ELECTROCHEMISTRY

• Relationship between electricity and chemical reactions

Two types of electrochemical processes: 

• Generation of electric current from spontaneous chemical reactions
  • Electrochemical Cells
  • Exothermic
• Use of electricity to cause non-spontaneous reactions to occur
  • Electrolysis
  • Endothermic

Oxidation:

• Loss of electrons
• Gain of Oxygen
• Loss of Hydrogen
• Increase in Oxidation number

Reduction:

•  Gain of electrons
• Loss of Oxygen
• Gain of Hydrogen
• Decrease in Oxidation number

• Reactions between a metal & a solution containing ions of a different metal
• Reaction will only occur if a reactive metal is place in a less reactive solution
  • use activity series to determine if reactions will occur

Process:

• Metal dissolves and the ions of the other metal are reduced to elemental state and deposits out of the solution
• 

• Elemental State = 0
• Monatomic Ions = charge on ion
• The sum of oxidation states of atoms in a neautral molecule/ ionic compound= 0
• The sum of oxidation states of atoms in a polyatomic ion = charge on ion
• Group 1= +1
• Group 2 = +2
• Oxygen = -2
  • except peroxides = -1
  • and F20 = +2
• Hydrogen = +1
  • except metal hydrides = -1

• The Zinc anode becomes smaller as ions go into solution
• Electrons flow from anode to cathode
• Copper ions in solution become copper atoms which deposit on *****
• Cations move from anode to cathode
• Anions move from cathode to anode
• Anode is oxidation
• Cathode is reduction
• Anode (-)
  • electrons are generated by oxidation
• Cathode (+)
  • electrons are accpeted in reduction

Salt Bridge

• filter paper containing an electrolytic solution ( KNO3 or NaNO3)
• it completes the cicuit and allows ions to move between each half-cell

• Galvanic cells are a source of direct current & provide portable electrcity

Classified into two groups:

• Primary Cells
  • cannot be recharged
    • Lelanche Cell
    • Alkaline Dry Cell
• 
• Secondary Cells
  • can be recharged- supply electrical energy to reverse the cell reaction
    • Lead-acid Battery

• Dry cells have electrolytes in the form of pastes
• Lelanche Cell

Lelanche Cell-

• The outer zinc casing is the anode
• The graphite rod is surrounded by a paste containing manganese dioxide which is the cathode
• Ammonium Chloride and Zinc Chloride are in the paste and act as the electrolyte (salt bridge)

Advantage: used in many applications

Disadvantages:

• short shelf life (zinc and ammonium in paste react and cause leaking and deterioration in the cell)
• Rapid drawing of current from the cell causes ammonia to buld up and this causes a drop in voltage

• Alkaline Dry Cell
  • Composition is similar to Lelanche cell-
  • except powered zinc anode is used 
  • electrolyte is 7mol/L of KOH

Advantages:

• has a longer working life
• supplies current more rapidly

Disadvantages:

• is more expensive though

Lead-Acid Batteries are commonly used in motor vehicles

• Electrodes in each cell consist of a bank of lead grids supporting a large surface area of the electrode material
• Negative electrode (anode) grid is filled with spongy metallic lead
• Positive electrode (cathode) grid is filled with brown lead oxide
• Eletrolyte is sulfuric acid

Advantages:

•  Reachargeable & Cheap

Disadvantages:

• Enclosed in heavy casing-limits use in application
• Lead can do body damage

• process where an electric current is used to drive a non spontaneous chemical reaction
• electrical energy is converted to chemical energy
• products have higher potential energy than reactants
• current comes from power pack/battery
• a cell which electrolysis occurs is called an electrolytic cell
• electrons travel through external circuit

Uses:

• electroplating-improve appearance, reduce corrosin
• electrorefining-higher levels of purity
• extraction of reactive metals from their ores
• manufacturing of aluminium, sodium hydroxide, chlorine, hydrogen

Electrodes:

• electrodes may/ may not take part in reaction
  • Inert electrodes do not take part in reaction
• chemical reactions occur at interface between electrode and electrolyte
• anode (+)
• cathode (-)
• battery pulls electrons from anode and pushes them to cathode

Electrolyte:

• produces ions in the solution
• ions may/may not take part in reaction
  • is water is present it is preferentially oxidised
  • or a reactive metal cathode (Zinc Chloride)
• Cations to Anions

Factors:

• the nature of the element
• the concentration of its ions in solution
• the temperature of the solution

Amount of metal does not influence e.m.f. Temperature & concentration must be stated when comparing electrode potentials of different elements. Standard conditions are 298 K (25 degrees) and 1mol/dm3

•  Increasing concentration increases the electrode potential
• Temperature and concentration affect the voltage

For Molten Salts

• cations are reduced at cathode
• anion is oxidised at anode

For Aqueous Solutions-at anode

• Nitrate & Sulfate are never oxidised
• Water is oxidised to produce Oxygen
• Anions may be oxidised (depends on water)
• Chloride is oxidised if concentrated
• If a non-inert metal is used as electrode; it may oxidise & go into solution as ions

For Aqueous Solutions- at cathode

• Metal ions of less reactive metal (Cu& Ag) are reduced in preference to water
• Water is reduced in preference to metal ions of reactive metals (Na, Mg, Al)

Anode Reaction:

• More positive e.m.f value is oxidised (of the two competing)
• Water is oxidised producing Oxygen gas
• 

Cathode Reaction:

• Ions with more positive e.m.f value is reduced (of the two competing)
• Water is reduced producing Hydrogen gas

• Pure water does not conduct electricity-therefore KNO3 or Dilute Sulfuric acid is added.
• Hydrogen gas is produced at the cathode
  • H ions are discharged-accepting electrons to form hydrogen molecules
• Oxygen gas is produced at the anode
  • OH ions are discharged-form water and oxygen molecules