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USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL: TO...

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USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL: TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING AGENTS. TO CALCULATE THE STANDARD CELL POTENTIALS. TO PREDICT POSSIBLE REACTIONS.
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Page 1: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

USES OF STANDARD ELECTRODE POTENTIALS

THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:

TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING AGENTS.

TO CALCULATE THE STANDARD CELL POTENTIALS.

TO PREDICT POSSIBLE REACTIONS.

Page 2: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING AGENTS

• THE STANDARD ELECTRODE POTENTIALS ARE ARRANGED IN ORDER OD ITS VALUES TO GIVE THE ELECTROCHECMICAL SERIES.

• THIS GIVES A MEASURE OF THE STRENGTHS OF SPECIES (REALTIVE OXIDIZING OR REDUCING).

Table showing examples of standard electrode potentials.

• THE MORE NEGATIVE (-VE) VALUE INDICATES A STRONGER REDUCING AGENT.

• THE MORE POSITIVE(+VE)VALUE INDICATES A STRONGER OXISIZING AGENT.

R E/V

K+ + e- K -2.92 strongest

2H+ +2e- H2 0.00 reducing

Cl2(g) + 2e- 2Cl- 1.36 agent.

Page 3: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

TO CALCULATE THE STANDARD CELL POTENTIALS.

THE STANDARD CELL POTENTIAL IS THE MAXIMUM POTENTIAL DIFFERENCE BETWEEN TWO HALF CELLS CONNECTED UNDER STANDARD CONDITIONS.

THE STANDARD CELL POTENTIAL CAN BE CALCULATED BY USING THE STANDARD ELECTRODE POTENTIALS OF THE HALF CELL.

THE STANDARD POTENTIAL WITH THE MOST NEGATIVE VALUE HAS TO BE REVERSED .

• EQUATION FOR STANDARD ELECTRODE POTENIAL IS: E(CELL)= E(ANODE)+E(CATHODE).

Page 4: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

QUESTIONCALCULATE THE E (cell) VALUE

IF THE STANDARD ELECTRODE POTENTIAL OF CHLORINE IS 1.36V AND POTASSIUM IS -2.92 V .

ANSWER

Page 5: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

ANSWER.• SINCE POTASSIUM HAS THE MORE NEGATIVE

ELECTRODE POTENTIAL THE SIGN IS REVERSED.

E(CELL)=E(ANODE)+E(CATHODE)

E(CELL)=+2.92V + 1.36V

E(CELL)=+4.28V.

Page 6: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

TO PREDICT POSSIBLE REACTIONS.

• A REACTION WILL ONLY OCCUR IF:

• THE E.M.F OF THE REACTION IS POSITIVE ,THE REACTION IS SAID TO BE FEASIBLE.

• THE REACTION WILL NOT OCCUR IF:

• THE EMF OF THE REACTION IS NEGATIVE.

Page 7: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

TO DETERMINE IF A REACTION IS FEASIBLE OR NOT:

1. DETERMINE THE STANDARD ELECTRODE POTENTIAL OF THE 2 HALF CELLS.

2. THE VALUES ABOVE TO CALCULATE THE STANDARD CELL POTENTIAL .

IT IS IMPORTANT TO NOTE THAT THE STANDARD CELL

POTENTIAL INDICATES THE FEASIBILITY OF THE REACTION .

Page 8: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

EFFECT OF CONCENTRATION ON ELECTRODE POTENTIAL.

• THE ELECTRODE POTENTIAL IS MEASURED UNDER STANDARD CONDITIONS, i.e MAINTAINED CONCENTRATION ,TEMPERATURE AND PRESSURE.

• WHEN THERE IS A CHANGE IN THESE VALUES IT AFFECTS THE ELECTRODE POTENTIAL,TEMPERATURE AND PRESSURE AND BE KEPT CONSTANT BUT CONCENTRATION CANNOT.

Page 9: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

USING LE CHARTERLIE’S PRINCIPLE.

• WHEN A REACTION OCCURS THE AMOUNT OF A SUBSTANCE USED IS NOT THE AMOUNT PRODUCED,HENCE THERE IS A CONCENTRATION CHANGE.

• IF THERE IS A DECREASE IN THE CONC IN THE CATHODE ,TAKE IN TO CONSIDERATION LE CHARTERLIE’S PRINCIPLE.

Page 10: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

LE CHARTERLIE’S PRINCIPLE.

