction of reactive metals such as sodium, magnesium and aluminium from their ores industrial preparation of sodium hydroxide, chlorine and hydrogen. recharging of car batteries and other rechargeable cells such as Ni-Cads. refining copper metal increasing the thickness of the surface oxide layer on aluminium metal to improve its resistance to corrosion. In an electrolytic cell the reactions that occur and redox reactions. Consider the electrolysis of a solution of zinc iodide. If an electric current is passed through a solution of zinc iodide, a grey deposit is formed on the cathode (negative electrode) and a brown colour forms at the anode (positive electrode). It is easy to show that the grey deposit is zinc and the brown substance is iodine. The reactions that have occurred are: At the cathode: At the anode: Zn 2+ (aq) + 2e -- 2I (aq) -- Zn(s) I2(aq) + 2e -- The cathode reaction is reduction and the anode reaction is oxidation. In electrolysis Oxidation occurs at the anode Reduction occurs at the cathode Questions 1. What is the main difference between electrochemical and electrolytic cells? 2. What happens during electrolysis ? 3. Give at least three practical uses of electrolysis. COMPARISON OF ELECTROLYTIC AND ELECTROCHEMICAL CELLS Electrolytic cells A non-spontaneous reaction. Reactions are forced by an applied voltage which must be greater than that which the cell can produce. Electrons flow from the positive electrode. Converts electrical energy to a chemical reaction. Electrodes are usually immersed in a common electrode. Electrochemical cells A spontaneous reaction. Electrons flow from the negative electrode. Converts a chemical reaction to electrical energy. Two half-cells are often used with separate electrodes. Electrolytic cell: External voltage source pushes electrons through circuit. Electrons absorbed from the external circuit (cathode) Migration of negative ions Electrons released to external circuit (anode) Migration of positive ions Electrochemical cell: Electrons flow through metallic conductor of external circuit Electrons absorbed Electrons released (anode) Migration of negative ions Migration of positive ions (cathode) Similarities: The electrolyte is the substance that conducts electricity within the cell. Electric charge is carried by anions to the anode and by cations to the cathode. The electrode where oxidation occurs is called the anode. The electrode where reduction occurs is called the cathode. In the external circuit, electrons travel through the wire from the anode to the cathode. Key ideas of electrolysis: The products of electrolysis depend on: The reactivity of the ions present The type of electrode used The concentration of the reactants The amount of current used The temperature e.g. molten NaCl produces Na and Cl2 dilute aqueous NaCl produces H2 and O2 (as for H2O) For electrolytic cells: Oxidation takes place at the anode (sign +ve) Reduction takes place at the cathode (sign –ve) Electrons flow from the anode to the cathode Electric current is used to cause a chemical reaction The applied voltage must be greater then the voltage of the cell If more than one reaction is possible the following rules apply Metal ions: Whether a metal is produced depends on its position in the reactivity series relative to the position of water. In a concentrated solution the metal will form if it is below water in the activity series. If it is above water, hydrogen gas forms. If there are 2 possible products, the one lower on the activity series will form. Non-Metal ions: A simple ion will always form its element. If a polyatomic ion is present, oxygen will form in preference.