Using and Controlling Reactions 1. 1. Assign oxidation numbers and balance atom whose oxidation...

ElectrochemistryUsing and Controlling Reactions

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Redox Half Equations

1. Assign oxidation numbers and balance atom whose oxidation number changes

2. Balance oxygen by adding water3. Balance hydrogen by adding H+ 4. Balance charges by adding

electrons (always on the same side as the added H+)

5. Check the equation

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Balancing Redox Equations

1. Multiply one or both equations by appropriate numbers so that the number of electrons lost or gained in each equation is equal

2. Add the two equations cancelling electrons (and other species as necessary)

3. CHECK THE EQUATION!!!!!!!!!!!!!!

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Electrochemistry

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Electrochemicalcells

Galvaniccells

Electrolyticcells

Primarycells

Secondarycells

Fuelcells

Galvanic Cells

Produce electrical energy from spontaneous redox reactions

Consist of two half cells (metal or solution) where the oxidising agent and reducing agent are not in contact with each other.

The two half cells are connected via a conducting wire (connects the electrodes) and the salt bridge (connects the solutions)

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Galvanic Cells

Salt bridge consists of a concentrated solution of a salt which is not easily oxidised or reduced

Oxidation occurs at the ANODE (negative electrode)

Reduction occurs at the CATHODE (positive electrode)

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Galvanic Cells

Electrons flow from anode to cathode through the external wire

Positive ions move from the salt bridge into the reduction half cell

Negative ions move from the salt bridge into the oxidation half cell

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Metal Half Cells Solid metal electrode Solution containing ions of the

same metal (usually a sulfate salt) More reactive metal is oxidised at

the anode: M Mx+ + xe Less reactive metal is reduced at

the cathode: Ny+ + ye N (x and y represent number of electrons gained

or lost by metal/ metal ion)

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Galvanic Cell using Metal Half Cells

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Solution Half Cells

Inert electrodes (Graphite or Platinum)

The reacting solutions may contain an oxidant (e.g. MnO4

–) or a reductant (e.g. I–)

Sulfuric acid is used to acidify solutions in half cell where necessary for a reaction to occur

Electrons are donated or accepted from the solution, not the electrode

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Fuel Cells

Gaseous fuel (most often H2 gas) is oxidised at the anode.

H2(g) 2H+(aq) + 2e

Oxidant (oxygen gas) is reduced at the cathode.

O2(g) + 4H+(aq) + 4e 2H2O(l)

Overall reaction2H2(g) + O2(g) 2H2O(l)

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Fuel Cells

Electrodes: Porous graphite, containing platinum based catalyst. (To increase rate of reaction)

Salt Bridge: Five main types which identifies the fuel cell type. (Alkaline, Solid Polymer (PEM), Phosphoric acid, Molten carbonate, Solid oxide) These allow passage of ions but block the passage of electrons.

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Advantages of Fuel Cells

High operating efficiency Environmentally friendly (don’t

produce SO2, NOx) Quiet and reliable. Will run as long as

the fuel is available and require minimal maintenance.

Better mass to power output compared to conventional galvanic cells

Fuel and oxidant readily available13

Advantages of Fuel Cells

Products are removed as formed, rather than staying inside the cell.

Require minimal maintenance as there are no moving parts.

Can be used for a large range of applications.

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Disadvantages

High purity fuels and oxidants are expensive and are often produced using natural gas as a feedstock.

Impurities in the fuel can “poison” the catalyst in the electrodes

Electrodes are expensive due to the catalyst

Many of the electrolytes are corrosive Rate of reaction is slow. Medium to high

temperatures are required for the cell to function.

