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Dr Damian MoranE-mail: Damian.Moran@mq.edu.au

Electrochemistry

The Production of Materials (part 2)

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Applications – Energy Storage

Tesla.com

Source: https://www.theguardian.com/world/2018/may/23/hamburg-first-german-city-ban-older-diesel-cars-air-quality-pollution

Materials production (e.g. electrolytic refining of metals)Corrosion (built and natural environment)

Energy storage and production

Source: http://www.time.com/4846761/france-to-ban-sales-of-gas-diesel-vehicles-by-2040

Source: https://www.theguardian.com/politics/2017/jul/25/britain-to-

ban-sale-of-all-diesel-and-petrol-cars-and-vans-from-2040

Source: Macquarie Bank

Source: Macquarie Bank

Source: Macquarie Bank

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Applications – Industrial Storage

• Large batteries are capable of supporting renewable energy sources, night and day. However, they also have their limits (episode 848).

Source: https://www.news.com.au/technology/environment/thats-a-record-south-australias-tesla-battery-responds-to-coalfired-plant-failure

Source: https://www.npr.org/sections/money/2018/06/15/620298266/episode-848-the-world-s-biggest-battery

Oxidation and Reduction (Redox)

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Source: BOSTES

BOSTES Glossary of Key Wordshttp://www.boardofstudies.nsw.edu.au/syllabus_hsc/glossary_keywords.html

Oxidation:• The complete removal of one or more electrons

from a molecular entity.

Reduction:• The complete transfer of one or more electrons

to a molecular entity.

Definitions: IUPAC Goldbook.

A ⇌ An+ + ne−

Bn+ + ne− ⇌ B

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Terminology

Oxidation is a loss of electrons (e.g. A ⇌ An+ + ne−)

Reduction is a gain of electrons (e.g. Bn+ + ne− ⇌ B)

• An oxidising agent (or oxidant) is a substance which oxidises some other substance by removing its electrons. Hence, the oxidising agent is reduced.

• A reducing agent (or reductant) is a substance which reducing some other substance by giving it electrons. Hence, the reducing agent is oxidised.

• Redox reactions involve transfer of electrons from one substance to another and always occur in pairs. That is, a redox reaction is the reaction between an oxidising agent (or oxidant) and a reducing agent (reductant).

Cu(s) ⇌ Cu2+(aq) + 2e–

Al3+(aq) + 3e– ⇌ Al(s)

Displacement Reactions

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• Displacement reactions are examples of redox reactions.

• In a displacement reaction a more “reactive” metal reduces the ion of a less reactive metal to the neutral atom.

• Copper is more easily oxidised than silver and displaces the silver from solution. Copper is said to be more “active” than silver.

• The activity series below summarises the ability of metals to displace aqueous metal ions from solution (ordered from most easily oxidised to least easily oxidised):

Li, K, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Ni, Sn, Pb, H2, Cu, Ag, Au

Cu(s) + 2 Ag+(aq) ⇌ 2 Ag(s) + Cu2+

(aq)

React with acids to produce H2 (g).

A(s) + Bn+ ⇌ An+ + B(s)

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Oxidation Number

Atoms in elements e.g. O2(g), C(s) 0

Sum of the oxidation states of atoms in neutral species 0

Sum of the oxidation states of atoms in charged species charge on ion

Fluorine in all compounds –1

Hydrogen in compounds with non-metals e.g. HCl, CH4 +1

Oxygen in all compounds except peroxides (–1) –2

Elements in group 17 (aka 7A) –1

Valence of the element

• Redox reactions are just one type of chemical reaction and sometimes it is not clear whether a reaction involves oxidation and reduction.

• Oxidation numbers, also known as oxidation states, enable us to quickly assess whether a redox reaction has occurred.

• There are useful rules for determining oxidation number:*

*Note that these are hierarchical.

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Oxidation Number

(a) Cu(s) ⇌ Cu2+(aq) + 2e–

(b) Al3+(aq) + 3e– ⇌ Al(s)

0 +2

+3 0

‘Oil rig’ mnemonic- oxidation is loss of electrons (oil)

- reduction is gain of electrons (rig)

⇒ increase in oxidation number⇒ Cu(s) has been oxidised to Cu2+(aq)

⇒ decrease in oxidation number⇒ Al3+(aq) has been reduced to Al(s)

• Oxidation is a loss of electrons and an increase in oxidation number• Reduction is a gain of electrons and a reduction in oxidation number

http://acrigs.com/

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Half Reactions

a. Cu(s) Cu2+(aq) + _ e–

b. Al3+(aq) + _ e– Al(s)

c. _ F–(aq) F2(g) + _ e–

d. Ag+(aq) + _ e– Ag(s)

e. _ H2O(l) O2(g) + _H+(aq) + _ e–

2

3

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2 4 4

0 +2

+3 0

−1 0

+1 0

+2 −2 0

1

(+1 per H)

• Half reactions (or half equations) are reactions which describe oxidation and reduction processes separately in terms of electrons lost or gained.

