REDOX EQUILIBRIA AND FEASIBILITY OF A REACTION Oxidizing agent Reducing agent
Oxidation-Reduction Reactions Electron transfer reactions Electrons transferred from one substance to another Change in oxidation number of element during reaction indicates redox reaction has occurred Important class of chemical reactions that occur in all areas of chemistry & biology
Oxidation Reduction Reactions
Oxidation Reduction Reaction Oxidizing Agent Substance that accepts e 's Accepts e 's from another substance Substance that is reduced Cu + + e Cu Reducing Agent Substance that donates e 's Releases e 's to another substance Substance that is oxidized Zn Zn + + e
Common oxidizing and reducing agents Common oxidizing agents Common reducing agents O (oxygen) H (hydrogen) O 3 (ozone) CO (carbon (II)oxide) F (fluorine) Fe (iron) Br (bromine) Zn (zinc) H SO 4 (sulfuric acid) Li (lithium) H O (hydrogen peroxide) C (carbon) K Cr O 7 (potassium dichromate) SO (sulfur(iv) oxide) KMnO 4 (potassium permanganate) KI (potassium iodine) 5
Guidelines For Redox Reactions Involves processes. Oxidation & reduction always occur simultaneously. Can't have one without the other Total number of electrons lost by one substance = total number of electrons gained by second substance For a redox reaction to occur, something must accept electrons that are lost by another substance
Ox 1 Red 1 Cu Cu + +ē Oxidized form Reduced form Conjugate redox pair Red Ox Zn -ē Zn + Reduced form Oxidized form Conjugate redox pair Cu + + Zn Zn + + Cu Ox 1 + Red Red 1 + Ox Net Conjugate redox process
If a piece of zinc metal is placed in a solution of copper (II) sulphate, the blue colour of the copper sulphate slowly fades, the zinc dissolves and red-brown copper metal takes its place. This reaction produces energy, which is lost as heat if we simply carry out the process in a single reaction vessel. However, if we separate the reaction in two halfcells, we can harness the energy through the flow of electrons taking place between the cells. This combination of half-cells is an electrochemical cell (galvanic cells or voltaic cells).
If a rod of metal is dipped into a solution of its own ions, an equilibrium is set up. For example: Zn Zn + + e- zinc metal strip zinc sulfate solution (1 mol dm -3 ) 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.
17_36 e e e Porous disk e Oxidation Reduction e Reducing agent Oxidizing agent e (a) Anode (b) Cathode
Galvanic Cell A device in which chemical energy is converted to electrical energy. It uses a spontaneous redox reaction to produce a current that can be used to generate energy or to do work.
Zinc (Zn) Zn Zn + + e- Zn +
Components Electrodes: conduct electricity between cell and surroundings Electrolyte: mixture of ions involved in reaction or carrying charge Salt bridge: completes circuit (provides charge balance)
Zn + + Cu Zn + + Zn Zn gives up electrons to Cu pushes harder on e - greater potential energy greater electrical potential Spontaneous reaction due to relative difference in metals abilities to give e - ability of e - to flow
Electromotive Force (e.m.f.) Electromotive Force (e.m.f.): The pull or driving force on electrons. Measured voltage (potential difference) By using a high resistance voltmeter, the current in the external circuit is virtually zero and the cell registers its maximum potential difference. This maximum potential difference is called the e.m.f. This e.m.f. gives a quantitative measure of the likelihood of the redox reaction taking pace in the cell.
17_363 e e e E cell = +1.1 V e Zn(s) Zn + SO 4 1. M Zn + solution Cu + SO 4 1. M Cu + solution Cu(s) Anode Cathode
It is impossible to obtain the electrode potential for a single half-cell. E.m.f. values can only be measured for a complete circuit with two electrodes. This can be done by arbitrarily assigning an electrode potential of zero to one particular half-cell. Other half-cell can then be compared with this standard. The standard chosen for electrode potentials is standard hydrogen electrode (SHE). This consists of H gas at one atmosphere pressure and 5 o C bubbling around a platinised platinum electrode. The electrode is immersed in a 1. moldm -3 solution of H + ions. Hydrogen is adsorbed on the platinum and an equilibrium is established between the adsorbed layer of H and H + ions in the solution.
Standard Reduction Potentials The standard electrode potential (E ) of an electrode ( halfcell) is thus defined as the e.m.f. of a cell in which the electrode on the left is a standard hydrogen electrode and that on the right is the standard electrode in question. E values for reduction half-reactions with solutes at 1M and gases at 1 atm Standard electrode potentials are sometimes called standard reduction potentials because they related to the reduction of the more oxidized species. It is also a measure of the strength of the oxidising agent to accept electrons.
The following table shows standard electrode potentials of some common redox systems :
THE NERNST EQUATION The processes that occur at an electrode surface are reversible and are governed by the equilibrium law. If the concentration of a reactant or product is altered, the position of equilibrium also alters and therefore the electrode potential changes. The exact relationship between the electrode potential of a half-cell and the concentrations of the ionic species involved is given by the Nernst equation : E E RT a,59 Ox Ox / Red ln EOx / Red lg nf ared n a a Ox Red where E = electrode potential E = standard electrode potential R = gas constant T = absolute temperature n = number of electrons transferred F = Faraday s constant (the charge of one mole of electrons) An increase in temperature results in the electrode potential becoming more positive compared with the standard hydrogen electrode.
The standard electrode potentials enable us to predict whether a redox reaction is spontaneously or not. In general, redox reactions with an overall positive standard cell potential ( cell e.m.f. ) are energetically feasible whereas those with an overall negative value are not so.
