Electrochemistry

Question 1

Define a redox reaction.

Answer 1

A chemical reaction where the oxidation states of the atoms change that there is an oxidation and a reduction.

Question 2

Describe a basic hydrogen fuel cell.

Answer 2

There are 2 platinum electrodes connected by wire. H₂ gas is fed into one side at the anode where oxidation occurs to form 2H⁺, the electrons from the reaction pass to the other electrode where O₂ is fed in which is reduced in the presence of H⁺ ions to form water. The chemical energy can be thought to be converted into electrical energy by the current flowing though the wire.

Question 3

Define the anode, cathode and the cell potential.

Answer 3

The anode is where the oxidation occurs (a⇒o vowels), the cathode is where the reduction occurs (c⇒r constanents).

The cell potential, E_cell, is defined as E_cathode - E_anode and must be positive for a spontaneous reaction. Represent this graphically.

Question 4

Give the key equations for Gibbs free energy.

Answer 4

ΔG_cell = -nFE_cell where n is the number of electrons transferred, F is the Faraday constant and E_cell is the charge potential.

ΔG = -RTlnK where K is the equlibrium constant.

Question 5

Define standard electrode potentials.

Answer 5

The measure of the individual potential of a reversible electrode under standard conditions which are: stated temp. usually 298 K, unit activity (simplified to 1 M for soluble species), partial pressur eof 1 bar for each gas and metals in their pure state.

Standard electrode potentials are defined relative to a standard hydrogen electrode which is assigned the E° value of 0 V.

Question 6

Explain how to write out a cell diagram and give an example for the reactio between a Zn and Cu battery where Cu²⁺ is reduced to Cu_(s) and the opposite happens for zinc.

Answer 6

1. The anode is described first, then the cathode, with the reactants specified first and the products after.
2. A single vertical line seperates species in contact but in a different phase, a double vertical line represents a salt bridge seperating the cells.
3. The chemical phase is shown in parentheses as are any conditions that are non-standard.

Zn_(s) | Zn²⁺_(aq) || Cu²⁺_(aq) | Cu_(s)

Question 7

Describe a Latimer diagram.

How can you calculate the cell potential for two of the states?

Answer 7

For one element the oxidation forms are drawn out with the most oxidised species on the left and the most reduced on the right. The states are connected by arrows from left to right that have labelled the numerical value of the standard potential for that reduction.

The standard electrode potentials cannot be added, only the ΔG values can be added. The ΔG values can be worked out by -nFE, so the difference between states is multiplied by the number of electrons in the transition and the Faraday constant. Then the standard electrode potential for the system can be worked out by ΔG/-nF where n is the number of electrons transferred between the states.

Question 8

Describe a Frost diagram and the information that can be gained from it.

Answer 8

A Frost diagram is a plot of oxidation state on the X, and the nE value for converting the state to the zero oxidation state on the Y, where n is the number of electrons transferred. The zero oxidation state will always be the origin.

The Y axis is proportional to ΔG to convert to the elemental form, it represents the relative stability of the oxidation state. The lower the state is, the more stable it is, the zero ox. state is often not the most stable.

The gradient connecting two points is the potential for the reaction, E. Any species that lies above the points either side of it can undergo disproportionation to form both the states. If the gradient is above a point, then the two states around it may conproportionate into that one state.

Question 9

How are Frost diagrams worked out from Latimer diagrams?

Answer 9

Construct a table with the columns: species, oxidation state, reduction potential to the next state (on the Latimer diagram) x number of electrons transferred, nE for converting from oxidation state to the zero state (adding the stepwise potentials down to zero together).

Question 10

Give the Nernst equation. What is it used for?

Answer 10

E = Eº - (RT/nF)lnQ

E is the potential under non-standard conditions, Eº is the standard potential and Q is the reaction quotient which is the products of the activities of the products/the activities of the reactants (with each to the power of their empirical ratio).

Activity of a pure solid/liquid/gas = 1, aqueous liquid = concentration, gas = partial pressure.

This is used to adjust a potential to non-standard conditions.

Question 11

How can the Nernst equation be used with a m-proton, n-electron, reversible process?

Answer 11

Defining the reaction as, X + ne^- + mH⁺ ⇔ XH_m^(m-n)+

The Nernst equation can be substituted with [H⁺] = 10^-pH, the terms seperated, and differentiated to find dE/dpH = - (mRT/nF) ln10

This gives the change in electrode potentials with the change in pH. Such as for a m = 1, n = 1, process, the electrode potential will change with pH by -59 mV pH^-1 at 298 K.

Question 12

How do Frost and Latimer diagrams change when the pH changes?

Answer 12

Using the equation, dE/dpH = - (mRT/nF) ln10, any process which includes a change in protons will depend on pH and temperature. This can drastically change the potentials ofa reaction.

Question 13

What is a Pourbaix diagram?

