Topic 14 Redox II

What is meant by standard electrode potential?

• the voltage/potential difference of a cell when a half-cell is connected to a standard hydrogen electrode measured under standard conditions

Conditions of standard electrode potential:

• 298K temperature

• 100kPa pressure of gases if gases are involved

• 1.0 moldm^-3 concentration of ions

Features of a standard hydrogen electrode

• more complex and less reliable as it is difficult to maintain the pressure of gas at 100kPa (as a gas is involved)

Why is a reference electrode necessary?

• contains a standard electrode potential of 0V under standard conditions

• used to measure the standard electrode potential of another half cell

Calculate the emf by combining two standard electrode potentials

Write cell diagrams using conventional representation of half-cells

Importance of conditions when measuring the electrode potential, E

• temperature of 298K

• If using a gas, 100kPa

• Solutions of 1.0moldm^-3

Predict the thermodynamic feasibility of a reaction using standard electrode potentials

In order for a reaction to be thermodynamically feasible, the Ecell value must be positive. This does not mean that the reaction will happen, as kinetic factors, such as a high Ea, could cause the rate of reaction to be effectively 0.

Why can’t HCl be used to acidify potassium permanganate?

What is the electrode potential of a cell directly proportional to?

• the total entropy change for a reaction

• lnK for a reaction

(therefore a positive E cell will have an overall positive entropy change - spontaneous)

What are the limitations of predictions made using standard electrode potentials

• kinetic inhibition

• departure from standard conditions

While a reaction may be thermodynamically feasible it may occur at such a slow rate that it doesn’t actually occur.

Limitations are non-standard conditions or high activation energy.

What can standard electrode potentials be listed as?

• an electrochemical series

How does disproportionation relate to standard electrode potentials?

Disproportionation definition => an element in a single species is simultaneously oxidised and reduced

What are different methods used to measure standard electrode potentials of?

• metals or non-metals in contact with their ions in aqueous solution

• ions of the same element with different oxidation numbers

What is the energy released on the reaction of a fuel with oxygen utilised in?

• a fuel cell to generate a voltage

• Methanol and other hydrogen-rich fuels

Electrode reactions that occur in a hydrogen-oxygen fuel cell:

Carry out both structured and non-structured titration calculations including Fe²⁺/MnO₄^-:

Methods used in redox titrations (MnO₄^- ):

• redox titration uses the same techniques as an acid-base titration but is applied to the reactants in a redox reaction

MnO₄^- ion acts an as oxidising agent:

• acts as an oxidising agent in acidic solution

• The mole ratio is 5 mol Fe²⁺: 1 mol MnO₄^-

• potassium manganate (VII) acts as its own indicators during titration with aqueous iron (II) ions

• dilute H₂SO₄ must also be added to the conical flask to supply the H⁺ ions

• purple MnO₄^- (aq) ions are reduced to colourless Mn²⁺ (aq) ions

• pale green Fe²⁺ (aq) ions are oxidised to pale yellow Fe³⁺ (aq) ions

• the end-point of the reaction occurs when a hint of pink is observed in the flask showing a slight (one drop) excess of MnO₄^- (aq) when all Fe²⁺ (aq) has reacted

Half equations and overall equation of MnO₄^-/Fe²⁺:

Carry out both structured and non-structured titration calculations including I₂/S₂O₃^2-:

Methods used in redox titrations (S₂O₃^2-):

Sodium thiosulphate can act as a reducing agent:

• mole ratio is therefore I mole of I₂: 2 mole S₂O₃^2-

• none of the ions resent in the iodine / thiosulphate reaction are coloured

• starch, which forms a deep blue complex with iodine, can be used as an indicator to show a clear end point for the reaction

• a small quantity of starch solution is added just before the end point when the iodine solution is a pale straw colour (i.e. when most of the iodine has reacted)

• the end point is when the intense blue colour of the starch-iodine complex finally disappears

Half equations and overall equation of I₂/S₂O₃^2-

Using redox titrations to work out the oxidation state of an element - Example 1:

Using redox titrations to work out the oxidation state of an element - Example 2:

Using redox titrations to work out the oxidation state of an element - Example 3:

Using redox titrations to work out the oxidation state of an element - Example 4:

Using redox titrations to work out the oxidation state of an element - Example 5:

What must a half-cell contain?