Reduction half-reaction: Bn+ + ne− ⇌ B

Oxidation half-reaction: A ⇌ An+ + ne−

[ ( 2+) + (2-) = 0 ]

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Given Pb(s) Pb2+(aq) + 2e–

Cl2(g) + 2e– 2Cl–(aq)

Pb(s) Pb2+(aq) + 2e–

+) Cl2(g) + 2e– 2Cl–(aq)

Pb(s) + Cl2(g) Pb2+(aq) + 2Cl–(aq)0 0 +2 -1

Cl2(g) is reduced and acting as an oxidantPb(s) is oxidised and acting as a reductant

Redox Reactions

• A complete redox reaction is derived by combining two half-reactions. Note that it is necessary to balance all species, including the number of electrons, as the complete redox equation must not include electrons.

Oxidation & Reduction (Redox)

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Source: BOSTES

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Galvanic Cells

Zn|Zn2+||Sn2+|Sn

e– e–

Zn(s)

Zn2+(aq)

Sn(s)

Sn2+(aq)

Cl– K+Anode Cathode

Oxidataion Reduction Cathode Anode

An electrochemical cell that derives electrical energy from spontaneous redox reactions

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i. In the left-hand cell, electrons are generated.

ii. For the reaction taking place in the right-hand cell, electrons are consumed.

iii. Electrons travel from the left-hand cell to the right-hand cell.

iv. When electrons move along a conducting wire, current is generated along the conducting wire.

v. When electrons move, there must be a difference in voltage/potential between the left-hand electrode and the right-hand electrode.

vi. As the reaction in the left-hand cell progresses, the electrode will lose mass.

e– e–

Zn(s) Sn(s)

Cl– K+

anode cathode

Galvanic Cells

Zn2+(aq) Sn2+(aq)

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i. Placed between two half-cells, a salt bridge consists of Agar gel entrapped with cations (e.g. K+, Na+) and anions (e.g. Cl-, NO3

-).

ii. In the absence of a salt bridge, as the reaction progresses, we expect that:- the anode will become increasingly positive- electrons will stop flowing into the electrolyte solution of the right hand cell.

iii. Using the salt bridge, anions will move towards the left-hand cell, while cations will move towards the right-hand cell.

iv. In the presence of a salt bridge a complete circuit is formed and current can flow through the entire electrochemical cell.

Salt Bridge

e– e–

Zn(s) Sn(s)

Cl– K+

anode cathodeZn2+(aq) Sn2+(aq)

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Setting-up a Galvanic Cell

First, calculate the theoretical standard potential for the galvanic cell using tabulated standard electrode potentials:

This is the maximum voltage that a cell can deliver. The voltage arises due to the potential difference across the electrodes.

The potential of an electrode relative to the Standard Hydrogen Electrode (SHE).

0 VSHE

+3 V

−2 V

5 V

Substance A

Substance B

• Voltage is a relative measurement. The standard potentials found in the HSC data sheet are measured relative to the standard hydrogen electrode (SHE).

• By definition, the SHE has a potential of 0.00 V.

2H+(aq) + 2e− ∏ H2(g) (Eʅ = 0.00 V)

• Conditions are fixed:

- [H+] 1.000 mol L-1

- Black platinum electrode- H2(g) 100.0 kPa- Temperature 298.15 K

• Conditions are fixed:

Standard Hydrogen Electrode (SHE)

mV

H2(g)

Pt(s)

H+ (aq)1 mol L−1

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EO (1 Atm) Eʅ (100 kPa) versus

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2.36 V

Mg| Mg2+ half-cell

Mg is oxidised

solutions = 1.000 mol L-1 gases = 100.0 kPa

Mg2+ (aq)1 mol L−1

Mg(s)

V

H2(g)

Pt(s)

H+ (aq)1 mol L−1

Standard Electrode Potential (Eʅ)

Nb:| = phase change|| = salt bridge

H+ | Pt, H2

H+ is reduced

Standard conditions:temperature = 298.15°C pure solids or liquids

• Voltage is a relative measurement. The standard potentials found in the HSC data sheet are measured relative to the standard hydrogen electrode (SHE).

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Standard Electrode Potential (Eʅ)

Source: BOSTES

• Voltage is a relative measurement. The standard potentials found in the HSC data sheet are measured relative to the standard hydrogen electrode (SHE).

