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CHM 120 CHAPTER 21 Electrochemistry: Chemical Change and Electrical Work Dr. Floyd Beckford Lyon College.

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Presentation on theme: "CHM 120 CHAPTER 21 Electrochemistry: Chemical Change and Electrical Work Dr. Floyd Beckford Lyon College."— Presentation transcript:

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2 CHM 120 CHAPTER 21 Electrochemistry: Chemical Change and Electrical Work Dr. Floyd Beckford Lyon College

3 REVIEW Oxidation: the loss of electrons by a species, leading to an increase in oxidation number of one or more atoms Reduction: the gain of electrons by a species, leading to an decrease in oxidation number of one or more atoms Oxidizing agents: the species that is reduced in a redox reaction

4 Reducing agents: the species that is oxidized in a redox reaction

5 In acidic solution: add H + or H 2 O only In basic solution: add OH - or H 2 O only To balance O To balance H In acidic solution: Add H 2 Oand then Add H + For each O needed For each H needed In basic solution 1. add 2 OH - to the side needing O and and then 1. add 1 H 2 O to the side needing H and 2. add 1 H 2 O to the other side 2. add 1 OH - to the other side

6 THE HALF-REACTION METHOD This method breaks the overall reaction into its two components – half-reactions Each half-reaction is balanced separately and then added Use the following guidelines to help 1. Write as much of the unbalanced net ionic equation as possible 2. Decide which atoms are oxidized and which are reduce – write the two unbalanced half-reactions

7 3. Balance by inspection all atoms in each half- reaction except H and O 4. Use the rules mentioned previously to balance H and O in each half-reaction 5. Make equal the number of electrons involved in both half-reactions Take a look at the breathalyzer reaction H + (aq) + Cr 2 O 7 2- (aq) + C 2 H 5 OH(l)  Cr 3+ (aq) + C 2 H 4 O(l) + H 2 O(l)

8 Balance the following net ionic equation in basic solution.

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10 ELECTROCHEMISTRY Deals with chemical changes produced by an electric current and with the production of electricity by chemical reactions All electrochemical reactions involve transfer of electrons and are redox reactions EChem reactions take place in electrochemical cell (an apparatus that allows a reaction to occur through an external conductor)

11 ELECTROCHEMICAL CELLS Two types: 1. Electrolytic cells: - these are cells in which an external electrical source forces a nonspontaneous reaction to occur 2. Voltaic cells: - also called galvanic cells. In these cells spontaneous chemical reactions generate electrical energy and supply it to an external circuit

12 Electric current enters and exits the cell by electrodes - electrodes are surfaces upon which oxidation or reduction half-reactions occur Inert electrodes: - electrodes that don’t react Two kinds of electrodes: 1. Cathode: - electrode at which reduction occurs (electrons are gained by a species) 2. Anode: - electrode at which oxidation occurs (as electrons are lost by some species)

13 VOLTAIC CELLS Cells in which spontaneous reactions produces electrical energy The two half-cells are separated so that electron transfer occurs through an external circuit Each half-cell contains the oxidized and reduced forms of a species in contact with each other Half-cells linked by a piece of wire and a salt bridge

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15 A salt bridge has three functions: 1. It allows electrical contact between the two half-cells 2. It prevents mixing of the electrode solutions 3. It maintains electrical neutrality in each half-cell as ions flow into and out of the salt bridge Point 2 is important – no current would flow if if both solutions were in the same cell

16 Point 3 is also important – anions flow into the oxidation half-cell to counter the build-up of positive charge Current flow spontaneously from negative to the positive electrode In all voltaic cells the anode is negative and the cathode is positive

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18 In voltaic cells, voltage drops as the reaction proceeds. When voltage = 0, the reaction is at equilibrium

19 The Silver-Copper cell Composed of two half-cells: 1. A strip of copper immersed in 1 M CuSO 4 2. A strip of silver immersed in 1 M AgNO 3 Experimentally we see: : - Initial voltage is 0.46 volts : - The mass of the copper electrode decreases : - The mass of the silver electrode increases : - [Cu 2+ ] increases and [Ag + ] decreases

20 Cu  Cu e - (oxidation, anode) 2(Ag + + e -  Ag)(reduction, cathode) 2Ag + + Cu  Cu 2+ + Ag(Overall cell reaction) Cu |Cu 2+ (1.0 M) ||Ag + (1.0 M) | Ag Notice that in this case the copper electrode is the anode

21 STANDARD ELECTRODE POTENTIALS Associated with each voltaic cell is a potential difference called the cell potential, E cell E measures the spontaneity of the cell’s redox reaction Higher (more positive) cell potentials indicate a greater driving force for the reaction as written All electrode potentials are measured versus the Standard Hydrogen Electrode (SHE): E° = 0.00 V

22 The E° cell calculated is for the cell operating under standard state conditions For electrochemical cell standard conditions are: -solutes at 1 M concentrations - gases at 1 atm partial pressure - solids and liquids in pure form All at some specified temperature, usually 298 K

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24 The electrode potential for each half-reaction is written as a reduction process The more positive the E° value for a half- reaction the greater the tendency for the reaction to proceed as written The more negative the E° value, the more likely is the reverse of the reaction as written

25 Prediction of Spontaneity 1. First write the HR equation with the more positive (less negative) E° for the reduction along with its potential 2. Write the other HR as an oxidation and include its oxidation potential 3. Balance the electron transfer 4. Add the reduction and oxidation HR and add the corresponding electrode potentials to get the overall cell potential, E° cell

26 Important points to note: 1. E° for oxidation half-reactions are equal to but opposite in sign to reduction half-reactions 2. Half-reaction potentials are the same regardless of the species’ stoichiometric coefficient in the balanced equation E° cell > 0 Forward reaction is spontaneous E° cell < 0 Backward reaction is spontaneous

