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 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

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

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+

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

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

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

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

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

Galvanic Cells – + Zn2+ Cu2+ Flow of electrons zinc copper Salt bridge Zn2+ Cu2+ Negative Electrode (ANODE) Positive Electrode (CATHODE)

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

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

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

Standard Electrode Potentials The Standard Cell Potential Difference (E0 cell) is the measured cell potential difference when the concentration of each species = 1M, pressure = 1 atm and Temp = 25 C E0 cell = E0 oxidant – E0 reductant

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

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

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

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

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

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

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

Primary Cells Dry Cells Alkaline Cells Button Cells

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)

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

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

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 –

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

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

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 –

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

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

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

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

Lead Acid Battery Positive electrode is coated with PbO2 Lead (IV) Oxide The electrolyte is a solution of 4M sulfuric acid

Lead Acid Battery Anode Oxidation (–) Cathode Reduction (+) 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)

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

Nickel Cadmium Cells Anode (Oxidation) (– ) Cathode (Reduction) (+) 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)

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

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

Fuel Cells Anode Oxidation (–) Cathode Reduction (+) Overall Equation H2(g) + 2OH –(aq)  2H2O (l) + 2e– Cathode Reduction (+) O2(g) + 2H2O(l) + 4e–  4OH–(aq) Overall Equation H2(g) + O2(g)  2H2O (l)

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