A + B C +D ΔG = ΔG + RTlnKp Me n+ + ne - Me K p a a Me Me n a n e
1 mol madde 6.2 x 1 23 atom elektron yükü 1.62 x 1-19 C FARADAY SABİTİ: 6.2 x 1 23 x 1.62 x 1-19 = 96485 A.sn (= coulomb) 1 Faraday 965 A.sn = coulomb. Bir sistemden her geçen 965 coulomb yük için 1 ekivalen mol açığa çıkar.
Me n+ + ne - Me G nf G nf RT nf ln Kp ln a Me a Me n. a n e = ln a ln a n nln a Me Me e E E RT nf ln Kp G E Nernst Denklemi nf ΔE > da reaksiyon gerçekleşir. Çünkü, ΔG yi n e böldük. Q A. I. t nf
Electromotive Force emf potentialdifference (V) work (J) charge (C) 1 joule of work is produced or required when 1 coulomb of charge is transferred between two points in the circuit that differ by a potential of 1 volt.
Galvanic Cells (Voltaic Cells ) Galvanic cell - an electric cell that generates an electromotive force by an irreversible conversion of chemical to electrical energy; cannot be recharged. The electron flow from the anode to the cathode is what creates electricity. In a galvanic cell, the cathode is positive while the anode is negative, while in a electrolytic cell, the cathode is negative while the anode is positive.
Galvanic Cells
Standard Reduction Potentials
Standard Hydrogen Electrode 2H + (aq)+zn(s) Zn +2 (aq)+h 2 (g) Oxidation half-reaction Zn(s) Zn +2 (aq)+2e - Standard hydrogen electrode 2H + (aq)+2e - H 2 (g) The cathode consists of a platinum electrode in contact with 1 M H + ions and bath by hydrogen gas at 1 atm. We assign the reaction having a potential of exactly volts. E cell E H H 2 E ZnZn 2.76.76( V )
Copper-Zinc Voltaic Cells
The Cell Potentials E cell=e (cathode)-e (anode) reduction oxidation The value of E is not changed when a half reaction is multiplied by an integer. 2Fe +3 +2e - 2Fe +2 E (cathode)=.77 V Cu Cu +2 +2e - -E (anode)=-.34 V Cu+2Fe +3 2Fe +2 +Cu + E cell=e (cathode)-e (anode)=.43 V
Cell Diagrams For a copper-zinc voltaic cells Cu Zn ZnSO 4 (aq) CuSO 4 (aq) Cu Dashed line 1. A vertical line indicates a phase boundary. 2. A dashed vertical line indicates the phase boundary between two miscible liquid.
Pt L H 2 (g) HCl(aq) AgCl(s) Ag Pt R Anode: H 2 (g)=2h + +2e - (Pt L ) Cathode: [AgCl(s)+e - (Pt R )=Ag+Cl - ] 2 Overall: 2AgCl(s)+ H 2 (g)=2ag+ 2H + +2Cl - Cu L Cd(s) CdCl 2 (.1M) AgCl(s) Ag(s) Cu R Anode: Cd=Cd +2 +2e - Cathode: [AgCl+e - =Ag + +Cl - ] 2 Overall: Cd+2AgCl=2Ag+Cd +2 +2Cl -
E cell Nernst Equation E RT ln Q nf E : standard reduction potential n: moles of electrons F: Faraday constant 96485 C/mol
Thermodynamic-Free Energy The maximum cell potential is directly related to the free energy difference between the reactants and the products in the cell. W max G qe max q nf G nfe max
Calculation of Equilibrium Constants for Redox Reactions E cell E when Q RT nf K, E ln Q cell RT nfe E ln K ln K nf RT G -nfe G RT ln K
Reaction Quotient (Q) E G E E -nfe RT nf RT nf ln ln K Q E RT nf ln G nf Q RT nf RT ln nf ln K K Q The positive E (f R >f L ) means that Q<K. As Q increases toward K, the cell emf decreases, reaching zero when Q=K
The Equilibrium Constant of a Cu-Zn Cell Zn+Cu +2 (aq)=zn +2 (aq)+cu E =.34V-(-.76V)=1.1V lnk 2(96485 C mol (8.314 J mol -1-1 K )(1.1J -1 C -1 ) )(298 K) 85.6 K 1.51 37 G 212 KJ/mol
Concentration Cells
Ag L Ag + (.1M) Ag + (1M) Ag R E E E RT nf E ln Q RT nf ln CL 8. 314 298 1. ln. 591 V C 196485 1. R Pt L Cl 2 (P L ) HCl(aq) Cl 2 (P R ) Pt R E RT 2F ln P P L R
Corrosion of Iron
Anodic Region Fe Fe +2 +2e - Cathodic Region O 2 +2H 2 O+4e - 4OH - Overall Reaction 4Fe +2 (aq)+o 2 (g)+(4+2n) H 2 O(l) 2Fe 2 O 3 nh 2 O(s)+8H + (aq)
Electrolysis Electrolytic Cell: use electrical energy to produce chemical change The process of electrolysis involves forcing a current through a cell to produce a chemical change for which the cell potential is negative.
