Electrochemistry Notes

Electrochemistry

WACE 2019 Question 28

• Chlorine in calcium hypochlorite and chloride from hydrochloric acid are converted to chlorine gas.

(a) Oxidation number for chlorine in:

• Calcium hypochlorite, Ca(OCl)_2

• Hydrochloric acid, HCl

(b) Half-Equations:

• Chlorine gas produced by oxidation of one substance and reduction of the other.

• Oxidation half-equation.

• Reduction half-equation.

Half-Equation Details

• Correctly identifying oxidation and reduction half-equations (1 mark).

Oxidation Half-Equation

• 1 mark for correct reactants and products.

• 1 mark for correct balancing.

• Example: 2Cl^-(aq) \rightarrow Cl_2(g) + 2e^-

Reduction Half-Equation

• 1 mark for correct reactants and products.

• 1 mark for correct balancing.

• Example: Ca(OCl)2(s) + 4H^+(aq) + 2e^- \rightarrow Ca^{2+}(aq) + 2H2O(aq) + Cl_2(g)

Standard Electrode Potentials

• Comparing standard electrode potentials.

• Predicting reactivity from SRP.

Syllabus: Redox Reactions

• Oxidation-reduction (redox) reactions involve the transfer of one or more electrons from one species to another.

• Oxidation: loss of electrons.

• Reduction: gain of electrons.

• Half-equations and redox equations (acidic conditions only).

• Identifying species oxidized and reduced using oxidation numbers.

• Relative strength of oxidizing and reducing agents determined by comparing standard electrode potentials; used to predict reaction tendency.

Reactivity Series

• Some metals react readily (e.g., sodium, potassium) and are stored under paraffin oil.

• Oxidation of iron is a significant problem in the oil/gas industry.

• Inert metals do not corrode or react easily with oxygen.

• Reactivity series: ranks metals by their reactivity or ability to act as reducing agents.

• Half-equations are shown in the series.

Reactivity Series Examples

• Potassium, sodium, calcium, magnesium, aluminum, carbon, zinc, iron, tin, lead, hydrogen, copper, silver, gold, platinum.

• Calcium nitrate, aluminum nitrate, tin nitrate, silver nitrate.

• Word equations for reactions between metals and compounds.

• Displacement reactions.

Reduction Potential

• Reduction potential (redox potential, oxidation/reduction potential, or Eh) measures the tendency of a chemical species to acquire electrons and be reduced.

• Measured in volts (V) or millivolts (mV).

• Each species has its own intrinsic reduction potential.

• The more positive the potential, the greater the species’ affinity for electrons, or the more the species tends to be reduced.

Standard Reduction Potential

• Defined relative to a standard hydrogen electrode (SHE) reference electrode, arbitrarily given a potential of 0.00V.

• Values are standard reduction potentials for half-reactions measured at 25 °C, 100 kPa, and pH 7 in aqueous solution.

• Examples:

  • CH3COOH + 2H^+ + 2e^- \rightarrow CH3CHO + H_2O (-0.58 V)

  • 2H^+ + 2e^- \rightarrow H_2 (0.0 V)

  • O2 + 2H^+ + 2e^- \rightarrow H2O_2 (+0.7 V)

  • O2 + 4H^+ + 4e^- \rightarrow 2H2O (+1.64 V)

• Comparing standard reduction potentials helps determine how a reaction will proceed.

SRP: Standard Reduction Potential

• Reduction Potential: tendency for the half-reaction to occur as a reduction half-reaction in an electrochemical cell.

• Electrode Potential: the difference in potential between an electrode and its solution.

• Potential Difference (Voltage): measure of energy required to move a certain electric charge between electrodes, measured in volts.

• Standard Reaction Potential (E°): half-cell measured relative to a potential of zero for the standard hydrogen electrode (SHE), measured at STP at 1 mol/L.

WACE 2018 Questions 14 and 15

• Students investigated the reactivity of rhenium, vanadium, zirconium, and tantalum by placing metals in test tubes with nitrate salt solutions (1.00 mol/L).

Results Summary Table

Question 14

• Which metal is most easily oxidized?

  • (a) rhenium, (b) vanadium, (c) zirconium, (d) tantalum

Question 15

• Which metal is the weakest reducing agent?