• Note that half-reactions are written as reduction reactions

Setting-up a Galvanic Cell

First, calculate the theoretical standard potential for the galvanic cell using tabulated standard electrode potentials:

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Cd2+(aq) + 2e– Cd(s) Eʅ = –0.40 VAg+(aq) + e– Ag(s) Eʅ = +0.80 V

Reverse the Cd2+|Cd equation because it has a less positive E°

Cd(s) Cd2+(aq) + 2e– Eʅ = +0.40 V +) 2Ag+(aq) + 2e– 2Ag(s) Eʅ = +0.80 V

Cd(s) + 2Ag+(aq) Cd2+(aq) + 2Ag(s) Eʅ = 1.20 V

Positive value ⇒ spontaneous net reaction!

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A galvanic cell is constructed such that one electrode component consists of an aluminium strip placed in a solution of Al(NO3)3, and the other has a nickel strip placed in a solution of NiSO4. The overall cell reaction is 2Al(s) + 3Ni2+(aq) → 2Al3+(aq) + 3Ni(s)Assume the Al is not coated with its oxide.

(a) What is being oxidised, and what is being reduced?Al(s) is oxidised; Ni2+(aq) is reduced.

(b) Write the half-reactions that occur in the two electrode compartments.Al(s) → Al3+(aq) + 3e–; Ni2+(aq) + 2e– → Ni(s)

(c) Which electrode is the anode, and which is the cathode?Al(s) is the anode; Ni(s) is the cathode

Practice question

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A galvanic cell is constructed such that one electrode component consists of an aluminium strip placed in a solution of Al(NO3)3, and the other has a nickel strip placed in a solution of NiSO4. The overall cell reaction is 2Al(s) + 3Ni2+(aq) → 2Al3+(aq) + 3Ni(s)

(d) Indicate the signs of the electrodes.Al(s) is negative; Ni(s) is positive

(e) Do electrons flow from the Ni to the Al electrode?No, from the Al (−ve) electrode to the Ni (+ve) electrode.

(f) In which directions do the ions migrate through the solution?Cations migrate toward the Ni(s) cathode.Anions migrate toward the Al(s) anode.

Practice question

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HSC practice question

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HSC practice question

0 +2

0

+2 +3

+2

+3 +3

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HSC practice question

EʅCell = 1.36 V − 0.16 V= 1.20 V

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HSC practice question

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HSC practice question

Cathode half-cell reaction: Cu2+(aq) + 2e− ∏ Cu(s) (Eʅ = +0.34 V)

As the reaction proceeds, Cu(s) forms at the cathode. The solutions blue colour disappears, as the Cu2+(aq) concentration is decreased.

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HSC practice question

EʅCell = EʅCathode − EʅAnode

EʅAnode = EʅCathode − EʅCell EʅAnode = 0.34 V − 2.02 V = −1.68 V

Cu2+(aq) + 2e− ∏ Cu(s) (Eʅ = +0.34 V)

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HSC practice question

Anode reaction: Zn(s) ∏ Zn2+(aq) + 2e− ; Cathode reaction: Ag+(aq) + e− ∏ Ag(s)Balanced redox reaction: Zn(s) + 2 Ag+(aq) ∏ Zn2+(aq) + 2 Ag(s)

Ratio: Zn(s) 2 Ag(s)

2 Ag(s)Zn(s)

or• As Zn is anode (left half-cell), it loses mass.

As Ag is cathode (right half-cell), it gains mass.

• Mass Zn lost is 1.00 g, which is:(1.00 g / 65.38 g mol-1) = 0.0152… mol

• Ratio Zn(s) consumed to Ag(s) produced:

• Mass Ag(s) produced in cathode half-cell:(0.0305…mol * 107.9 g mol-1) = 3.30 g

2 Ag(s)Zn(s) * 0.0152…mol = 0.0305…mol

⸫Increase in mass of the Ag(s) electrode:10.0 g + 3.30 g = 13.3 g

Oxidation & Reduction (Redox)

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Source: BOSTES

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DRY CELL

• Inexpensive• Low energy density

(cannot deliver high currents)

• Anodes corrode (may leak)• Not rechargeable

Graphite rod (metal cap)

Insulating packaging

Cathode +

Anode − Zinc cylinder

NH4Cl paste

MnO2, NH4Cl, C(s)

NH4+(aq) + MnO2(s) + H2O(l) + e− ∏ Mn(OH)3(s) + NH3(aq)

Zn(s) ∏ Zn2+(aq) + 2e−0 +2

+4 +3

(Leclanché dry cell)

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LITHIUM CELL

• High energy density (light and compact)

• Long shelf life• Wide temperature range

MnO2(s) + Li+(soln) + e− ∏ LiMnO2(s)

Li(s) ∏ Li+(soln) + e−0 +1

+4 +3

Cathode +

Anode −

Li

MnO2

Li saltOrganic solvent

(Lithium-manganese dioxide cell)

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YOUR FUTURE IS ELECTRIC

Good luck!