Answer 13

A plot of electrode potential on the Y and pH on the X, mapping out the stable phases at certain conditions and are often overlapped with water.

At some reducing potentials, protons will form H₂, at some oxidising potentials, H₂O will form O₂ and 2H⁺. Both are m=n processes to change by -59 mV per pH unit.

A similar graph can be drawn but with mole fraction of each oxidation state of an element. The fraction will increase when the pH favours its reaction.

Question 14

Describe how to construct a Pourbaix diagram for the reduction of Zn²⁺ to Zn across a range of pH.

Answer 14

Establish the reactions that occur at the different pH values. At low pH, Zn²⁺ will be in that state in the solution. At pH 9 to 11, Zn(OH)₂ will form and at pH 11 to 12, Zn(OH)₃^- will form.

All these states have different reduction reactions. The potential of Zn²⁺ will not be affected by pH as there are no protons, so up to pH 9, there will be a flat line between the states. Between pH 9 to 11, there is a m=n reaction giving the line a -59 mV pH^-1 gradient and between 11 and 12 there is a 3H, 2e process so the line will be slightly steeper.

Lines are drawn between each state, as the pH changes and the potential where the reduction occurs.

Question 15

How can you predict reactions that will occur from the Pourbaix diagram? What does the Pourbaix diagram assume?

Answer 15

When overlaid with another species, like water, the reactions can be visualised. A metal with a zero oxidation state below the H₂ line on water shows that in water, the metal will react to form H₂. The E_cell for the reaction is the difference between the metal phase line and the H₂ phase line.

The diagram assumes no activation energy.

Question 16

Define overpotential. Describe how this manafests in batteries.

Answer 16

Theoretically, a reaction can be forced to happen if the potential of the backwards reaction is applied. In theory, an additional overpotential is required to cause the reaction to occur. Therefore the energy required to make an oxidation reaction run is E_applied > E_reaction + E_overpotenial and the opposite is true for a reduction.

In batteries, the kinetic barrier to reaction takes away from E_cell.

E_cell = E_cathode - E_anode - E_{overpotential}

Question 17

How can the amount of charge a reaction transfers be worked out?

Answer 17

The moles of electrons transferred x F. Moles of electrons transferred is found from the amount of reactants that complete the reaction which can be found from the equlibrium constant.

Question 18

What is the role of dynamic electrochemistry?

Answer 18

To moniter the rate and energetics of an electrochemical reaction.

Question 19

Draw a diagram of a 3-electrode cell to take dynamic measurements and define each of the components.

Answer 19

Counter electrode: Generates current to counter that flowing at the working electrode.

Working electrode: Where you moniter the redox reaction of interest.

Referemce electrode: Provides a stable electrode voltage.

For the working electrode, negative current is defined as electrons moving out of the working electrode surface with positive current as the opposite.

Question 20

Define positive and negative faradaic current and non-faradaic current.

Answer 20

Positive faradaic current indicates an oxidation reaction is occuring. (pO =Ox)

Negative faradaic current indicates a reduction reaction is occuring. (nE = rE)

Current not due to a redox reactions is described as non-faradaic.

Question 21

Give the use and an example of reference electrodes.

Answer 21

Standard electrode potentials are defined relative to the standard hydrogen electrode but having to use H₂ gas is very inconvinient so other electrodes are used with very stable redox potentials. Most common is the a AgCl/Ag cell with a AgCl coated Ag wire in a 3 M Cl^- solution. This has the potential of 0.207 V vs SHE.

Question 22

What is the role of the counter electrode in a 3-electrode cell?

What electrolyte is used in the cell?

Answer 22

This is added to stop current flowing through the reference electrode and changing the redox potential used as the standard.

The only requirement is a system with high ionic conductivity. In aqueous solution NaCl is a common choice. In organic solvents a common salt is tetrabutylammonium hexafluorophosphate.

Question 23

Define chronoamperometry and give the control and measured variables.

Answer 23

Chonoamperometry measures the current (amperometry) as a function of time, t_hold, (chrono) with a constant voltage, E_app. The current is measured at the working electrode. Charge transferred can be determined from the intergral of the current-time graph.

Question 24

Define a voltammetry experiment and the graphs that are drawn from it.

What types of voltammetry are there?

Answer 24

Voltage is changed as a function of time while current is monitered. A voltammogram is a plot of current vs voltage.

Linear sweep and cyclic.

Question 25

Describe the two types of voltammetry.

Answer 25

Linear sweep: A linear voltage sweep is applied with either a positive (anodic) or negative (cathodic) gradient. A single sweep is made over a voltage range with a constant scan rate.

Cyclic voltammetry: A forward and backward voltage sweeps are applied with both a positive and negative gradient. Sweeps are also

Question 26

Describe the source of non-faradaic current in a cell.

Answer 26

When the voltage of the working electrode changes, the ions in solution that are attracted to it.