  • (a) rhenium, (b) vanadium, (c) zirconium, (d) tantalum

SYLLABUS Redox Reactions

• Metal and halogen displacement reactions, and combustion in both limited and excess oxygen environments.

Metal Displacement

• A metal ion higher in the reactivity series (an oxidizing agent) will react spontaneously with a metal lower in the series (a reducing agent).

• Example: Copper metal in silver(I) nitrate solution.

Predicting Metal Displacement Reactions: Example

• Pearson p204 Worked example 8.4.1

• Predict whether zinc will displace copper from a solution containing copper(II) ions. Write the overall redox equation if appropriate.

Predicting Metal Displacement Reactions: Try Yourself

• Pearson p204 Worked example: Try yourself 8.4.1

• Predict whether copper will displace gold from a solution containing gold ions. Write the overall redox equation if appropriate.

• Pearson Key Questions Ch8.4 p205

CELLS - Introduction

• An electric cell converts chemical energy into electrical energy.

• Alessandro Volta invented the first electric cell, inspired by Luigi Galvani.

• Galvani observed that two different metals could make a frog's legs twitch.

• Galvani thought this was due to “animal electricity”.

• Volta recognized the experiment's potential.

• A battery is defined as two or more electric cells connected in series for steady current.

• Volta’s first battery: bowls of brine (NaCl(aq)) connected by metals.

• Revised design: a sandwich of two metals separated by paper soaked in salt water.

Introduction (Continued)

• Volta’s invention was an immediate technological success.

• Produced electric current more simply and reliably than static electricity methods and also produced a steady electric current.

• Electric cells are composed of two electrodes (solid electrical conductors) and at least one electrolyte (aqueous electrical conductor).

• In current cells, the electrolyte is often a moist paste.

• Positive electrode: cathode.

• Negative electrode: anode.

• Electrons flow through the external circuit from the anode to the cathode.

• To test voltage, red(+) lead to cathode (+ electrode) and black(-) lead to anode (- electrode).

Electrochemical Cells

• Define anode, cathode, anion, cation, salt bridge/porous cup, electrolyte, and voltaic cell.

• Predict and write the half-reaction equation that occurs at each electrode in an electrochemical cell.

Galvanic Cells (Voltaic Cells)

• A device that spontaneously produces electricity by redox.

• Uses chemical substances that participate in a spontaneous redox reaction.

• Reduction half-reaction will be above the oxidation half-reaction in the activity series (SRP).

• Composed of two half-cells: metal rod or strip immersed in a solution of its own ions or an inert electrolyte.

• Electrodes: solid conductors connecting the cell to an external circuit.

• Anode: electrode where oxidation occurs (-).

• Cathode: electrode where reduction occurs (+).

• Electrons flow from anode to cathode (“a before c”) through an electrical circuit.

• Porous boundary separates the two electrolytes, allowing ions to flow to maintain cell neutrality.

• Often a salt bridge (inert aqueous electrolyte such as Na2SO4(aq) or KNO_3(aq)).

• Or a porous cup containing one electrolyte within a container of a second electrolyte.

Cell Notation

• Single line: phase boundary (electrode to electrolyte).

• Double line: physical boundary (porous boundary).

• RED CAT AN OX

Observations

Zn Half Cell

• Zinc.

• Zn \rightarrow Zn^{2+} + 2e^-

Cu Half Cell

• Copper.

• Cu^{2+} + 2e^- \rightarrow Cu
  *Electron flow

Questions

• Sn(s) | Sn^{4+}(aq) || Cu^{2+}(aq) | Cu(s)

• Mg(s) | MgCl2(aq) || SnCl4(aq) | Sn(s)

• Sn(s) | SnCl2(aq) || CuCl2(aq) | Cu(s)

• Mg(s) | Mg^{2+}(aq) || Cu^{2+}(aq) | Cu(s)

• Mg(s) | Mg^{2+}(aq) || Sn^{2+}(aq) | Sn(s)

• Sn(s) | SnCl2(aq) || SnCl4(aq) | Sn(s)

• Copper placed in copper(II) chloride and tin metal placed in tin(II) ions.

• A copper-magnesium cell.