standard galvanic cell standard electrolytic cell Zn+Cu +2 Zn +2 +Cu Zn +2 +Cu Zn+Cu +2
electroplating
Electrochemistry An electrochemical cell produces electricity using a chemical reaction. It consists of two half-cells connected via an external wire with a salt bridge connecting the solutions An electrolytic cell uses an external electricity source to produce a chemical reaction. This is usually called electrolysis. 1
Historically Historically oxidation involved reaction with O 2. i.e., Rusting 4 Fe(s) + 3O 2 (g) Fe 2 O 3 (s) Another example Zn (s) + Cu 2+ (aq) Zn 2+ (aq) + Cu (s) In this reaction: Zn (s) Zn 2+ (aq) Oxidation Cu 2+ (aq) Cu (s) Reduction In a redox reaction, one process can t occur without the other. Oxidation and Reduction reactions must simultaneously occur. 2
Redox Between Zn and Cu If Zn (s) and Cu 2+ (aq) is in the same solution, then the electron is a transferred directly between the Zn and Cu. No useful work is obtained. However if the reactants are separated and the electrons shuttle through an external path... 3
Electrochemical Cells Voltaic / Galvanic Cell Apparatus which produce electricity Electrolytic Cell Apparatus which consumes electricity Consider: Zn Cu Initially there is a flow of e- After a very short time the process stops Electron transport stops because of charge build up on both sides Build up of positive charge Build up of negative charge The charge separation will lead to process where it cost too much energy to transfer electron. 4
Completing the Circuit Electron transfer can occur if the circuit is closed Parts: Two conductors Electrolyte solution Salt Bridge / Porous membrane Process that must happen if e - is to flow. A. e- transport through external circuit B. In the cell, ions must migrate C. Circuit must be closed using a salt bridge (no charge build up) Anode (-) Cathode (+) Green A Red Negative electrode generates electron B C Positive electrode accepts electron Oxidation Occurs Reduction Occurs 5 Anode/Anion (-) Cathode/Cation(+)
Voltaic Cell Electron transfer can occur if the circuit is closed Parts: Two conductors Electrolyte solution Salt Bridge / Porous membrane 3 process must happen if e - is to flow. A. e- transport through external circuit B. In the cell, ions a must migrate C. Circuit must be closed (no charge build up) Anode (-) Black Negative electrode generates electron Oxidation Occur Cathode (+) Red Positive electrode accepts electron Reduction Occur 6 Anode/Anion (-) Cathode/Cation(+)
Completing the Circuit: Salt Bridge In order for electrons to move through an external wire, charge must not build up at any cell. This is done by the salt bridge in which ions migrate to different compartments neutralize any charge build up. 7
Sign Convention of Voltaic Cell @ Anode: Negative Terminal (anions). Source of electron then repels electrons. Oxidation occurs. Zn (s) Zn +2 (aq) + 2e - : Electron source @ Cathode: Positive Terminal (cation) Attracts electron and then consumes electron. Reduction occurs. Electron target: 2e - + Cu +2 (aq) Cu (s) Overall: Zn (s) + Cu +2 (aq) Zn +2 (aq) + Cu (s) E = 1.1 V Note when the reaction is reverse: Cu (s) + Zn +2 (aq) Cu +2 (aq) + Zn (s) Sign of E is also reversed E = -1.1 V Oxidation: Zn (s) Zn +2 (aq) E =.76 V Reduction: Cu +2 (aq) Cu (s) E =.34 V 1.1 V = E CELL or E CELL = E red (Red-cathode) - E red (Oxid-anode) 8
Another Voltaic Cell Zn (s) + 2H + (aq) Zn +2 (aq) + H 2 (g) E =.76 V @ Anode: Negative Terminal (anions): Zn (s) Zn +2 (aq) + 2e - : Source of electron then repels electrons. Oxidation occurs. @ Cathode: Positive Terminal (cation): 2e - + 2H + (aq) H 2 (g) Attracts electron and then consumes electron. Reduction occurs. Overall: Zn (s) + 2H + (aq) Zn 2+ (aq) + H 2 (g) 9
Another Voltaic Cell Zn (s) + 2H + (aq) Zn +2 (aq) + H 2 (g) E =.76 V @ Anode: Negative Terminal (anions): Zn (s) Zn +2 (aq) + 2e - : Source of electron then repels electrons. Oxidation occurs. @ Cathode: Positive Terminal (cation): 2e - + 2H + (aq) H 2 (g) Attracts electron and then consumes electron. Reduction occurs. Overall: Zn (s) + 2H + (aq) Zn 2+ (aq) + H 2 (g) 1
