14. Redox II
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Define what is meant by a standard electrode potential of a half cell
The standard electrode potential of a half cell is the emf (potential difference) of the half cell connected to the standard hydrogen electrode under standard conditions - 298K, 100KPa and 1moldm-3 Symbol is Eө
Describe the standard hydrogen electrode
An inert platinum electrode is in contact with hydrogen gas and H+ ions —> In a diagram, the standard hydrogen electrode always goes on the left
What do you need to create an electrochemical cell
Salt bridge - usually made of filter paper soaked in concentrated solution of KNO3 High resistance voltmeter - this minimises the flow of electrons
Why is a salt bridge used in electrochemical cells
Allows ions to flow between the half cells to complete the circuit
Why might E cell you have in you experiment not match E cell standard (core 10)
–> Metal oxidation - ensure metal is sanded before starting experiment to remove oxide layer —> Not all concentrations are 1 moldm-3. May have to use less like 0.4 moldm-3 as it may be not safe to use 1moldm-3 —> Temperatures may not be standard- use thermometer to monitor temp as it may not be standard temp (273K)
What does a more positive standard electrode potential value indicate
More negative standard electrode potential value means the more positive you can be that the half cell will be a reduction More positive electrode potential means the species is more likely to gain electrons and be reduced, so equilibrium sits on the right
Explain half cell notation - how do you draw combined half cells
More positive standard electrode ie REDUCTION goes on the right Less positive electrode (or S.H.E) ie OXIDATION goes on the left ROOR acronym may help The reduced forms of each cell are on the outside: Eg: OXIDATION || REDUCTION Zn(s)| Zn2+ (aq) || Cu2+ (aq) | Cu(s) —> you will also need to include Pt right at the start/ end too for some electrodes: Fe(s) l Fe2+(aq) ll VO2+(aq), H+(aq), V3+(aq) l Pt(s) –> l represents a change in physical state and a comma is used to separate species in the same physical state
Give the 3 equations that can be used to calculate Ecell
Right hand electrode - left hand electrode Reduction - oxidation Positive - negative —> pick which to use depending on the wording of the question. Some questions specify which the reduction is, even if it is more negative
Describe the “flipping method”
—> decide which half cell is the reduction and look up the Ecell value —> Decide which half cell is oxidation and look up the Ecell value BUT reverse the sign —> Add the reduction value to the reversed value for oxidation to get Ecell
For a reaction to be feasible, E cell must be….
Positive
Give the equation that relates gibs free energy with Ecell
Gibs free energy = -nFEcell Gibs free energy = - number of moles of electrons being transferred x faraday constant x Ecell Units for gibs is Jmol-1 —> will need to divide by 1000
Give the different ways half cells can be made
Metal or non metal half cells eg: zinc rod electrode in Zn2+ solution Different oxidation states of the same element eg: a platinum electrode in a solution containing both Fe3+ and Fe2+ ions
Where is the oxidising agent found in the half equation
On the left - agent on the left is more easily reduced so it makes the oxidising agent
Where is the reducing agent found in the half equation
On the right - agent on the right is more easily oxidised so it makes the reducing agent Reducing agent - Right. Both R’s
Explain the difference between primary storage cells and secondary storage cells
Primary cells are non rechargeable as one of the reactants will eventually run out. The reaction is not reversible (irreversible) Secondary cells are rechargeable and the reaction is reversible
Give examples of primary storage cells and secondary storage cells
Primary - zinc - carbon battery, alkaline battery Secondary - NiCd, Nickle-metal-hydride cell, lithium ion battery
In storage cells, which is the charge and discharge reaction
Forward reaction - discharge Backward reaction - charge The half equations that occur during the charging are the reverse direction of the half equations that occur when discharging
Explain how a hydrogen-oxygen fuel cells generate electricity Suggest why a fuel cell does not need to be recharged?
The fuel (hydrogen) is oxidized at the anode (looses electrons) These electrons get transferred to the oxygen gas molecules from air to reduce oxygen to o2- ions at the cathode This creates a potential difference within the cell Fuel cell is constantly supplied with fuel
Give the equations for a hydrogen - oxygen fuel cells under alkaline conditions
Anode = : H2 (g) + 2OH– (aq) → 2H2O (l) + 2e– (NOT found in data book) Cathode= O2(g) + 2H2O(l) + 4e- –> 4OH-(aq) (found in data book)
Give the equations for a hydrogen - oxygen fuel cells under acidic conditions
Anode= H2(g) –> 2H+(aq) + 2e- (found in the data book) Cathode= O2(g) + 4H+(aq) + 4e- —> 2H2O(l) (found in data book)
What is the overall equation that is the same for both alkaline and acidic conditions
2H2(g) + O2(g) —> 2H2O(l) –> The fact that they have the same overall equation explains why they have the same E cell value
Give the advantages and disadvantages of hydrogen - oxygen fuel cells
Advantages: —> water is the only byproduct - hydrocarbons produce CO and CO2 which need to be removed by catalytic converters —> fuel cells are more effienct - 40-60% efficient but engines are only 20% Disadvantages: —> Hydrogen is a flammable gas so is difficult to store —> limited infrastructure —> energy from fossil fuels is often still needed as hydrogen is made by the electrolysis of water
Explain what happens if the concentration of the oxidised species in the half cell increases
Makes E more positive, as fewer electrons are produced as equllibrium shifts to the right
Give some reasons why a reaction may not occur, even if the E cell is positive
High activation energy Rate of reaction was too slow