• Magnesium in magnesium chloride and tin in tin(II) chloride.

• Tin(IV) ion solution with tin and magnesium ions with magnesium.

• Two tin electrodes in tin(II) chloride and tin(IV) chloride.

• Copper in copper(II) solution and tin in tin(IV) solution.

• Match notation to descriptions.

Example: Silver-Copper Cell

• Cu(s) | Cu^{2+}(aq) || Ag^+(aq) | Ag(s)

• Memory device: “An ox ate a red cat” (Anode oxidation; reduction cathode).

• Cathode: strongest oxidizing agent.

• Anode: strongest reducing agent.

• E{overall} = E{red} + E_{ox}

• Switch the sign of E_{ox}.

Explanation

• Silver ions are the strongest oxidizing agents, undergo reduction at the cathode, creating more Ag(s).

• Copper atoms are the strongest reducing agents, give up electrons in oxidation at the anode, enter solution as Cu^{2+} (blue ions).

• Electrons from copper oxidation travel to silver cathode.

• Ag^+(aq) wins tug-of-war for electrons over Cu^{2+}(aq).

• Positive silver ions removed from solution; solution would become negatively charged, but cations move from the salt bridge to maintain electrical neutrality.

• Anions move from the salt bridge into the anode compartment to maintain electrical neutrality.

Interpretation: Silver-Copper Cell

Interpretation Table

SUMMARY

• Electrode: Solid electrical conductor (metal, graphite).

• Electrolyte: Aqueous electrical conductor (SNAPE).

• Cathode: Positive electrode (+); reduction occurs.

• Anode: Negative electrode (-); oxidation occurs.

• Electron flow: anode → cathode (“a before c”).

• Cations towards the cathode.

• Anions towards the anode.

Voltaic Cells with Inert Electrodes

• Inert electrodes needed when the SOA or SRA is not solid.

• Graphite (C(s)) rod or platinum strip is used as the electrode.

• Inert electrodes provide a location to connect a wire and a surface on which a half-reaction can occur.

Example

• Write equations for the half-reactions and the overall reaction in the cell: C(s) | Cr2O7^{2-}(aq), H^+(aq) || Cu^{2+}(aq) | Cu(s)

Cathode

• Cr2O7^{2-}(aq) + 14H^+(aq) + 6e^- \rightarrow 2Cr^{3+}(aq) + 7H_2O(l)

Anode

• 3[Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-]

• 3Cu(s) + Cr2O7^{2-}(aq) + 14H^+(aq) \rightarrow 3Cu^{2+}(aq) + 2Cr^{3+}(aq) + 7H_2O(l)

Observations

• Copper electrode mass decreases, and blue color increases, indicating oxidation at the anode.

• Carbon electrode remains unchanged, but the orange dichromate solution becomes less intense and turns greenish-yellow, indicating reduction.

WACE 2018 Questions 5 and 6

• Indium metal, In(s), in acidified dichromate solution, Cr2O7^{2-}(aq)

• Reaction: Cr2O7^{2-}(aq) + 14H^+(aq) + 2In(s) \rightarrow 2Cr^{3+}(aq) + 7H_2O(l) + 2In^{3+}(aq)

• EMF at 25.0 °C is +1.70 V.

Question 5

• What is the calculated E° value for the In^{3+}/In half-equation?

  • (a) -0.34V (b) 0.34V (c) 1.36V (d) 3.06V

Question 6

• Which sets of metals cannot be oxidized by indium ion, In^{3+}, under standard conditions?

  • (a) Sn, Cd, Fe, Cr (b) Mg, Na, Ca, Sr (c) Mn, Ni, Sn, Cu (d) Ni, Sn, Cu, Ag

WACE 2018 Question 25

Which statement pairs can be used to distinguish between an electrolytic cell and a galvanic cell?

Pearson

• Ch9.2 p223

Predicting the Operation of a Galvanic Cell

• A cell is made from Ag^+(aq)/Ag(s) and Fe^{2+}(aq)/Fe(s) half-cells under standard conditions and at 25°C. Use the electrochemical series to predict the overall cell reaction, identify the anode and cathode, and determine the direction of electron flow.

Thinking

• Identify two relevant half-equations in the electrochemical series.

