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Module 5 - Chapter 23 - Redox and Electrode potentials II - Coggle Diagram
Module 5 - Chapter 23 - Redox and Electrode potentials II
Electrode potentials
Electrode with more reactive metal loses electrons and is oxidised - negative electrode
Electrode with less reactive metal gains electrons and is reduced - positive electrode
Tendecy to be reduced and gain electrons
Standard electrode potential
Half cell containing hydrogen gas and solution of H+ ions
Inert platinum electrode is used as they conduct electricity but are inert so won't interfere in the reaction
Conditions
Solutions have concentration of 1 moldm^-3
temperature of 298K
Pressue of 100kPa
EMF of a half-cell connected to a standard hydrogen hald cell under standard conditions
Sign of standard electrode potential shows the sign shows relative tendency to gain electrons compared with the hydrogen half cells.
Standard electrode potential = Standard electrode potential of positive electrode - negative electrode
Measuring standard electrode potential
Two electrodes are connected by wire to allow controlled flow of electrons
Solutions are connected with a salt bridge which allows ions to flow
Salt bridge is made from filter paper soaked in sodium nitrate. It completes the circuit and allows ions to pass between solutions, but it doesn't react with either solution
More negative
- greater tendency to lose electrons and be oxidised. Less tendency to gain electrons and be reduced
More positive
, greater tendency to gain electrons and undergo reduction. Less tendency to lose electrons and undergo oxidation
Metals tend to have negative
values and lose electrons
Non metals tend to have positive
values and gain electrons
More negative
, greater reactivity of a metal in losing electrons
More positive
, greater reactivity of a non metal in gaining electrons
Zinc-copper cell
Copper half cell has more positive electrode potential - greater tendency to be reduced and gain electrons
Zinc half cell has more negative electrode potenital - is oxidised and loses electrons
Electrons flow form negative zinc half cell to less negative copper half cell
Limitations
Reaction rate
Electrode potentials indicate thermodynamic feasibility of a reaction but don't give an indication of the rate
Reaction may seem feasible, but large activation energy prevents it from happening
Concentration
Not all reactios use 1M - this leads to different values electrode potentials
If the reactant concnetration increases ,it shift equilibrium to the right removing electrons and making the system less negative
(-0.76V)
Other factors
Actual conditions for the reaction may not be standard
Standard electrde potentials apply to aqueous equilibria, many reactison take place that aren't aquous
Predicting reactions
Redox systems have oxidising agents on the left and reducing agents on the left
Strongest reducing (most negative) is at the top
Strongest oxidising agent (most positive) is at the bottom
Redox system with more positive standard electrode potential reacts from left to right, gains electrons
Redox system witl the less positive standard electrode potential reacts from right to left, loses electrons
Storage and fuel cells
Primary cells
Non rechargeable
Electrical energy is produced by oxidation and reduction at electrodes - reaction can't be reversed
Chemicals will be used up, voltage falls and battery goes flat
Many are based on the reaction between zinc and manganese dioxide
Secondary cells
Rechargeable
Reaction producing electrical energy can be reversed
Lead-acid batteries used in car batteries
Fuel cells
Uses energy from reaction of a fuel with oxygen to create voltage
Fuel and oxygen are supplied to the cell
Either has acid or alkali electrolyte
Hydrogen fuel cells are most common as water is the only product
Overall equation is combustion of hydrogen
