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Electricity from Chemical Reactions

Electricity from Chemical Reactions. 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. Reduction.

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Electricity from Chemical Reactions

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  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 Zn2+  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  Zn2+ • 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, Cl2 • For a neutral molecule, the sum of the oxidation numbers are zero eg CO2 • 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 (H2O2 & BaO2) it is – 1 • Hydrogen takes + 1 in compounds, except in hydrides (NaH, CaH2) 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) Zn2+(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)  Zn2+(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 Flow of electrons – + zinc copper Salt bridge Zn2+ Cu2+ Negative Electrode (ANODE) Positive Electrode (CATHODE)

  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 (E0cell) is the measured cell potential difference when the concentration of each species = 1M, pressure = 1 atm and Temp = 25 C • E0cell = E0oxidant – E0reductant

  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 wire Glass sleeve H2 gas (1 Atm) Salt Bridge to Other half-cell 1.00M Acid solution Platinum electrode

  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 Mg2+/Mg and Cl2/Cl– half cells get a potential difference of 3.7V • Looking at the electrochemical series

  21. Potential Difference • Cl2 + 2e– Cl– has an E0 of 1.36V • Mg2+ + 2e– Mg has an E0 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 (+) • 2MnO2 (s) + NH4+(aq) + 2e– Mn2O2(s) + 2NH3(aq) + H2O (l)

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

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

  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 (+) • 2MnO2 (s) + H2O(l) + 2e– MnO2(s) + OH –(aq) + H2O (l) • Five times the life of the dry cell

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

  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) + H2O (l)+ 2e – Hg (l)+ 2OH –(aq) • Ag2O(s)+H2O (l)+ 2e – 2Ag (s) + 2OH(aq) • Produce an almost constant 1.35V

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

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

  34. Lead Acid Battery • Car Batteries p 211-2 • 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 PbO2 Lead (IV) Oxide • The electrolyte is a solution of 4M sulfuric acid

  36. Lead Acid Battery • Anode Oxidation (–) • Pb(s) + SO4 2- PbSO4(s) + 2e – • Cathode Reduction (+) • PbO2(s) + SO4 2- + 4H+ + 2e – PbSO4(s) + 2H2O (l) • Overall Reaction • Pb(s) + PbO2(s) + 2H2SO4  2PbSO4(s) +2H2O (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) + H2O (l) + e– Ni(OH)2(s) + OH–(aq) • Overall Reaction • Cd (s) +NiO-OH(s) + H2O(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 (–) • H2(g) + 2OH –(aq) 2H2O (l) + 2e– • Cathode Reduction (+) • O2(g) + 2H2O(l) + 4e–  4OH–(aq) • Overall Equation • H2(g) + O2(g)  2H2O (l)

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

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