• Identify strongest oxidizing agent (left with most positive E° value) and strongest reducing agent (bottom right).

• Write the two half-equations.

• The strongest oxidizing agent will react with the strongest reducing agent.

• Hint: Reduction equation has the most positive E° value and oxidation equation has the most negative E° value.

• Write the overall cell equation.

Working

• Ag^+(aq) + e^- \rightleftharpoons Ag(s) E° = +0.80 V

• Fe^{2+}(aq) + 2e^- \rightleftharpoons Fe(s) E° = -0.44 V

• Ag^+ is higher, thus strongest oxidizing agent.

• Fe is lower, thus strongest reducing agent.

Reduction

• Ag^+(aq) + e^- \rightleftharpoons Ag(s)

Oxidation

• Fe(s) \rightleftharpoons Fe^{2+}(aq) + 2e^-

Multiply the Ag/Ag^+ half-cell equation by two: [Ag^+(aq) + e^- \rightleftharpoons Ag(s)]*2 and then reverse to account of Oxidation

• Fe(s) \rightarrow Fe^{2+}(aq) + 2e^-

Overall cell reaction

• 2Ag^+(aq) + Fe(s) \rightarrow 2Ag(s) + Fe^{2+}(aq)

Identify the anode and the cathode in this cell.

• The anode is the electrode at which oxidation occurs.

• The cathode is the electrode at which reduction occurs.

The silver electrode will be the cathode and the iron electrode will be the anode.

Pearson Ch9.2 p223

Determine the polarities of the electrodes and the direction of electron flow in the cell.

The anode is negative; the cathode is positive.

Electrons flow from the negative electrode (anode) to the positive electrode (cathode), as shown in Figure 9.2.6.

Pearson Ch9.2 p224

• Worked example: Try yourself 9.2.1

Predicting the Operation of a Galvanic Cell

• A cell is made from Sn^{2+}(aq)/Sn(s) and Ni^{2+}(aq)/Ni(s) half-cells under standard conditions and at 25°C. Use the electrochemical series to predict the overall cell reaction, identify the anode and cathode, and determine the direction of electron flow.

Standard Cell Potential, E°_{cell}

• Spontaneous reactions have a positive E° (“e naught”) value. E°{cell} = E°{cathode} - E°_{anode}

• The higher the E°_{electrode}, the stronger the oxidizing agent (OA).

• Negative E° means the metal electrode is more willing to give up its electron; this is not favoured. These reactions prefer oxidation over reduction.

• When a half-cell is multiplied by a constant (for balancing), the E° value is NOT multiplied.

• When a reaction is reversed (flipped), the sign of the E° value switches.

• In a voltaic cell, the half-rxn with the more negative standard electrode potential is the anode, where oxidation occurs.

Cell Potential

• The potential voltage a reaction can produce.

• Cu^{2+} + 2e^- \rightarrow Cu E° = 0.34 V

• Ag^+ + e^- \rightarrow Ag E° = 0.80 V

• Reduction potentials

• Because this is a spontaneous process:

  • (Ag^+ + e^- \rightarrow Ag) \times 2 E° = 0.80V

  • Cu \rightarrow Cu^{2+} + 2e^- E° = -0.34 V

  • Cu + 2Ag^+ \rightarrow Cu^{2+} + 2Ag E° = 0.46 V

• Since both reactions are reduction, one must be oxidation. Flip it. A positive voltage must result from spontaneous reactions.

Cell Potential (Continued)

• The potential voltage a reaction can produce.

• Na^+ + e^- \rightarrow Na E° = -2.71 V

• Cl_2 + 2e^- \rightarrow 2Cl^- E° = 1.36 V

• Because this is a nonspontaneous process:

  • 2Na^+ + 2Cl^- \rightarrow 2Na + Cl_2 E° = -4.07 V

  • (Na^+ + e^- \rightarrow Na) \times 2 E° = -2.71 V

  • 2Cl^- \rightarrow Cl_2 + 2e^- E° = -1.36 V

• Non-spontaneous, must end in negative voltage. Flip one to become oxidation!

Rules

• Determine which electrode is the cathode. The cathodes are the electrodes where the strongest oxidizing agent (REDUCTION) present in the cell reacts.