LE CHARTERLIE’S PRINCIPLE STATES THAT IF A SYSTEM IN EQUILIBRIUM IS DISTURBED BY CHANGES IN THE DETERMING FACTOR ,e.g. CONCENTRATION, THE SYSTEM WILL TEND TO SHIFT ITS EQULIBRIUM POSITION SO AS TO COUNTERACT THE EFFECT OF THE DISTURBANCE.

Page 11: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

ENERGY STORAGE DEVICES

• ENERGY STORAGE DEVICES CONVERT CHEMICAL ENERGY INTO ELECTRICAL, TWO SUCH EXAMPLES ARE BATTERIES AND CELLS.

• A BATTERY IS MANY CELLS JOINED IN SERIES OR PARALLEL.

• PRIMARY CELLS ARE BATTERIES THAT PRODUCE EMF FROM IRREVERSIBLE CHEMICAL REACTIONS.

• SECONDARY CELLS ARE BATTERIES WHICH PRODUCE EMF FROM REVERSIBLE CHEMICAL REACTIONS.

Page 12: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THERE ARE FIVE(5) TYPES OF CELLS /BATTERIES:

1.THE DANIEll CELL

2.THE DRY LECLANCHE CELL. 3.ALKALINE CELL. 4.LEAD –ACID ACCUMULATOR. 5.FEUL CELL.

Page 13: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THE DANIELL CELL.

• THIS IS THE FIRST CELL TO BE USED.

• IT WAS INVENTED BY JOHN FREDERIC DANIELL.

• IT CONSIST OF Cu POT FILLED WITH CuSO4 • A POROUS POT IS IMMERSED IN THE COPPER POT ABOVE.

• THE POROUS POT COINTAINED A ZINC ANODE WHICH TO THE CENTRE

• THE EMF UNDER IS 1.10VOLTS.

Page 14: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THE DANIELL CELL.

Page 15: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THE DRY LECLANCHE CELL.

• THE WAS INVENTED IN 1866 BY GEORGES LECLANCHE.

• POROUS POT CONSIST OF A CARBON ROD DIPPED INTO A MANGANESE AND CARBON POWDER PASTE.

• THE POROUS POT IS SURROUNDED BY ANOTHER PASTE OF ZINC CHLORIDE AND AMMONIUM CHLORIDE DISSOLVED IN WATER.

• THE ABOVE IS ENCLOSED IN A ZINC CASE WHICH ACTS AS THE ANODE.

Page 16: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THE DRY CELL.

Page 17: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

REACTIONS THAT OCCUR WHEN THE DRY CELL IS OPERATING.

• WHEN THE CELL IS OPERATING :

1. THE OUTER ZINC CASE OXIDISES ,THIS IS SEEN IN THE FOLLOWING EQUATION Zn(s) Zn2+ (aq) + 2e.

2. AT THE CATHODE, THE EQUATION BELOW SHOWS THE REACTION. 2MnO2(s) +2H+(aq) + 2e Mn2O(s) +H20. • THE ELECTROLYTE OF THE CELL LEAK OUT AS THE ZINC CONTAINER GETS THINNER .

• DUE TO THE AMMONIUM CHLORIDE IN THE CELL THE ZINC CASING WARES AWAY EVEN WHEN NOT IN USE.

• THE CELL HAS AN EMF OF 1.5 VOLTS AND CAN LAST UP TO 1.5 YEARS.

Page 18: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

ALKALINE BATTERY

• THESE ARE SIMILAR TO LECLANCHE DRY CELLS.

THE DRY CELL USES POTASSIUM HYDROXIDE INSTEAD OF AMMONIUM CHLORIDE AS THE ELECTROLYTE.

• THE ANODE IS MADE UP OF ZINC POWDER WHICH INCREASES THE RATE OF REACTION. • THE CATHODE IS MADE UP OF MANGANESE DIOXIDE. • THE HALF EQUATIONS FOR THE REACTION ARE AS FOLLOWS:

Zn(s) + 2OH-(aq) Zn(OH)2(aq) + 2e AND 2MnO2(s) +H2O+2e Mn2O3(S) + 2OH-(aq)

NOTE: BATTERIES SHOULD BE REMOVED FROM DEVICES BEFORE STORAGE.

Page 19: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

ALKALINE BATTERIES.

Page 20: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THE LEAD ACID ACCUMULATOR.

• IT WAS INVENTED IN 1859 BY GASTON PLANTE.

• THEY ARE THE OLDEST TYPE OF RECHARGEABLE BATTERIES.