Decreasing the voltage of the working electrode will ‘charge’ the surface with cations due to the increasing attraction between them. This results in electrons moving away from the electrode.

Scan rate controls the rate of ion movement so the current depends on the scan rate.

Question 27

Define and describe film electrochemistry.

Answer 27

Where the analyte is adsorbed to the working electrode.

The film can be oxidised in the reaction and then be reduced by the potential applied to the working electrode. The film is quite often a catalyst for another reaction occuring in solution.

Question 28

For electrocatalytic film voltammetry, what shape is the graph predicted to be and how is this prediction made?

Quantatively describe the shape.

Answer 28

The non-faradaic component will form two lines slightly apart, which follow the shape of the faradaic current.

The Nernst equation can be used where the reaction quotient is represented by the amount of catalyst and the conversion between the two catalytic species. The conversion between the species depends on the voltage applied. Simplification leads to a sigmoidal relation between the species. (pg 45)

In the two extreme potentials, the coversion of the film species is at its maximum conversion rate or not being converting at all so there is no current. In between there is a dramatic change in potential is the ratio of the redox species is highly dependant on the potential.

Question 29

What equation describes the rate of catalysis of film electrochemistry?

Answer 29

i_cat = nFm_catk_cat

Where i_cat is the current as a function of potential of catalysed process, m_cat is the molar amount of a catalyst species as a function of potential and k_cat is the rate of transfer between catalyst species.

Question 30

Describe the shape of the faradaic only current of a non-catalytic, reversible redox reaction in a cyclic voltammetry experiment and how you can interpret the figure.

Answer 30

At high and low potentials, the species will remain in one state and between the two, there will be a curve where the electrons are transferred.

The potential at which both the reactions occur is the same and equal to the potential of the redox couple (For Cu^2+/1+, the potential is the E°.

A peak area of the current-time graph is equal to the charge passed which therefore gives a measure of the amount of redox active material at the electrode surface (peak area (C)/F (C per mol of electrons) x moles of electrons per species). The scan rate can be used to convert the X-axis.

Question 31

How do peaks heights change when the scan rate of voltammetry experiments change?

How do the faradaic peak potentials for catalytic species and reversible redox reaction species compare?

Answer 31

The current peaks increase as the scan rate increases due to the kinetic time delay of the redox processes.

The faradaic responses of the are much smaller when there is no catalytic regeneration.

Question 32

Why is there a seperation of E_ox and E_red at fast scan rates in voltammetry experiments?

Answer 32

The kinetic delay for the time at the which the potential is needed to induce electron transfer. The time delay is represented by the apparent voltage seperation.

Question 33

What causes the causes the ‘duck shaped’ voltammagram seen in solution cyclic voltammetry?

How does the shape change with with scan rate and concentration of the redox species?

Answer 33

The competition between the rate of electrolysis at the electrode surface and the rate of transport of the reacting chemical species to the surface by diffusion.

At the lower potentials the transfer between the species is slow enough that the diffusion doesn’t affect the rate of transfer. At the peak current the reaction is limited by the rate of diffusion of the species at the electrode.

Increasing the scan rate and the contration each increase the peak heights in the voltammogram. This is described in the Randles-Sevcik equation.

Question 34

Describe what can be interpreted from the currents of a cyclic solution voltammetry experiment of a reversible redox process.

Answer 34

The individual peak voltages and currents can be defined for the reduction and oxidation processes individually. The peak currents should be opposite but the same magnitude.

The current depends on scan rate by the Randles-Sevcik equation

i_p = 0.4463nFAC • (nFνD/RT)^-0.5

Where A = electrode area, C = concentration of redox species.

The key part of this is that the peak current is proportional to the square root of the scan rate. This is because faster scan rates cause the diffusion layer to be established over a shorter distance with a higher concentration gradient.

Question 35

Describe what can be interpreted from the redox potentials of a cyclic solution voltammetry experiment of a reversible redox process.

Answer 35

The average of the oxidative and reductive potentials for the process is called the midpoint.

The number of electrons transferred in the reaction, n, is calculated by:

E_p^ox - E_p^red = 2.22 (RT/nF)

However this relies on a low resistance solution and isn’t perfect.

Question 36

How can the rate determining step of a redox reaction be found from a voltammogram?

Answer 36

By comparing the peak potentials and working out the predicted seperation by E_p^ox - E_p^red = 2.2RT/nF

If the seperation is larger then there must be some chemical RDS.

Question 37

Describe when a multiple peak cyclic voltammogram would be seen.

Answer 37

When successive redox reactions are occuring, and that there is a point before full reduction/oxidation where the species is stable.

Question 38

What is the appearence of a irreversible redox reaction and when would it occur?

Answer 38

Like the normal shape but with only one faradaic peak. This occurs when one species can react in the conditions to form a new species that cannot undergo the reverse process.