  • I.e. The OA that is closet to the top on the left side of the redox table = SOA

  • If required, copy the reduction half-reaction for the strongest oxidizing agent and its reduction potential

• Determine which electrode is the anode. The anode is the electrode where the strongest reducing agent (OXIDATION) present in the cell reacts.

  • I.e. The RA that is closet to the bottom on the right side of the redox table = SRA

  • If required, copy the oxidation half-reaction (reverse the half-reaction)

• Determine the overall cell reaction. Balance the electrons for the two half reactions (but DO NOT change the E°_r) and add the half-reaction equations.

• Determine the standard cell potential, E°{ocell} using the equation: E°{cell} = E°{cathode} – E°{anode}

Cell Potential Example

• What is the standard potential of the cell represented below:

• Determine the cathode and anode

• Determine the overall cell reaction

• Determine the standard cell potential

Cell Potential Example (Continued)

• What is the standard potential of an electrochemical cell made of a cadmium electrode in a 1.0 mol/L cadmium nitrate solution and a chromium electrode in a 1.0 mol/L chromium(III) nitrate solution?
  Cd2+(aq) Cd(s) Cr2+(aq) Cr(s) H2O(l)

• E°{cell} = E°{rcathode} – E°_{ranode} = (-0.40V) – (-0.91V) = + 0.51V

• The E°_{cell} is positive, therefore the reaction is spontaneous.
  SOA SRA
  cathode anode

Cell Potential Example (Continued)

• A standard lead-dichromate cell is constructed. Write the cell notation, label the electrodes, and calculate the standard cell potential.
  *
  \ P b(s) P b^{2+}(a q) C r{2} O{7}^{2-}(a q) H^{+}(a q) C r^{3+}(a q) C(s)

• E°{cell} = E°{r cathode} – E°_{r anode} = (+1.36V) – (-0.13V) = + 1.49V

• The E°_{cell} is positive, therefore the reaction is spontaneous.

• SRA SOA

• anode cathode

Cell Potential Example (Continued)

• A standard scandium-copper cell is constructed and the cell potential is measured. The voltmeter indicates that copper the copper electrode is positive.
  \ S c(s) S c^{3+}(a q) C u^{2+}(a q) C u(s) E_{cell} = +2.36V

• Write and label the half-reaction and net equations, and calculate the standard reduction potential of the scandium ion.

• E{cell} = E{r cathode} - E_{r anode}
  2.36V = (+0.34V) - (x)

• anode cathode

Observations

Cl half cell

• chlorine gas

• platinum cathode

• anions

• Cl^− solution

• Cl_2(g) + 2e^− \rightarrow 2Cl(aq)

Cu half cell

• direction of electron flow

• galvanometer

• cations

• Cu^{2+} solution

• Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-

• copper anode

Pearson Ch9.3 p226

• Worked example 9.3.1

Predicting Direct Redox Reactions

Consider the following equations that appear in the order shown in the electrochemical series.
\ B r_{2}(a q)+2 e^{-} 2 B r(a q) E^{\circ}=+1.09 V
\ N i^{2+}(a q)+2 e^{-} N i(s) Е^\circ=-0.23 V
\ M g^{2+}(a q)+2 e^{-} M g(s) E^{\circ}=-2.34 V

Use the electrochemical series to predict the effect of mixing:

• a B r_{2}(a q) and M g^{2+}(a q)

• b M g^{2+}(a q) and Ni(s)

• c N i^{2+}(a q) and Mg(s)

Thinking

• Identify the two relevant half-equations in the electrochemical series.

• Predict whether or not a reaction occurs. A chemical species on the left (an oxidizing agent) of the electrochemical series reacts with a chemical species on the right (a reducing agent) that is lower in the series.

• Write the overall equation.

Working

• \ B r_{2}(a q)+2 e^{-} 2 B r(a q) E^{\circ}=+1.08 V

• \ M g^{2+}(a q)+2 e^{-} M g(s) E^{\circ}=-2.36 V

• No reaction occurs because both B r_{2}(a q) and M g^{2+}(a q) are oxidizing agents.

