redox reactions

oxidation

• Gain of oxygen

  • Ex: C + O2 → CO2

• Also loss of hydrogen

  • Ex: CH4 + O2 → CO2 + H2O

reduction

• Loss of oxygen

  • CuO + H2 → Cu + H2O

• Also gain of hydrogen

  • Ex: N2 + H2 → NH3

redox

• Oxidation is loss of electrons

• Reduction is gain of electrons

• Reactions never take place alone, must occur simultaneously

oxidizing agent

• Oxidizes others causing them to lose electrons

• it is reduced, gains electrons

• Oxidation number of oxidizing agent decreases

reducing agent

• Reduces others cousing them to gain electrons

• it is oxidized, loses electrons

• Oxidation number of reducing agent increases

balancing redox reactions

• Write half equation: divide oxidation and reduction into separate equations

• Step 1: balance all atoms except hydrogen & oxygen

• Step 2: balance oxygen by adding water 

• Step 3: balance hydrogen by adding H+

• Step 4: balance charge by adding electrons 

• Step 5: add up & cancel out (electrons must cancel out)

writing equation in neutral medium

• Neutralize H+ by adding OH -

• Combine H and OH -

metals & metal ions

• Metals higher in reactivity can displace less reactive metals

• The more reactive metal acts as a reducing agent

  • The more reactive a metal is the better it is at pushing electrons onto less reactive metal ions

• More reactive atom will have less reactive ion

electrochemical ce

• Devices that interconvert electrical & chemical energy

voltaic cells: primary

converts chemical energy from spontaneous redox reactions to electrical energy, not rechargable

voltaic cells: secondary

• (rechargeable cells) that involve redox reactions that can be reversed using electrical energy

voltaic cells: fuel cells

•  convert chemical energy from a fuel , typically hydrogen, and an oxidant such as oxygen into electrical energy through redox reactions, with water and heat as byproducts

electrolytic cells

•  converts electrical energy to chemical energy by bringing about non spontaneous redox reactions

half cells

• Half cells: a strip of metal in solution of it's salt solution 

• Metalls will reach equilibrium with their cations, releasing electrons on surface of metal

• Strip of metal is an electrode

• The position of the equilibrium determines the potential difference between the metal strip and the solution of metal

• When two half cells are connected, a voltaic cell is formed

electrode

• Each half of a voltaic cell is called a half cell or electrodes

• There are two types of electrones: anodes & cathodes

anode (voltaic)

• Where oxidation occurs (more reactive meta, RA)

• It is negative because oxidation is causing the build up of electrons

• Oxidation is driving the current

• Anode loses mass as the solid metal is converted to it's ions

• Electron flows from anode to cathode

cathode (voltaic)

•  Where reduction occurs (less reactive metal, OA)

• Cathode is positive because reduction is using electrons

• Cathode gains masses as the ion is converted to the metal

voltaic cells

• Electrons travel through the wires that connect the two half cells 

• The more reactive metal is oxidized faster, causing electrons to travel to the less reactive metal 

• This shifts the equilibrium in this less reactive half cell causing it to be reduced

without salt bridge

• The anode is negative

• As oxidation occurs, more ions build up in the solution`this cancels out the charge of the anode, preventing the current from flowing

• The cathode is positive

• The solution contains metl ions & anions, as the metal gets reduced the negative ions build up, canceling out the positive charge

• This prevents current from flowing

salt bridge

• Voltage will only be conducted between electrodes when the circuit is complete

• Complete circuit requires: external circuit connected to each electrode, a salt bridge 

• Salt bridge is a glass tube that contains a concentrated solution of a strong electrolyte

• The ions in the electrolyte can flow between electrodes

• This maintains the negative charge of the anode and the positive charge of the cathode 

voltage

• Different half cells make voltaic cells with different voltages

• Direction of electron flow & voltage generated will be determined by the difference in reducing strength of the two metals 

• This can be judged by the position of metals in the reactivity series 

• Greater the gap between metals, greater voltage produced (more reactive metal = higher voltage)

electrolytic cells

• Electrons are pushed by an outside power source, causing non spontaneous chemical reaction

• Anode is positive because electrons are being pulled away from it by the outside power source

• Electrical energy converted to chemical energy

electrolytic cells vs voltaic cells

• In voltaic cells, the current is the result of a spontaneous redox reaction

• In electrolytic cells, the current is produced by an outside source causing a non spontaneous redox reaction

source of electric power

• A battery or a DC power source

• Represented as 2 lines - long line is positive terminal, short line is negative terminal