PREDICTING THE DIRECTION OF REDOX REACTIONS It is possible to use standard electrode potentials to decide on the feasibility of a reaction. Electrodes with more negative electrode potentials have a lower tendency to accept electrons. Zn + + e- Zn E ө = -.76 V Cu + + e- Cu E ө = +.34 V When a pair of electrodes are connected, electrons flow from the more negative to the more positive. The signs of the electrodes can be used to predict the direction of the reaction. Zn + Cu + Zn + + Cu Zn + + Cu Zn + Cu + feasible not feasible
EXAMPLE What reaction would occur if Fe 3+ /Fe + and Cu + /Cu half cells were connected? Step 1: write the equations for the two half reactions: Fe 3+ + e- Fe + E θ = +.77 V Cu + + e- Cu E θ = +.34 V Electrons flow to the more positive terminal of a cell, which is Fe 3+ /Fe +. This half cell will accept electrons, a reduction reaction occurs, and the half equation is: Fe 3+ + e- Fe +
An oxidation reaction occurs in the other half cell. Electrons are produced and the half equation is: Cu Cu + + e- Step : combine the two half equations to give a full equation for the reaction. The number of electrons supplied and donated must be equal, so the reaction in the Fe 3+ /Fe + half cell must occur twice for each Cu + /Cu half cell reaction: Fe 3+ + Cu Cu + + Fe + The reverse reaction is not feasible.
Example What reaction would occur? E E 1. 37 Cl / Cl Fe 3 / Fe. 77 E 3. 3 F / F E. 54 I / I V V V V Cl Br Br I Cl Cl Cl I Cl 3 Fe Cl Fe E 1. 8 Br / Br V Fe 3 I I Fe aox Only in case, lg a Red Cl F Cl F
Conclusions: Cl is a stronger oxidizing agent than Br ion, oxidizing iodide, chloride and bromide ions. The first reaction therefore happens spontaneously in the direction written, with a decrease in the energy of the system. I is a weaker oxidizing agents than Cl, so the reaction written here do not happen spontaneously. Instead, the reverse reactions are the energetically favorable ones, in which Cl ions oxidize iodide ions to I. Analogically the directions of other reactions can be explained.
Limitation of prediction : Theoretically, redox reactions having a positive cell e.m.f. are feasible reactions and will take place spontaneously. However, since electrode potentials are affected by changes in concentration and temperature, a prediction based on standard electrode potential values may not be valid under conditions which are not standard.
Solve a problem Given the following reduction potentials: Cu + + e - Cu; Ni + + e - Ni; E =.34 V E = -.3 V Predict whether the following reaction will take place: Cu + Ni + Cu + + Ni
FACTORS THAT AFFECT THE DIRECTION OF REDOX REACTIONS а) Precipitation of reducing or oxidizing agents Cu 4I CuI I E.17 В E. 54В Cu Cu I / I / According to the standard potentials reverse reaction should occur, but in reality the formation of CuI precipitate leads to the decreasing of Cu / Cu potential and as the result the direct reaction takes place
b) Complexation of reducing or oxidizing agents For example Fe 3+ /Fe + Е = +,77В if there are F -, ions in the solution, the potential of the cell became less because of complex ion [FeF 6 ] 3- formation.
в) рн HAsO 4 I H HAsO 3 I HO E. 56 HAsO4 / HAsO3 В > E. 54 I / I В The equilibrium is slightly shifted from lift to right but
The change of the рн affects only the following potential 3 E HAsO4 / HAsO in acidic medium > >,54V and the reaction will occur from left to right, in basic medium (in the presence of NaHCO 3 ) the reaction will occur from right to left (рн = 8, с(н + ) = 1-8 ) E.59 c( HAsO4 )(1 lg c( HAsO ) 8.56. 9 HAsO4 / HAsO3 3 ) В.9 V <,54 V
EQUILIBRIUM CONSTANTS FOR REDOX REACTIONS Many redox reactions are reversible and the Low of mass actions can be applied to describe them. aox 1 bre d are d 1 box K c c a a b (Re d1) c ( Ox b ( Ox ) c (Re d 1 ) )
Two simultaneous processes are characteristic for redox reactions a Ox1 ab e a Re d 1 Reduction process b Re d ab e b Ox Oxidation process Let s write Nerst expression for both processes E равн. E 1.59 lg ab a c ( Ox1) a c (Re d ) 1 E равн. E.59 lg ab b c ( Ox) b c (Re d )
At equilibrium: E 1 a.59 c ( Ox1) lg a ab c (Re d ) 1 E b.59 c ( Ox) lg b ab c (Re d ) E 1 E b.59 c ( Ox) (lg b ab c (Re d ) a c ( Ox1) lg ) a c (Re d ) 1 ab( E E.59 ) b c ( Ox) lg b c (Re d ) a c ( Ox1) lg a c (Re d ) 1 1 lg K K 1 ab( E 1 E.59 ) where Е 1 and Е standard electrode potentials The bigger value of К the bigger difference between Е 1 and Е.
Example: For the following redox reaction: Fe 3+ + Sn II Sn IV + Fe + K 1 1( E Fe 3 / Fe E.59 Sn IV / Sn II ) 1 (.77.15).59 1 1 The value of K is very large (К >>> 1), it means that this reaction is reactant-favored.
APPLICATION OF REDOX REACTIONS IN CHEMICAL ANALYSIS 1. For ion separation Example: chromium and aluminium ions can be separated using hydrogen peroxide in ammonia solution Al 3 3NH 3 3HO Al( OH ) 3 3NH 4 3 Cr 3HO 1OH CrO 4 8HO. For ion detection (manganese, bismuth, chromium etc.) Example: Mn 5PbO 4H MnO 4 5Pb HO 3. For dissolution of different metals and salts Example: 3 3 4 HgS 1HCl HNO H [ HgCl ] 3S 4H O NO 4. For stabilization of reducing agent solutions.