• N i^{2+}(a q)+2 e^{-} N i(s) E^{\circ}=-0.24 V

• \ M g^{2+}(a q)+2 e^{-} M g(s) E^{\circ}=-2.36 V

• No reaction occurs because the oxidizing agent, M g^{2+}, is below the reducing agent, Ni, in the electrochemical series.
  \ N i^{2+}(a q)+2 e^{-} N i(s) E^{\circ}=-0.24 V

• \ M g^{2+}(a q)+2 e^{-} M g(s) E^{\circ}=-2.36 V

• A reaction occurs because the oxidizing agent, N i^{2+}, is above the reducing agent, Mg, in the electrochemical series.

• The higher half-equation occurs in the forward direction:

Worked example 9.3.1

Predicting Direct Redox Reactions

Consider the following equations that appear in the order shown in the electrochemical series.
\ B r_{2}(a q)+2 e^{-} 2 B r(a q) E^{\circ} +1.09 V
\ N i^{2+}(a q)+2 e^{-} N i(s) Е^\circ=-0.23 V
\ M g^{2+}(a q)+2 e^{-} M g(s) E^{\circ}=-2.34 V

• Use the electrochemical series to predict the effect of mixing:

• a B r_{2}(a q) and M g^{2+}(a q)

• b M g^{2+}(a q) and Ni(s)
  \ N i^{2+}(a q)+2 e^{-} \rightarrow N i(s)
  The lower half-equation occurs in the reverse direction:

\ M g(s) \rightarrow M g^{2+}(a q)+2 e^{-}
The overall reaction equation is found by adding the half-equations:

\ N i^{2+}(a q)+M g(s) \rightarrow N i(s)+M g^{2+}(a q)

• c N i^{2+}(a q) and Mg(s).

Pearson Ch9.3 p227

• Worked example: Try yourself 9.3.1

Predicting Direct Redox Reactions

Consider the following equations, which appear in the order shown in the electrochemical series:
\ C l_{2}(g)+2 e^{-} 2 C I^{-}(a q) E^{\circ}=+1.36 V
\ 12(s)+2 e^{-} 21^{-}(a q) E^{\circ}=+0.54 V
\ P b^{2+}(a q)+2 e^{-} P b(s) E^{\circ}=-0.13 V

Use the electrochemical series to predict the effect of mixing:

• a l_{2}(s) and P b^{2+}(a q)

• b C l(a q) and l_{2}(s)

• c C l(a q) and Pb(s).

Pearson Ch9.3 p228

KEY QUESTIONS

1 Which one of the following metals would you expect to be coated with lead when immersed in a solution of lead(II) nitrate?
A copper
B cobalt
C silver
D gold

2 Which one of the following species would react with H{2} S(g) but not with H{2} O(1) under standard conditions?
A C l_{2}(g)
B M g(s)
C A g^{+}(a q)
D C u(s)

3 Using the electrochemical series, predict whether a reaction will occur in the following situations. If a reaction does occur, write the overall equation for the reaction.
a Chlorine gas is bubbled into a solution containing bromide ions.
b Chlorine gas is bubbled into a solution containing iodide ions.
c A bromine solution is added to a solution containing chloride ions.
d A bromine solution is added to a solution containing iodide ions.

4 A reaction occurs when a strip of zinc metal is placed in a silver nitrate solution.
a Write the overall equation for the reaction.
b Describe the energy change that takes place in this reaction.

5 Iron nails are placed into 1 mol L^{-1} solutions of CuSO4, MgCl2, Pb(NO3)2 and ZnCl_2. Use the electrochemical series to identify in which solution(s) you would expect a coating of a metal other than iron to appear on the nail.

Electrolytic Cell Rxns

• Reactions that require an energy source to react.

• When electric voltage is used to produce a redox reaction, it is called electrolysis.

Electrolytic Cell

• Electrolytic Cell – a cell in which a non-spontaneous redox reaction is forced to occur; a combination of two electrodes, an electrolyte, and an external power source.

• Electrolysis – the process of supplying electrical energy to force a non-spontaneous redox reaction to occur

• The external power source acts as an “electron pump”; the electric energy is used to do work on the electrons to cause an electron transfer

• Electrons are pulled from the anode and pushed to the cathode by the battery or power supply

Electro