3.1.11.1 electrode potentials + cells

deifne an oxidising agent + reducing agent

outline what an half-cell is

• when a piece of metal is dipped into a solution of its metal ions, an equilibrium is set up that represents a half equation

• this is known as a half cell

for example, describe what happens for a zinc half cell (a zinc rod immersed in a solution containing Zn²⁺ ions (eg a solution ZnSO₄)

give the half equations for this half cell

• the Zn atoms on the rod can deposit two electrons on the rod + move into solution as Zn²⁺ ions

• Zn → Zn²⁺ + 2e⁻

• alternatively, the Zn²⁺ ions in solution could accept two electrons from the rod + move onto the rod to become Zn atoms

• Zn²⁺ + 2e⁻ → Zn

• this can be written as an equilibrium reaction

• Zn²⁺ + 2e⁻ ⇌ Zn

if the equilibrium lies to the left hand side, what does this mean?

that there are more electrons on the surface of the rod

for example, describe what happens for a copper half cell (a copper rod immersed in a solution containing Cu²⁺ ions (eg a solution Zn(NO₃)₂)

give the half equation for this half cell

• this can also be written as an equilibrium reaction

• Cu²⁺ + 2e⁻ ⇌ Cu

• in this half-cell the equilibrium could lie further to the right hand side so fewer electrons would be deposited on the rod

what is the electrode potential?

what the potential difference set up between the rod + solution is known as

what does measuring the electrode potential tell us?

how readily electrons are released by the metal so how good a reducing agent the metal is

outline what will happen if two different electrodes are connected

• the potential difference between the two electrodes will cause a current to flow between them

• an electromotive force (EMF) is established + the system can generate electrical energy

draw a representation of a cell + outline how one is created

• each half-cell is placed in a separate beaker

• the two beakers are then connected by a salt bridge + a voltmeter

(negative electrode is always on left side of drawing)

give 4 features of a cell

1. voltmeter

2. wire

3. electrodes

4. salt bridge

what is the purpose of the voltmeter?

measures the potential pushing power of electrons through the circuit but keeps the current at zero

what is the purpose of the wire

allows the movement of electrons

what are the electrodes?

• these are where the half equations are taking place

• and are so referred to as the half cells

what is the salt bridge?

its usually a filter paper soaked in a solution of KNO₃

what is the purpose of the salt bridge?

to complete the circuit + allows the movement of ions

why is KNO₃ a suitable solution?

• it does not react with any of the ions in solution so it doesnt interfere with the redox reaction

• and its ions can move

why is KCl not used as a salt bridge solution for a cell that contains Ag⁺ ions?

it contains Cl⁻ ions which would react with Ag⁺ to form a white ppt

there are other types of half cells, outline these half cells

• there is no solid metal involved in the half-equation

• for these half cells, a metal electrode is required + usually platinum is used as it is so unreactive (an inert electrode) + conducts electricity

outline the 3 general types of electrode

what is the effect of surface area of the electrode on the EMF?

the surface area of the electrode will not affect the EMF of the cell

how can a typical electrochemical cell be made?

by combining a Zn/Zn²⁺ electrode and a Cu/Cu²⁺ electrode

outline what the positive + negative electrodes are

• the positive electrode is the one which most favours reduction (to the right)

• the negative electrode is the one which most favours oxidation (to the left)

in terms of example of the copper, zinc electrochemical cell, state which is the positive + negative electrode, which way the electrons flow + give half equations

• copper electrode is positive + zinc electrode is negative

• electrons flow from the zinc electrode to the copper electrode

it is not possible to measure…. it is only possible to measure the…

• the absolute potential of a half electrode on its own

• the potential difference between two electrodes

how can we measure the potential difference between two electrodes?

the electrode has to be connected to another half-cell of known potential + the potential difference between the two half cells is measured

what will measuring the potential difference help us workout?

• which electrode will be positive or negative

• + therefore its tendency to release electrons

by convention, what can we assign to each electrode + how?