Line Notation Convention basically: metal ionic solution ionic solution metal This is a convenient way of representing cells 1. Anode Cathode [oxidation (-) ] [reduction (+)] 2. represents a phase boundary 3. represents the salt bridge 4. Concentration of component 1 4 Zn (s) ZnSO 4 (aq,1.m) CuSO 4 (aq,1.m) Cu (s) 11 2 3
Line Notation Examples Consider : Zn (s) + Cu +2 (aq) Zn +2 (aq) + Cu (s Anode: Zn Zn +2 + 2e - Cathode: Cu +2 + 2e - Cu Shorthand Line notation Zn (s) Zn +2 (aq) (1.M) Cu+2 (aq) (1.M) Cu(s) 2nd Example : Zn (s) + 2H + (aq) Zn +2 (aq) + H 2(g) Anode: Zn Zn +2 + 2e - Cathode: 2H + + 2e - H 2 (g) Shorthand Line notation Zn (s) Zn +2 (aq) (1.M) H+ (aq) (1.M), H 2 (g, 1atm) Pt (s) 12
EMF - ElectroMotive Force Potential energy of electron is higher at the anode. This is the driving force for the reaction (e- transfer) Larger the gap, the greater the potential (Voltage) Anode e (-) e- flow toward cathode (+) Cathode D P.E. = V = J e - C 13
For electrochemical cells Emf electromotive force ElectroMotive Force Potential energy difference between the two electrodes Units of emf: V - Volts : 1V - 1 Joule / Coulomb 1 Joule of work per coulomb of charge transferred. For electrolysis When a current of one amp flow for 1 second this equals one coulomb of charge passed. Charge (C) = Current (A) x Time (s) 96,4 coulombs = 1 Faraday of charge 1 Faraday is equal to 1 mole of electrons 14
Written as reduction Standard Electrode Potentials Cell Potential is written as a reduction equation. M + + e - M E = std red. potential 15
Written as reduction Zoom of Std. Electrode Potentials Cell Potential is written as a reduction equation. M + + e - M E = F 2 (g) + 2e - 2 F - (aq) 2.87 V Ce 4+ + e - Ce 3+ (aq) 1.61 V 2H + + 2e - H 2 (g). V Li + (aq) + e - Li (s) -3.45 V All reaction written as reduction reaction. But in electrochemistry, there can t be just a reduction reaction. It must be coupled with an oxidation reaction. 16
E Cell Evaluation E Cell Function of the reaction Oxidation Process (Anode reaction) Reduction Process (Cathode reaction) E Cell = E red (cathode) - E red (anode) Simple to remember: Eº (cell) = Eº (red) Eº (ox) E = E redox The potential difference of the cell is equal to the electrode potential of the reduced component minus the electrode potential of the oxidised component. 17
Standard Electrode Potential How is E red (Cathode) and E red (Anode) determine. E (EMF) - State Function; there is no absolute scale Absolute E value can t be measured experimentally The method of establishing a scale is to measure the difference in potential between two half-cells. Consider: Zn Zn +2 + 2e - E =? Can t determine because the reaction must be coupled Question? - How can a scale of reduction potential be determined? Answer - Use a half reaction as reference and assign it a potential of zero. Measure all other half cells in comparison to this reaction. 18
Voltaic Vs. Electrolytic Cells Anode (Oxidation) Voltaic Cell Energy is released from spontaneous redox reaction System does work on load (surroundings) Oxidation Reaction X X + + e- Reduction Reaction e- + Y + Y Overall (Cell) Reaction X + Y + X + + Y, DG = Electrolytic Cell Energy is absorbed to drive nonspontaneous redox reaction Surrounding (power supply) do work on system (cell) Oxidation Reaction A - A + e- Reduction Reaction e- + B + B Overall (Cell) Reaction A - + B + A + B, DG> General characteristics of voltaic and electrolytic cells. A voltaic cell generates energy from a spontaneous reaction (DG<), whereas an electrolytic cell requires energy to drive a nonspontaneous reaction (DG>). In both types of cell, two external circuits provides the means or electrons to flow. Oxidation takes place all the anode, and reduction takes place at the cathode, but the relative electrode changes are opposite in the two cells. 19