Both the oxidised and reduced species from a half-equation

Half-cell with metal in contact with a solution that contains metal ions:

What is an electrochemical cell?

2 half cells joined together

How do you set up an electrochemical cell?

• use sandpaper or emery paper to rub off any metal oxides formed on surface of metal during oxidation of the metal which acts as an insulator

• place each metal in a solution of ions of the same metal for example using a zinc electrode placed in solution of zinc nitrate

• place a salt-bridge (filter paper soaked in saturated potassium nitrate) across both beakers - has to be submerged in solution

• allows ions to move and carry charge to complete circuit (electrons can’t move across this)

Why use a high resistance voltmeter in an electrochemical cell?

• prevents current from flowing

• if replaced by lightbulb, the current will flow, and the bulb will light

• the light bulb will go out eventually when one of the reagents is used up

• as voltmeter measures the voltage between two half cells (Ecell)

Why is platinum used as an electrode when 2 aqueous ions are involved?

• metallic so conduct electricity but are also inert so won’t interfere with reaction

Explain what happens at a Zn²⁺/Zn half cell:

• although it looks as though nothing happens in a half-cell, some of the metal atoms get oxidised forming ions, leaving the electrons on the piece of metal

• the reversible arrow shows that the reaction happens in both directions

Half equation of previous reaction:

Explain what happens at a Cu²⁺/Cu half-cell:

• as Cu is less readily oxidised compared to Zn, there’d be less of a build-up of electrons on the piece of Cu (s) metal

Half equation of previous reaction:

Details relating to a Zinc/Copper voltaic cell

• when the two half-cells are connected they form a voltaic cell

• a voltaic cell is an example of an electrochemical cell

• a wire allows electrons to flow from one half-cell to the other

• a salt bridge completes the circuit by allowing the flow of ions

What happens at the Zn electrode?

• the Zn electrode is the negative electrode, as there is a greater build-up of electrons on the Zn electrode compared to the Cu electrode

Why is the fact that the voltmeter is high resistance important?

• the voltmeter is a high resistance voltmeter which prevents much current from flowing

• this enables us to measure the pd (or the cell emf) when two-half cells are combined

Explanation of this voltaic cell

• no reaction/ very little reaction takes place while the high-resistance voltmeter is in the circuit

• if it’s replaced by a bulb, or if a wire joins the two electrode together, then electrons will flow freely from left to right

• once this happens, the size of the Zn electrode will get smaller due to Zn→ Zn²⁺ + 2e^- (oxidation) and the size of the Cu electrode will get bigger due to the Cu²⁺ + 2e^- → Cu (reduction).

• the salt bridge completes the circuit

• it is usually made from a strip of filter paper soaked in saturate aqueous KNO₃

• in the LH half-cell, zinc metal is oxidised to Zn²⁺ ions

• the Zn²⁺ ions enter the solution, two NO₃^- ions from the salt bridge enter the solution to maintain neutrality

• in the RH half-cell, Cu²⁺ ions are reduced to Cu metal

• the [Cu²⁺] in the solution therefore decreases, and to maintain neutrality for each Cu²⁺ ion that is reduced, two K⁺ ions enter the solution from the salt bridge

Write down the key features of a half cell where the metal is in contact with the solution containing its metal ions (using Cu and Zn as an example)

Write down the key features of a half-cell which involves a gas, using H2 as an example

Write down the key features of a half-cell which involves a gas, using Cl2 as an example

Write down the key features of a half-cell involving both the aqueous ions of the same element, using Fe²⁺ and Fe³⁺ as an example

(when you add two equal volumes of 2 solutions, it halves the concentration of each ions)

Write down the key features of a half-cell involving both the aqueous ions of the same element, using V²⁺ and V³⁺ as an example

Write down the key features of a half-cell when it contains an oxidising agent (using MnO₄^- as an example)

For oxidising agents, it is important that the solution is acidic

Write down the key features of a half-cell when it contains an oxidising agent (using Cr₂O₇^2- as an example)

For oxidising agents, it is important that the solution is acidic

Define standard electrode potential

the potential difference of a cell when a half-cell is connected to the standard electrode under standard conditions

What is the value of the electrode potential affected by?