• THIS CONSIST OR 2-6 CELLS (CONNECTED IN SERIES),GENERATING 6/12 VOLTS .

• THE ABOVE PROVIDES ENOUGH ELECTRICITY FOR STARTING AND IN ENGINE IN A VEHICLE

• THE CELL CONSIST OF A LEAD ANODE AND A LEAD ANODE AND A LEAD (IV) OXIDE IN CONCENTRATED H2SO4.

Page 21: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

LEAD-ACID ACCUMULATOR.

Page 22: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

REACTIONS THAT OCCUR WHEN THE LEAD-ACID ACCUMULATOR IS OPERATING.

• WHEN THE CELL IS OPERATING:

1. THE ANODE IS OXIDISED FROM Pb TO Pb2+ IONS.THE EQUATION FOR THIS REACTION IS : PB(s) Pb2+ (aq) + 2e.

2. AT THE CATHODE THE PbO2 REACTS WITH THE H+ IONS IN THE H2SO4 TO FORM Pb2+ IONS.THE EQUATION FOR THSI REACTION IS : PbO2(s) +4H+(aq) +2e PB2+(aq) +2H2O(l).

3. THE Pb2+ IONS THAT IS FORMED REACTS WITH THE SO4 IONS IN THE ACID AND LEAD SULPHATE IS FORMED.THE EQUATION FOR THIS REACTION IS: Pb2+(aq) +SO4(aq) PbSO4(s).

4. THE ALTERNATOR RECHARGES THE BATTERY BY PASSING AND ELECTRIC CURRENT THUS RESTORING THE BATTERY TO ITS ORIGINAL CONDITION.

NOTE:IF THERE IS A BUILD UP OF PbSO4 WHICH BECOMES COARSER AND INERT THEN THIS PREVENTS THE BATTERY FROM RETURNING TO ITS ORIGINAL CONDITION BY RECHARGING.

Page 23: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THE FUEL CELL.

• THIS IS A DEVICE THAT CONVERTS THE CHEMICAL ENERGY FROM A FUEL INTO ELECTRICITY THROUGH A CHEMICAL REACTION.

• IT IS A PRIMARY CELL , IT IS DIFFERENT IN THE SENSE THAT THE ELECTRICITY SOURCE CAN BE REPLACED CONSTANTLY.

• THE ELECTRODES ARE RELATIVELY INERT AND ONLY CATALYZE THE CELL REACTIONS.

• PRESENTLY THE FUEL CELL IN USE IS THE HYDROGEN-OXYGEN FUEL CELL.

• THE H-O FUEL CELL IS HIGHLY EFFICIENT AND POLLUTION FREE.

• THEY ARE USED IN SPACE CRAFTS FOR HEAT, ELECTRICITY AND TO PROVIDE DRINKING WATER FOR ASTRONAUTS.

Page 24: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THE FUEL CELL.

IT CONSISTS OF WARM POTASSIUM HYDROXIDE SOLUTION WHICH IS HELD BETWEEN POROUS CARBON ELECTRODES.

THE ELECTRODES ARE COATED WITH A CATALYST E.G PLATINUM OR NICKEL THIS INCREASES THE RATE OF REACTION.

Page 25: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

THE FUEL CELL.

Page 26: USES OF STANDARD ELECTRODE POTENTIALS THERE ARE THREE(3) USES OF STANDARD ELECTRODE POTENTIAL:  TO MEASURE THE RELATIVE STRENGTHS OF OXIDIZING AND REDUCING.

REACTIONS OCCURING WHEN THE CELL IS OPERATING .

1.HYDROGEN ENTERS INTO THE –VE COMPARTMENT OF THE CELL AND DIFFUSES THROUGH THE CARBON ANODE.

2. THE CATALYST SEPARATES THE GAS INTO IONS AND ELECTRONS. 3. THE EQUATION FOR THE REACTION AT THE ANODE IS: H2(g) +2OH(aq) 2H2O(l) +2e.4. OXYGEN ENTERS INTO THE +VE COMPARTMENT OF THE CELL AND DIFFUSES INTO

THE POTASSIUM HYDROXIDE.

5. OXYGEN ,WATER AND ELECTRONS COMBINE CATALYTICALLY TO FORM HYDROXIDE IONS ,THIS CAN BE SEEN IN THE EQUATION BELOW.

O2(g) +2H2O(l) +4e 4OH(aq). THE OVERALL EQUATION FOR THE HYDROGEN OXYGEN CELL IS : 2H2(g) +O2(g) 2H2O(l).


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