Presentation is loading. Please wait.

Presentation is loading. Please wait.

Electricity from Chemical Reactions. Electrochemistry The production of electrical energy from chemical reactions Redox reactions involve the transfer.

Similar presentations

Presentation on theme: "Electricity from Chemical Reactions. Electrochemistry The production of electrical energy from chemical reactions Redox reactions involve the transfer."— Presentation transcript:

1 Electricity from Chemical Reactions

2 Electrochemistry The production of electrical energy from chemical reactions Redox reactions involve the transfer of electrons Redox means that reduction and oxidation are occurring simultaneously

3 Reduction Occurs when there is a decrease in oxidation number Zn 2+  Zn Gains electrons Loses Oxygen Converting a complex substance into a simpler form i.e. smelting iron to produce the pure metal iron

4 Oxidation Occurs where there is an increase in oxidation number Zn  Zn 2+ Loses electrons Gains oxygen The reaction used to describe the reaction of any substance with oxygen

5 Determining Oxidation Numbers The atoms in elements have an Oxidation Number of zero eg Fe, C, Cl 2 For a neutral molecule, the sum of the oxidation numbers are zero eg CO 2 For a monatomic ion, the oxidation number is the same as it’s charge Cl –, Na +

6 Determining Oxidation Numbers Oxygen usually takes – 2 in compounds. In peroxides (H 2 O 2 & BaO 2 ) it is – 1 Hydrogen takes + 1 in compounds, except in hydrides (NaH, CaH 2 ) where it takes – 1

7 Determining Oxidation Numbers For a polyatomic ion, the sum of the oxidation numbers of its component atoms is the same as its charges For polyatomic molecules or ions, the, most electronegative element has a negative oxidation number and the least electronegative element has a positive oxidation number

8 Redox Half Reactions Consider the reaction when a strip of zinc is dropped in a solution of Copper Sulphate Zn (s) + Cu 2+ (aq)  Zn 2+ (aq) + Cu (s) Electrons are transferred from zinc atoms to copper ions Reaction occurs spontaneously, that is with no external force or energy being applied

9 Redox Half Reactions Redox reactions consist of two half reactions Oxidation Zn (s)  Zn 2+ (aq) + 2e –1 Reduction Cu 2+ (aq) + 2e –1  Cu (s) It is possible to use redox reactions to produce electricity

10 Galvanic Cells Also called Electrochemical Cells Achieved by separating the half equations into half cells Transferred electrons are forced to pass through an external circuit Such an apparatus is called a Galvanic Cell

11 Galvanic Cells Cu 2+ Zn 2+ zinc copper –+ Salt bridge Negative Electrode (ANODE) Positive Electrode (CATHODE) Flow of electrons

12 Standard Electrode Potentials The electrical potential of a galvanic cell is the ability of the cell to produce an electric current. Electrical potential is measured in volts Cannot measure the electrode potential of an isolated half cell Can measure the difference in in potential between two connected half cells

13 Standard Electrode Potentials Electrical potential of a cell results from competition between 2 half cells for electrons Half cell with the greatest tendency to attract electrons will undergo REDUCTION Other half cell will lose electrons and undergo OXIDATION

14 Standard Electrode Potentials The Reduction Potential of a half cell is a measure of the tendency of the oxidant to accept electrons and so undergo reduction The difference between the reduction potentials of the two half cells is called the Cell Potential Difference

15 Standard Electrode Potentials The Standard Cell Potential Difference (E 0 cell ) is the measured cell potential difference when the concentration of each species = 1M, pressure = 1 atm and Temp = 25 C E 0 cell = E 0 oxidant – E 0 reductant

16 Standard Electrode Potentials A Standard Hydrogen Half cell is used as a comparative measure the reduction potentials of other cells The SHE is given a value of 0.00 V All other half cells are given a reduction potential value in comparison to this SHE by being connected to it

17 Standard Hydrogen Electrode Platinum electrode H 2 gas (1 Atm) Salt Bridge to Other half-cell 1.00M Acid solution Platinum wire Glass sleeve

18 Standard Hydrogen Electrode SHE is used to measure reduction potential of other cells If a species accepts electrons more readily than hydrogen, its electrode potential is positive If a species accepts electrons less readily than hydrogen, its electrode potential is negative

19 Electrochemical Series The reaction that is higher on the electrochemical series will occur as it appears and will reverse the direction of the reaction that occurs lower on the table

20 Potential Difference Is measured by a volt meter Can be estimated by using electrochemical series Connect Mg 2+ /Mg and Cl 2 /Cl – half cells get a potential difference of 3.7V Looking at the electrochemical series

21 Potential Difference Cl 2 + 2e –  Cl – has an E 0 of 1.36V Mg e –  Mg has an E 0 of – 2.38V The potential difference can be calculated 1.36 – (– 2.38) = 3.74V