27 E° cell,  G° and K From thermodynamics, we know that,  G° = -RT lnK We can relate E° cell to free energy for that cell  G° = -nFE° cell n = number of moles of e - So-nFE° cell = -RT lnK and E° cell = (RT/nF) lnK

28 (Standard state conditions) Under nonstandard conditions  G = -nFE cell

29 THE NERNST EQUATION Usually concentrations of reactants differ from one another and also change during the course of a reaction As cell reaction proceeds, cell voltage drops so that E° cell is different from E cell E° cell and E cell are related by the Nernst Equation E cell = E° cell - (RT/nF) lnQ

30 E = potential under the nonstandard conditions E° = standard potential R = gas constant, J/mol.K T = absolute temperature n = number of moles of electrons transferred F = faraday, 96,485 J/V.mol e - Q = reaction quotient

31 BATTERIES Two type of batteries: : - Primary batteries cannot be “recharged” Once all the chemicals are consumed there is no more chemical reaction : - Secondary batteries can be regenerated Most common example is the lead storage battery used to power automobiles

32 The Lead Storage Battery Composed of two alternating groups of Pb plates; one group contains pure lead (anode) and the other group contains PbO 2 (cathode) The plates are immersed in 40 % sulfuric acid During discharge Pb  Pb e - (oxidation) Pb 2+ + SO 4 2-  PbSO 4 (precipitation) Net: Pb + SO 4 2-  PbSO 4 + 2e - (anode)

33 At the cathode PbO 2 + 4H + + 2e -  Pb H 2 O(reduction) Pb 2+ + SO 4 2-  PbSO 4 (precipitation) Net reaction: PbO 2 + 4H + + SO e -  PbSO 4 + 2H 2 O Adding the HR for the two half-cells, gives Pb + PbO 2 + 4H + + 2SO 4 2-  2PbSO 4 + 2H 2 O E° cell = V The battery can be recharged

34 Fuel Cells These are galvanic cells in which the reactants are continuously supplied to the cell and the products are continuously removed Best known example is the hydrogen-oxygen fuel cell Hydrogen is fed into the anode compartment and oxygen into the cathode compartment

35 Oxygen is reduced at the cathode – porous carbon doped with metallic catalysts At the anode hydrogen is oxidized to water Anode: 2H 2 (g) + 4OH - (aq)  4H 2 O(l) + 4e - Cathode: O 2 (g) + 2H 2 O(l) + 4e -  4OH - (aq) Overall: 2H 2 (g) + O 2 (g)  2H 2 O(g)

36 CORROSION Ordinary corrosion is a redox process in which metals are oxidized by oxygen in the presence of moisture A point of strain on the surface of the metal acts as an anode Areas on the metal surface exposed to air serves as cathodes

37 Anode: Fe(s)  Fe 2+ (aq) + 2e - Cathode: O 2 (g) + 4H + (aq) + 4e -  2H 2 O(l) 4Fe(s) + O 2 (g) + 4H + (aq)  4Fe 2+ (aq) + 2H 2 O(l) 2Fe 2+ (aq) + 4H 2 O(l)  Fe 2 O 3 H 2 O(s) + 6H + Rust Al also undergo corrosion – initial oxidation is stopped by a layer of Al 2 O 3

38 Corrosion prevention 1. Plating a metal with a thin layer of a less easily oxidized metal 2. Allow a protective film to form naturally on the surface of the metal 3. Galvanizing – coating the metal with zinc 4. Cathodic protection – connecting the metal to a “sacrificial anode”

39 ELECTROLYTIC CELLS Cells in which an electric current causes a nonspontaneous reaction to occur – one common process is called electrolysis In electrolytic cells the anode is the positive electrode and the cathode is the negative electrode Still : Anode = oxidation; cathode = reduction

40 The Down’ Cell: Electrolysis of molten NaCl Using graphite inert electrodes the following observations are made 1. Chlorine, Cl 2, is liberated at one electrode 2. Sodium metal forms at the other electrode Explanation 1. Chlorine is produced at the anode by the oxidation of Cl - ions

41 2. Metallic sodium is formed by reducing Na + ions at the cathode Electrons used at the cathode are reproduced at the anode The reaction is nonspontaneous and electricity is used to force the reaction to occur 2Cl -  Cl 2 (g) + 2e - (oxidation, anode HR) 2(Na + + e -  Na(l) (reduction, cathode HR) 2Na + + 2Cl -  2Na(l) + Cl 2 (g) Overall cell rxn. ________________________________________

42 Electrolysis of aqueous sodium chloride In an EChem cell containing aqueous NaCl : - H 2 gas is liberated at one electrode : - Cl 2 gas is liberated at the other electrode : - Solution at the cathode is basic Rationalization : - Chloride ions are oxidized at the anode and H 2 O is reduced at the cathode

43 2Cl -  Cl 2 + 2e - (oxidation, anode) 2H 2 O + 2e -  2OH - + H 2 (reduction, cathode) 2H 2 O + 2Cl -  2OH - + H 2 + Cl 2 Overall Sodium metal is more active than hydrogen metal and liberates H 2 from solution The hydroxide ions are responsible for the basicity around the cathode

44 FARADAY’S LAW States that the amount of substance that undergoes oxidation or reduction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the cell One faraday = the amount of electricity that reduces or oxidizes 1 equivalent of a substance

45 One equivalent of any substance is the amount of that substance that supplies or consumes one mole of electrons 1F= 1mole of electrons = x e - = 96,485 C C = It C= charge passed; I = current; t = time (in seconds)


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