• can assign a relative potential to each electrode

• by linking it to a standard electrode

what is the primary standard used?

the Standard Hydrogen Electrode (SHE)

what potential is the SHE assigned + why?

its assigned the potential of 0 volts, by definition, as a reference

draw a labelled diagram of the SHE

outline

a. how the SHE works

b. the half equation

c. standard conditions

d. components

a. hydrogen gas is bubbled into a solution of H+ ions and since hydrogen doesnt conduct, a platinum electrode is used

• metal electrode

for the substances given, draw a labelled diagram of the electrode

• redox electrode

for the substances given, draw a labelled diagram of the electrode

• gas electrode

what happens once the electrode potentials have been measured using the SHE?

the electrode potentials are displayed in a table called the electrochemical series

what is the electrochemical series?

a list of electrode potentials for different half cells in order of decreasing (or increasing) potentials

outline how redox is demonstrated in the electrochemical series

• all electrochemical series display reduction reactions only

• but all reactions are in equilibrium so can be reversed to become oxidation equations

which ones are the best reducing agents on the electrochemical series? outline why

• these have very negative potentials which mean that they are good at giving away electrons

• it is the species on the right of these equations → are the best reducing agents

which ones are the best oxidisng agents on the electrochemical series? outline why

• these have very positive potentials

• which means that they are very good at attracting electrons

• it is the specifies on the left of these equations - are the best oxidising agents

what can we use the values of electrode potentials to calculate?

to help calculate the overall electromotive force (EMF) of a cell + to predict whether reactions will take place

give the strongest reducing agent + explain why

what is it possible to use electrode potentials data from the electrochemical series to do?

to decide whether a redox reaction (and electron transfer) between two half-cells will work, as not all combinations will

those combinations that do react are called what?

feasible reactions

consider the reaction of Zn (s) and Cu²⁺ (aq)

explain which half equation is more likely to show oxidation

combine the half equations to give the overall equation

• the Zn²⁺/Zn half equation has a lower E⍬ value (-0.76) than the Cu²⁺/Cu half equation (+0.34) so will become the oxidation half equation

• this means the half equation must be reversed to show oxidation

• the two new half equations + overall equation is shown in the image

comment on the feasibility of the reaction between Zn(s) and Cu²⁺(aq)

explain why the reaction below are feasible + give an overall equation:

Co and Ag⁺

explain why the reaction below are feasible + give an overall equation:

Cl⁻ and MnO₄⁻

what is the final oxidation product of the reaction between Ag⁺ and Fe(s)?

in this question, check + see if the product of oxidation can be further oxidised by the oxidising agent

what is the final oxidation product of the reaction between Cu²⁺ and Fe(s)?

when will a reaction not be feasible?

when you construct the overall redox equation, if the suggested species are not reactants (ie on the same side of the equation)

eg Mn²⁺ and Zn²⁺, explain why they do not react together?

summary:

outline what happens at the negative and positive electrodes

negative electrode:

• oxidation

• equilibrium lies left

positive electrode:

• reduction

• equilibrium lies right

what is the E⍬ cell?

• also known as the e.m.f (electromotive force)

• its the potential difference across the two electrodes

outline how to calculate emf (E⍬ cell)

• first part is to choose two half equations + decide what the overall (feasible) redox reaction is when they combine

• then we consider the E⍬ values long with a straight forward calculation:

eg give the overall equation for the reaction between Br₂ and Cu and calculate the emf (E⍬ cell)

Br₂ + Cu → 2Br⁻ + Cu²⁺

how can we use the value of emf to predict the feasibility of a reaction?

• a positive value means the reaction is feasible

• a negative value means the reaction is not feasible

• the more positive a value, the more feasible a reaction

for reaction between Co and Ag⁺

for reaction between Cl- and MnO₄⁻

two different explanations (red and blue)

what is the conventional cell diagram?

a pictorial representation of the two electrodes that make up an electrochemical cell

outline what the single lines represent

outline what we use commas for in conventional cell diagrams

what is the double line used to represent?

the salt bridge

explain how to write a conventional cell diagram

• species on the LHS of diagram is species being oxidised (giving the electrons) → the negative electrode

• species on the RHS of diagram is species being reduced (accepting the electrons) → the positive electrode

• H⁺, H₂O are usually not included (unless theres a change in ox state) and e- from half equation are never included in the diagram

• Pt electrode is represented as |Pt in cell diagram and they go on the outside of diagram

• the order of elements is based on the direction of their half equation

what is used in a cell where there is no solid to act as the electrode? why?

• a platinum electrode will be used

• it provides a conducting surface for electron transfer

• Pt is used as it us unreactive + can conduct electricity

by convention, on what side is the SHE always written?

the left-hand side

SHE always goes on the LHS when writing conventional cell diagrams but what else is different as well?

the oxidation is still written on the left and reduction on the right even if the oxidation is not actually happening - eg shown in image

draw the conventional cell diagram for the following cell:

Mg²⁺/Mg and Co²⁺/Co

draw the conventional cell diagram for the following cell:

Fe²⁺/Fe and MnO₄⁻/Mn²⁺

draw the conventional cell diagram for the following cell:

Cl₂/Cl⁻ and Sn⁴⁺/Sn²⁺

outline what changing the concentrations of a species in a redox half equation will do?

• it will change the position of the equilibrium

• changing the position of equilibrium will change the E⍬ value for the half-cell as the number of electrons are either increasing or decreasing

explain how changing the position of equilibrium both ways changes the value of E⍬

*Cu⁺ raised

*Cu⁺

outline how you would increase and decrease Ecell

what changes to concentration would make the Ecell more positive + therefore the reaction more feasible?

electrochemical cells are ____ and can be used as what?

• batteries

• can be used as a commercial source of electrical energy

cells can be ….. or …..

rechargeable or non-rechargeable

outline what happens in non-rechargeable cells

• the chemicals are used up over time + the EMF drops

• once one or more of the chemicals have been completely used up, the cell is flat + the EMF is 0 volts

what then happens to non-rechargeable cells?

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

give an example of a non-rechargeable cell

• the Zn/Cu cell

• also known as ‘The Daniell Cell’

why is the Zn/Cu cell non-rechargeable?

• because the Zn electrode over time starts to run out meaning that the equation cannot be reversed

• also the solutions make it not practical for transport

another example of a non-rechargeable cell is Zinc/Carbon cells

outline what happens during this cell

• zinc is negative electrode + acts as a case + has a paste electrolyte solution

• carbon acts as positive electrode + is located as a rod down centre of canister — it behaves like the platinum in the hydrogen cell

what will happen to the emf of non-rechargeable cells over time?

it will decrease over time as the reactants are used up

the zinc is oxidised to Zn²⁺, causing the case to wear away

deduce one essential property of the non-reactive porous separator labelled in the diagram

to allow the movement of ions

alkaline batteries are a _____ cost cell but has a _____ life

• higher cost cell

• longer life

outline rechargeable cells

these are cells where the current (electrons) can be forced back the opposite direction by an applied external current (plugging them in) so that they can work again like new

what happens in the equation for rechargeable cells?

the half cell equations are reversed when it recharges

what is the main difference of rechargeable cells than non-rechargeable cells?

the electrodes are designed so that they do not decompose or wear down as quickly → these can be very small, light + portable

describe lithium ion batteries

• lithium is a lightweight metal + gives for a light weight, small battery (cell)

• the electrolyte phase (liquid) is a polymer which means the battery wont leak

• this cell can generate about 4V so can be found in most laptops, tablets + phones

give the half equations, + overall equations for the lithium ion battery during discharge and recharge

symbol is forwards arrow

describe fuel cells

• these are somewhat different to other cells

• they have a continuous supply of the chemicals into the cell + so neither run out of chemicals nor need re-charging

• but they do need to have a constant supply of the required chemicals

what is the most common fuel cell?

the hydrogen-oxygen fuel cell

outline the hydrogen oxygen fuel cell?

• this is the alternative to actually burning hydrogen as a fuel

• it consists of two platinum electrodes separated by a special polymer electrolyte which will allow ions to pass through it

in what conditions can the hydrogen-oxygen fuel cell be run in?

in alkaline or acidic conditions but the overall equation + overal emf is the same

give the two half equations + E⍬ for the hydrogen-oxygen fuel cell in acidic conditions

give the two half equations + E⍬ for the hydrogen-oxygen fuel cell in alkaline conditions