22 Galvanic Cells Primary Cells –Produce energy until one component is used up, then discarded Secondary Cells –Store energy and may be recharged

23 Primary Cells Dry Cells Alkaline Cells Button Cells

24 Dry Cells The ordinary zinc – carbon cell Anode oxidation (–) –Zn (s)  Zn 2+ (aq) + 2e – Cathode oxidation (+) – 2MnO 2 (s) + NH 4 + (aq) + 2e –  Mn 2 O 2 (s) + 2NH 3 (aq) + H 2 O (l)

25 Dry Cells The new cell produces about 1.5V Once reaction reaches equilibrium its “flat”

26 Dry Cell Metal Cap (+) Anode Zinc Case (–) Cathode Carbon Rod Mixture of Carbon & Manganese Dioxide Ammonium Chloride & Zinc Chloride Electrolyte

27 Alkaline Cells The ordinary zinc – carbon cell Anode oxidation (–) –Zn (s)  Zn 2+ (aq) + 2e – Immediately reacts with OH – ions in the electrolyte to form zinc hydroxide –Zn (s) + 2OH – (aq)  Zn(OH) 2 (s) + 2e –

28 Alkaline Cells Cathode reduction (+) – 2MnO 2 (s) + H 2 O (l) + 2e –  MnO 2 (s) + OH – (aq) + H 2 O (l) Five times the life of the dry cell

29 Alkaline Cell Potassium Hydroxide Electrolyte Metal Cap (+) Metal Base (–) Cathode outer steel case Powdered Zinc Anode Steel or Brass Mixture of Carbon & Manganese Dioxide

30 Button Cells Used in very small applications like watches, cameras etc. Two main types Mercury zinc and silver zinc Anode Oxidation (–) –Zn (s) + 2OH – (aq)  Zn(OH) 2 (s) + 2e –

31 Button Cells Cathode Reduction (+) depends on the type of battery –HgO (s) + H 2 O (l) + 2e –  Hg (l) + 2OH – (aq) –Ag 2 O (s) +H 2 O (l) + 2e –  2Ag (s) + 2OH (aq) Produce an almost constant 1.35V

32 Button Cell Cathode outer container of nickel or steel (+) Metal Cap (–) Electrolyte Zinc Powder Mercury Oxide

33 Secondary Cells Lead – Acid (Car Battery) Nickel cadmium Cells Fuel Cells

34 Lead Acid Battery Car Batteries p Also called storage batteries or accumulators Each cell produces 2 volts so typical 12 volt car battery contains 6 cells Both electrodes are lead plates separated by some porous material like cardboard

35 Lead Acid Battery Positive electrode is coated with PbO 2 Lead (IV) Oxide The electrolyte is a solution of 4M sulfuric acid

36 Lead Acid Battery Anode Oxidation (–) Pb (s) + SO 4 2-  PbSO 4 (s) + 2e – Cathode Reduction (+) PbO 2(s) + SO H + + 2e –  PbSO 4 (s) + 2H 2 O (l) Overall Reaction Pb (s) + PbO 2(s) + 2H 2 SO 4  2PbSO 4 (s) +2H 2 O (l)

37 Nickel Cadmium Cells Often called Nicads Electrodes are Nickel and Cadmium Electrolyte is Potassium Hydroxide Reactions involve the hydroxides of the two metals

38 Nickel Cadmium Cells Anode (Oxidation) (– ) Cd (s) + 2OH – (aq)  Cd(OH) 2 (s) + 2 e – Cathode (Reduction) (+) NiO-OH (s) + H 2 O (l) + e –  Ni(OH) 2 (s) + OH – (aq) Overall Reaction –Cd (s) +NiO-OH (s) + H 2 O(l)  Cd(OH) 2 (s) + Ni(OH) 2 (s)

39 Fuel Cells Limitation of dry cells looked at so far is that they contain reactants in small amounts and when they reach equilibrium. Primary Cells are then discarded, secondary cells are then recharged A cell that can be continually fed reactants would overcome this and allow for a continual supply of electricity

40 Fuel Cells Fuel cells transform chemical energy directly into electrical energy 60% efficiency Space Program uses hydrogen and oxygen with an electrolyte of Potassium Hydroxide

41 Fuel Cells Anode Oxidation (–) –H 2 (g) + 2OH – (aq)  2H 2 O (l) + 2e – Cathode Reduction (+) –O 2(g) + 2H 2 O (l) + 4e –  4OH – (aq) Overall Equation –H 2 (g) + O 2 (g)  2H 2 O (l)

42 Hydrogen Oxygen Fuel Cell – + Water outlet Oxygen Gas Inlet Hydrogen Gas Inlet Electrolyte Porous Anode Porous Cathode

Download ppt "Electricity from Chemical Reactions. Electrochemistry The production of electrical energy from chemical reactions Redox reactions involve the transfer."

Similar presentations

Ads by Google