• Electric wires are used to connect the electrodes tothe power supply

electrodes of electrolytic cells

• Made of graphite or inert metal which conducts electricut

• Anode is positive, oxidation occurs

• Cathode is negative, reduction oxxurs

• electrons flow from anode to cathode 

• Electric current enters through cathode leaves frim anode

electrolyte

• Aqueous or molten solution of ionic compounds tht will conduct electricity

direction of ions

• Cations in the solution will travel to the cathode, combines with the electron and becomes neutral

• Anion goes to anode, gives up electrons, and becomes neutral

conduction of current

• Current is conducted because: electrons move in the electrical circuit and ions move through the electrolyte to the electrodes

electrolysis of molten salt

• Molten sodium chloride breaks into Na+ and Cl - ions

  • 2NaCl —> 2Na+ + 2Cl-

• At the cathode, Na+ ions take in electrons to become Na atoms, reduction occurs

  • Na+ + e- —> Na

• At anode, Cl- ions give away electrons to become Cl₂, Cl- ions are discharged, oxidation occurs

  • 2Cl - —> Cl2 + 2e-

• When molten ionic compound is electrolized, metal is produce at the cathode, non metal is produced at the anode

secondary cells

• Voltaic cells that can be recharged & reused

• Chemical reactions are reversible

• When it's discharged → foreward reactions occur

• When it's recharged → backward reactions occur

• Example: lead-acid, nickel-cadmium, lithium ion

lead acid

• Used in motor vehicles due to it's ability to produce large electric current required to start an engine

• Cell potential is 2.02 V

• Typically battery has 6 cells in series - total voltage of 12

• Lead is anode and lead (II) oxide is cathode

• Electrolyte is sulfuric acid

lead acid battery reaactions

• At the anode lead is oxidied to form lead (II) sulfate , and at cathode, lead (II) oxide is reduced to formed lead (II) sulfuate 

• When being discharged the following reactions take place:

• Anode: Pb + HSO₄- → PbSO₄ + H⁺ + 2e-

• Cathode: PbO₂ + 3H⁺ + HSO₄- + 2e- → PbSO₄ + 2H₂O

• Overall: Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O

• When it's recharded the reactions are:

• Anode: PbSO₄ + 2H₂O → PbO₂ + 3H⁺ + HSO₄- + 2e-

• Cathode: PbSO₄ + H⁺ +2e- → Pb + HSO₄- 

• Overall: 2PbSO₄ + 2H₂O →   Pb + PbO₂ + 2H₂SO₄

fuel cells

• Need constant source of fuel to produce electric current

• Fuel cells will continue to supply electric current as long as they have fuel to do so

• Fuel is oxidized at anode using a catalyst, typically platinum (pt)

• Electrons flow from anode to cathode, producing current

• At cathode, oxygen gas is reduced, producing pure water

• Ex: hydrogen-oxygen fuel cells, methanol fuel cells

hydrogen-oxygen fuel cells

• Electrolyte is KOH

• Electrodes made of carbon  and incorporated metal catalyst such as Ni or Pt

• At anode, hydrogen gas is oxidized to produce water and electrons which flow through the circuit 

• At cathode, oxygen is reduced to form hydroxide ions

• Hydroxide ions will then migrate through the electrolyte to react with the hydrogen gas at the anode

hydrogen-oxide fuel cells

• Anode: H₂ + 2OH - → 2H₂O + 2e- 

• Cathode: ½ O₂ + H₂O + 2e- → 2OH-

OVERALL: H₂ +  ½ O₂ → H₂O

methanol fuel cells

• Methanol is used as a fuels source

• At the anode, mixture of methanol & water is oxidized to form CO₂ , hydrogen ions, and electrons

• Anode : CH₃OH + H₂O → CO₂ + 6H+ + 6e-

• The hydrogen ions migrate through the polymer electrolyte to the cathode and the electrons flow through the external circuit

• At the cathode, oxygen gas is reduced to form water

Cathode: 6H⁺ + 3/2 O₂ + 6e- → 3H₂O

• Methanol fuel cells do not have the same problems associated with storage of high pressure gases as alkaline fuel cells

• Because it's liquid it's easier to store and transport than compressed hydrogen gas

advantages & disadvantages

	Advantages	Disadvantages
Primary cells	Low cost, convenient, light weight, long shelf life, high energy density at low discharges	Can only be used once, waste batteries leads to large amount of recyling/environmental pollution
Secondary cells	Can be used many times through recharging, high power density, better performance at lower temps	Poorer charge retention, safety issues (Cd very poisonous), high initial cost
Fuel cells	Energy efficient & environmental friends	Expensive to produce, hydrogen difficult to store, need continuous supply of reactants