Redox Equilibria

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Redox Equations

Redox Reaction : Loss of electrons (Na --> Na+) is Oxidation. Gain of electrons (Cl2 --> Cl-)              Reduction.

Oxidation states go up and down when electrons are lost or gained e.g.:

                                                 V2O5  +  SO2  --->  V2O4  +  SO3

Oxidation state of V:                 +5                    --->  +4                                    Reduction

Oxidation state of S:                                 +4    --->                   +6                   Oxidation

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Electrode Potentials

Electrochemical Cells: Made from 2 different metals dipped in salt solutions of their own ions, it is a redox proccess.

Electrons flow through the wire from the most reactive metal to the least. A voltmeter in the external circuit shows the voltage between the 2 half-cells, this is the cell potential or e.m.f

Platinum electrodes are used when there are 2 solutions involved, this is because platinum is inert and conducts electricity.

ELECTRODE POTENTIALS: reactions occuring at each electrode are reversible. The direction the reaction takes place depends on how easy each metal loses electrons (oxidised). This is measured using electrode potentials. A metal that is easily oxidised has very negative electrode potential, while the one thats harder to oxidise has less negative electrode potential.




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Electrode Potentials

CALCULATING THE CELL POTENTIAL: E   cell  =  (E    right-hand side  -  E    left-hand side)

STANDARD ELECTRODE POTENTIALS: Equilibrium can be affected by changes in.. Temp. , Pressure and conc. To get around this, standard conditions are used to measure the electrode potentials .

STANDARD HYDROGEN ELECTRODE: Measure electrode potential of a half cell against a standard hydrogen electrode. In a standard hydrogen electrode, hydrogen gas is bubbled into a solution of aqueous H+ ions, Platinum electrode is used as a platform for the oxidation/reduction reactions.




STANDARD CONDITIONS : conc. 1.00 moldm-3 , 298 K and 100 kPa

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Electrode Potentials Continued







Electrochemical cell:

Standard hydrogen electrode always on the left

Use the equation to calculate standard electrode potentials. H electrode half cell has electrode potential of 0.00V .

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Electrochemical Series

Electrochemical series is a long list of electrode potentials for different electrochemical half-cells. They are written in order of electrode potential, starting at most negative and going down to most positive.

Half equations are always written as reduction equations, but they are reversable and can go other way. When 2 half equations are put together in the electrochemical cell, more negative electrode potential goes backwards (oxidation) and the more positive goes forward (reduction).

REACTIVITY: Shows what is reactive and what isn't, the more reactive a metal is, the more it wants to lose electrons to form a positive ion. The more reactive metals will have a more negative standard electrode potential.

PREDICTING OUTCOME: More positive (1) one stays forward and the more negative (2) one goes backwards.

So the reactants of 1 reduce the products of 2, to form the reactants of 2.

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Electrochemical Series


Use the formula -->

E   cell  =  (E   bottom  -  E   top)

Just put in the values.

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Electrochemical Cells

Electrochemical cells are used as a commmericial source of electrical energy

NON-RECHARGABLE CELLS: Use irreversiable reactions, common type is a dry alkaline battery. The half-equations have non-reversable arrows because it isn't practicle to reverse them in  a battery, can make it leak or explode. Zinc that forms the casing becomes thinner as zinc is oxidised. Another reason why is cant be recharged is because the ammonium ions produce hydrogen gas, which would escape from the battery.

RECHARGABLE CELLS: Use reversable reactions, found in things like phones. 2 other types include NiCad (nickle-cadmium) and L ion (lithium ion). To recharge these batteries, current is supplied to force the electrons to flow in the opposite direction around the circuitand reverse the reactions.

PROS AND CONS OF NON-RECHARGABLE CELLS: COST; cheaper than rechargable ones short term, in long run more expensive. LIFETIME; non-rechargable work longer than rechargable.  POWER; non-rechargables can't supply as much power as rechargables. TOXICITY; non-rechargables are less toxic than rechargables.

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Electrochemical Cells


Chemicals are stored seperatly outside of the cell and are fed in when electricity is required. An example of this is a hydrogen-oxygen fuel cell, can be used to power vehicles..




Hydrogen and oxygen gases are fed into 2 seperate platinum-containing electrodes. normally made from porous ceramic material with thin layer of platinum (cheaper and larger surface area). Electrodes are seperated by an ion-exchange membrane allows protons to pass through it, stops electrons going through. Hydrogen is fed to negative electrode: H2 ---> 2H+    +   2e-                        The electons flow from the negative electrode through an external circuit to the positive electrode. H+ ions pass through the ion-exchange membrane towards the positive electrode. Oxygen fed to + elextrode:  O2  +  4H+  +  4e-   ---> 2H2O                                                                                      Overall effect is that H2 and O2 react to make water :  2H2  +  O2  --->  2H2O

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Electrochemical Cells


+  They don't need electrical recharging. As long as hydrogen and oxygen are supplied, the cell will continue to produce electricity.

+  Only waste product is water, no nasty waste products to dispose of.

-  Need energy to produce a supply of hydrogen and oxygen. Produced by electrolysis of water. This requires electricity, normally generated by fossil fuels. Not carbon neutral

-  Hydrogen is highly flammable, needs to be handled carefully.

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