Safety and Storage of Hydrogen? 15

Mercedes NECAR Hydrogen Fuel Cell Car

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http://www.cardesignonline.com/technology/necar-fuel-cell.php

Hydrogen Fuel Cell Bicycles

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http://www.alternative-energy-news.info/hydrogen-fuel-cell-bikes

Portable fuel cell powered by water and Aluminium

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http://pinktentacle.com/2006/04/portable-fuel-cell-powered-by-water-and-aluminum/

Sony Exhibiting Hybrid Fuel Cell Batteries in Tokyo

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http://cleantechnica.com/2009/02/26/sony-exhibiting-hybrid-fuel-cell-batteries-in-tokyo/

World's smallest fuel cell promises greener gadgets

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http://www.newscientist.com/article/dn16370-worlds-smallest-fuel-cell-promises-greener-gadgets.html

Similarities between Fuel Cells and Conventional Cells

Redox reactions used to produce direct current.

Electrolyte between electrodes. No pollutants emitted. Anode is negative and cathode is

positive electrode.

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Differences between Fuel Cells and conventional cells

Conventional galvanic cells

Fuel cells

Limited quantities of reactants stored in cell

Continuous external supply of reactants

Must be discarded or recharged when fully discharged

Never discharge or run down

Limited life Virtually unlimited life

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Rechargeable cells

Referred to as storage cells or accumulators

Act as galvanic cells when discharging

During recharging an electric current reforms the original substances

Common types include the lead acid accumulator and the NICAD (nickel cadmium cell)

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Example: Lead Acid Accumulator

Power source in motor vehicles Six lead acid cells connected in

series (generate 2V each) Anode: Lead Cathode: Lead oxide on lead Electrolyte: Sulfuric acid (38%w/v)

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Lead Acid Accumulator

Discharging Anode(-): Pb(s) Pb2+ + 2e Cathode(+): PbO2(s)+ 4H+

(aq)+ 2e Pb2+(aq)+

2H2O(l)

The lead ions react with sulfate ions to form insoluble lead sulfate:

Pb2+(aq) + SO4

2-(aq) PbSO4(s) 25

Lead Acid Accumulator

Overall:PbO2(s)+ Pb(s)+ 2SO4

2-(aq)+ 4H+ 2PbSO4(s)+

2H2O(l)

Anode, cathode and electrolyte are consumed in the reaction

The state of charge/discharge of the battery can be measured by the density of the electrolyte

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Lead Acid Accumulator

Charging: Anode(-) when discharging becomes

the cathode(-) when charging:PbSO4(s) + 2e Pb(s) + SO4

2-(aq)

Cathode(+) when discharging becomes the anode(+) when charging:

PbSO4(s)+ 2H2O(l) PbO2(s)+ 4H+(aq)+ SO4

2-

(aq)+2e 27

Lead Acid Accumulator

Overall: (opposite reaction to discharging)

2PbSO4(s)+2H2O(l) PbO2(s)+ Pb(s)+2SO42-

(aq)+4H+

This regenerates the anode and cathode and increases the density of the electrolyte

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Electrolytic Cells

Change electrical energy into chemical energy

Cause a non spontaneous redox reaction to occur

Electrodes can be reactive or inert Electrolyte is a solution or molten

liquid. The chemicals reactivity related to the reactivity of water determines which is used.

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Electrolytic Cells

Oxidation occurs at the anode (+) and reduction occurs at the cathode (-)

If the electrolyte is molten then the anions (-ve ion) are oxidised at the anode and the cations (+ve ion) are reduced at the cathode.

If the electrolyte is aqueous then the reactions could involve the cations, anions or water. 31

Electrolytic Cells

Reduction Water will be reduced in preference

to the metals in the activity series Al and above:

2H2O + 2e → 2OH- + H2 Zn and below will undergo reduction

in an aqueous solution: M2+ + 2e → M (M represents metal)

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Electrolytic Cells

Oxidation Chloride, bromide and iodide are

oxidised in preference to water: 2X- → X2 + 2e (X represents halogen) Nitrate and sulfate ions will not

oxidise. (N and S already in max oxidation state)

When these ions are present water will oxidise:

H2O → 4H+ + O2 + 4e

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Uses of Electrolytic Cells

Extraction of metals from molten salts

Refining metals Electroplating for protection or

decoration Recharging secondary cells Production of chemicals (NaOH, H2,

Cl2, O2)

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