• temperature

• pressure of gases

• concentration of the solution

By convention, what side does the standard hydrogen electrode go on?

the left hand side

How to measure the potential difference between the standard hydrogen electrode and the Zn²⁺ (aq)/ Zn(s) half cell

• set up voltmeter as shown

• Zn is a better reducing agent than H₂, so Zn is oxidised and electrons flow around the circuit from right to left

• E_cell = E_RHS - E_LHS = -0.76V

How to measure the potential difference between the standard hydrogen electrode and the Cu²⁺ (aq)/ Cu(s) half cell

• set up voltmeter as shown

• H₂ is a better reducing agent than Cu, so H₂ is oxidised and the electrons flow around the circuit from left to right

• Cu²⁺ (aq) + 2e^- → Cu (s)

• E_cell = E_RHS - E_LHS = +0.34V

A question regarding disproportionation and E_cell:

Name the oxidation states and their associated colours of Vanadium:

Show the reaction that occurs when iodide ions are used to reduce vanadium +5

Show the reaction that occurs when tin is used to reduce vanadium +5

Show the reaction that occurs when zinc is used to reduce vanadium +5

What must be true if a reaction is thermodynamically feasible?

• the total entropy must be positive

• gibbs free energy must be negative

• Ecell must be positive

What is true about the relationship between the Ecell of a redox reaction and the total entropy change and lnK for a reaction?

they are directly proportional

What must you think about if non-standard conditions are used?

le Chatelier’s principle - equilibrium

What happens if you increase the concentration of Cu²⁺?

This would cause the equilibrium to shift to the RHS so E_cell for the forward reaction would become more positive.

Any change that would result in the position of equilibrium shifting to the RHS will result in the value of the E_cell becoming more positive.

What is the effect of changing the temperature on the value of E_cell?

The majority of voltaic cells are exothermic in the spontaneous direction. You would have to be told the value of delta H or whether the reaction was exothermic to justify changing the temperature.

Assuming the reaction is exothermic, an increase in temperature would cause the position of equilibrium to shift in the backward direction, so the value of E_cell, for a spontaneous reaction, would be less positive.

An increase in [Cl^-] would cause the position of equilibrium to shift to the RHS and so the value of E_cell will become more positive / less negative. If the value of E_cell>0V, then the reaction becomes thermodynamically feasible, and that’s why MnO₂ can oxidise conc HCl to Cl₂ gas.

Commercial cells (batteries)

• electrochemical cells can be used as a commerical source of electrical energy. Important of cell include non-rechargeable, rechargeable and fuel cells.

• the great advantage of this is that it is a source of portable electricity.

• note that a battery is more than once cell joined together (e.g. a car battery is made from six cells joined together) - what the general public call a battery, chemists call cells.

Non-rechargeable cells

• in these cells, the chemicals are used up over time and the E_cell drops. Once one or more of the chemicals have been completely used up, the cell is flat and the E_cell is 0 volts

• note the lack of the standard symbol sign. Conditions in these cells are not standard.

• these cells cannot be recharged and have to be disposed of after their single use

• E_cell of batteries is sometimes referred to as emf (electromotive force)

• you are not required to learn any of these half-equations, just to be able to use them

What are fuel cells?

• continuous supply of the chemicals into the cell and so neither run out of chemicals nor need re-charging

• use energy from the reaction of a fuel with oxygen to create a voltage

• most common fuel cell is a hydrogen-oxygen fuel cell

Hydrogen-oxygen fuel cell in acidic conditions:

Pros and cons of using cells:

Ethanol fuel cells:

Methanol fuel cells: