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Physical Chemistry Year 2 > Electrode Potentials + Electrochemical Cells > Flashcards

Electrode Potentials + Electrochemical Cells Flashcards

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1
Q

Basics

A
  • when electrons move , you get electricity
  • electrons move from more reactive metal to less reactive metal
  • in an electrochemical cell : ALWAYS 2 REACTIONS - REDOX
  • anode = positive = oxidation
  • cathode = negative = reduction
  • electrode potential + electromotive forever are the same thing - E°
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2
Q

Half cells / electrodes

A
  • when a rod of metal is dipped into a solution / its own ions , an equilibrium is set up
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3
Q

E.g. dipping zinc into zinc sulfate solution sets up the following equilibrium:

A

Zn (s) = Zn2+ (aq) + 2e-

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4
Q

What makes an electrochemical cell

A

HALF - CELL + HALF - CELL = ELECTROCHEMICAL CELL

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5
Q

An electrochemical cell is made from …

A
  • 2 different metals
  • both dipped in salt so,it ions of their own ions
  • connected by a wire ( the external circuit )
  • each one is a half - cell
  • circuit completed by a salt bridge
  • use a voltmeter in the external circuit to measure the voltage between 2 half - cells
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6
Q

What are salt bridges

A
  • complete circuit
  • piece of filter paper soaked in a solution of salt ( usually saturated potassium nitrate KNO3)
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7
Q

What are salt bridges used for ?

A
  • it allows ions to flow through + balance the charges to maintain neutrality
  • stops both going into equilibrium
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8
Q

Why are salt bridges used rather than a piece of wire ?

A
  • to avoid further metal / ion potentials in the circuit
  • they must be inert + not react with the solutions in the half - cells
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9
Q

What happens in the zinc / copper electrochemical cell ( zinc side ) ?

A
  1. Zinc loses electrons more easily than copper
    • ZN (S)= ZN2+ (aq) + 2e-
    • releases electrons into the external circuit
    • more negative electrode potential , showing it loses electrons more easily
    • better reducing agent
    • is oxidised
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10
Q

What happens in the zinc / copper electrochemical cell ( copper side ) ?

A
  1. Copper loses electrons less easily than zinc
    - Cu(s) + 2e- =Cu2+ (aq)
    - gains electrons from external circuit (same number as zinc lost )
    - more positive electrode potential , harder to lose electrons
    - is reduced
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11
Q

Overall equation in the cell

A

Cu 2+ (aq) + Zn (s) = Cu (s) + Zn2+ (aq)

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12
Q

What happens in electrochemical cell when voltmeter removed ?

A

Electrons allowed to flow , they would do so from zinc to copper

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13
Q

What following changes would take place ?

A
  1. Zinc would dissolve to form Zn2+ (aq) in solution -> increasing the concentration of Zn2+ (aq)
  2. The electrons would flow through the wire to the copper rod
  3. They would combine with Cu2+ (aq) ions from the copper sulfate solution -> depositing fresh copper on the rod + decreasing the conc of Cu2+ (aq) in solution
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14
Q

Daniel Cell

A

An electrochemical cell when two half cells are connected they generate electricity

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15
Q

Electrode potentials

A

These measure how easily a meta is oxidised
- how good of a reducing agent it is
- how easily it loses electrons
- if it’s more negative when connected to hydrogen , it is better reducing agent

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16
Q

What is a standard electrode potential

A

THE STANDARD ELECTRODE POTENTIAL, E°, OF A HALF - CELL IS THE VOLTAGE MEASURED UNDER STANDARD CONDITIONS WHEN THE HALF - CELL IS CONNECTED TO A STANDARD HYDROGEN ELECTRODE

Easy to oxidised = very negative potential
Harder to oxidise = less negative / positive potential

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17
Q

Use if the SHE

A
  • you measure the electrode potential of a half - cell against the SHE
  • this is to COMPARE THE TENDENCY OF DIFFERENT METALS TO RELEASE ELECTRONS
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18
Q

What is SHE defined as

A

The potential of SHE is defined as zero
- so if it’s connected to another electrode , the measured voltage ( the electromotive force) is the electrode potential of that cell if it’s at standard conditions

19
Q

How do they work

A
  • hydrogen gas is bubbled into a solution of H+ (aq) ions
  • since the hydrogen doesn’t conduct , electrical contact is made via a piece of unreactive platinum metal
  • metal is coated with finely divided platinum to increase the surface area + allow any reaction to proceed rapidly
20
Q

What are electrodes with negative E° values better at + show cell notation for a SHE

A
  • releasing electrons ( better reducing agents) than hydrogen

Pt | H2 | 2H+ || (whatever is being measured against )

21
Q

Standard conditions

A
  1. Solutions have conc of 1.00 moldm-3
  2. Temperature if solution is 298k (25°C)
  3. Pressure is 100 kPa

Changing these conditions will change the electrode potential

22
Q

Representing cells

A
  • more -ve electrode potential goes left
  • when giving the value of the emf state the polarity (I.e. wether it is positive/ negative ) of the right - hand electrode , as the cell representation is written
  • charges go closest to salt bridge
23
Q

Electrochemical series

A
  • a list of standard electrode potentials for electrochemical half - cells
  • arranged in this order with the most negative values at the top
  • the number of electrons involved in the reaction has no effect on the value of E°
24
Q

More +ve electrode potentials show that ….

A
  • left hand substance is more easily reduced
  • right hand substance is more stable
  • best oxidising agent (best at being reduced)
  • Ag+ is the best oxidising agent
25
Q

More -ve electrode potentials show that ….

A
  • right hand substance more easily oxidised
  • left hand substance is more stable
  • best reducing agent (best at being oxidised)
  • Li is the best reducing agent
26
Q

What is it written as and what do you need to do to it

A

Written as reduction eq , when working out any potentials , the more -ve value equation needs to be flipped

27
Q

Calculating standard cell potentials

A

The voltage obtained by connecting two standard electrodes together is found by the difference between the two E° values .

           E° cell = E° reduced - E° oxidised 
 Always more +ve - more -ve
28
Q

Predicting direction of redox rxns

A

Using standard electrode potentials we can predict if a rxn is feasible
- electrons move from more -ve to more +ve electrode

Using electrode potentials :
- the species OXIDISED is the REDUCING AGENT , this needs to have a more -ve potential value for it to be feasible
- the species REDUCED is OXIDISING AGENT , needs to have a more +ve potential for it to be feasible

29
Q

To work out if feasible

A
  • work out which species is being oxidised (loses electrons) - this is positive cell
  • work out which species is being reduced (gains electrons) - this is negative cell
  • look on electrochemical series at the emf values
  • more negative cell electrode potential value ( may have more + value) minus the more positive cell electrode potential ( may have more -ve value)
30
Q

How to tell feasibility after calculation

A

Feasible = +ve value
Not feasible = -ve value

31
Q

Example :
Fe3+ (aq) + Cl- (aq) -> Fe2+ (aq) + 1/2Cl2 (aq)

A

Iron being reduced so gaining electrons so this is the -ve electrode
Chlorine is being oxidised so losing electrons so is the +ve electrode

Fe3+ (aq) + e- -> Fe2+ (aq) = +0.77
Cl2 (aq) + 2e- -> 2Cl- (aq) = + 1.36

Negative electrode - positive electrode
0.77 - (+1.36) =-0.59

SO NOT FEASIBLE

32
Q

Types of electrochemical cells

A

Batteries are a type of electrochemical cell which provide the electricity we use to power things like watches + phones . A battery refers to a number of cells connected together.

RECHARGEABLE
- these can be recharged by reversing the cell rxns
- this is done by applying an external voltage greater than the voltage of the cell to drive the electrons in the opposite direction

NON-RECHARGEABLE
- non -rechargeable batteries are cheaper than rechargeable ones ,however, are worse is the long run as the don’t last longer

FUELL CELLS

33
Q

Lithium ion rechargeable battery

A
  • positive electrode is made of lithium cobalt oxide ,LiCoO2
  • negative electrode is carbon
  • these are arranged in layers with a sandwich of solid electrolyte in between
  • on charging , electrons are forced through the external circuit from positive to negative electrode + at the same time lithium ions move through the electrolyte towards the positive electrode to maintain the balance if charge
  • on discharging, electrons + Li+ ions move from the negative electrode back to the positive electrode
  • a single cell gives a voltage of between 3.5 v and 4.0 v , compared with around 1.5 v for most other cells , sed in laptops, tablets , smartphones
34
Q

Strength

A

Light ( lithium is a dense material) , won’t leak (electrode is solid polymer)

35
Q

Weaknesses

A

Limited life span , loses capacity,expensive , fire / explosion risk

36
Q

Eq @ +ve electrode

A

Li+ + CoO2 + e- -> Li+[CoO2]-

37
Q

Eq @ -ve electrode

A

Li -> Li+ + e-

38
Q

Alkaline hydrogen- oxygen fuel cell

A

-the cell has two electrodes if a porous platinum- based material
- hydrogen + oxygen are fed into 2 separate electrodes
- they are separated by a semipermeable anion membrane which allows OH- and water to pass through but not hydrogen or oxygen gas
- the electrolyte is sodium hydroxide solution

39
Q

Hydrogen enters at the -ve electrode and the following half reaction takes place :

A

2H2(g) + 4OH- -> 4H2O (l) + 4e-
EMF = -0.83V

40
Q

This releases electrons which flow through the circuit from the -ve hydrogen electrode to the +ve oxygen electrode . Oxygen enters + the following half rxn takes place :

A

O2 (g) + 2H20 (g) + 4e- -> 4OH-
EMF = +0.40 V

41
Q

What does this do

A

This accepts electrons from the other electrode + releases OH- ions which travel through the semi - permeable membrane

42
Q

Final overall effect / eq

A

2H2(g) + O2(g) -> 2H20 (g)
EMF = 1.23 V

43
Q

Benefits of fuel cells

A
  • more efficient - convert more available energy to kinetic energy
  • takes place at lower temp so no nitrogen produced
  • only waste product is water
  • no nasty toxic chemicals
  • no CO2 emissions from cell itself
  • don’t need to be recharged as long as hydrogen and oxygen are supplied
44
Q

Problems with fuel cells

A
  • need energy to produce hydrogen and oxygen
  • can be supplied by electrolysis but this requires energy from burning fossil fuels - co2 emissions
  • hydrogen us highly flammable
  • hard to store + transport hydrogen gas
  • burning 1g of hydrogen gives out around three times as much energy as burning 1g of petrol but the hydrogen takes up around 8000 times as much space

Physical Chemistry Year 2 (2 decks)

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