Classwork Series and Exercises {Chemistry- SS2}: Electrochemistry

Introduction

Electrochemistry is the study of electricity and how it relates to chemical reactions. In electrochemistry, electricity can be generated by movements of electrons from one element to another in a reaction known as redox reaction, or oxidation-reduction reaction.  In other word, electrochemistry is the study of changes that cause electrons to move. This movement of electrons is called electricity.

Electrolytic and Electrochemical Cells

An electrochemical cell is simply a device that converts chemical energy into electrical energy when a chemical reaction is occurring in a cell. An electrolytic cell converts electrical energy into chemical energy. In an electrochemical cell the reaction occurs spontaneously at the electrodes, while an electrolytic cell reaction is not spontaneous at the electrodes – the reaction has to be forced by applying an external electrical current.

Electrodes & Charge

The anode of an electrolytic cell is positive (cathode is negative), since the anode attracts anions from the solution. However, the anode of a electrochemical cell is negatively charged, since the spontaneous oxidation at the anode is the source of the cell’s electrons or negative charge. The cathode of a electrochemical cell is its positive terminal. In both electrochemical and electrolytic cells, oxidation takes place at the anode and electrons flow from the anode to the cathode.

 Galvanic or Voltaic Cells

Voltaic (Galvanic) Cells are electrochemical cells which contain a spontaneous reaction, and always have a positive voltage. The redox reaction in a galvanic cell is a spontaneous reaction. For this reason, galvanic cells are commonly used as batteries. Galvanic cell reactions supply energy which is used to perform work. The energy is harnessed by situating the oxidation and reduction reactions in separate containers, joined by an apparatus that allows electrons to flow. A common galvanic cell is the Daniel cell, shown below.

 

• Anode (negative terminal): Oxidation / Zinc strip immerses in zinc sulphate solution.
  Zn(s) –> Zn²⁺(aq) + 2e / Zinc strip becomes thinner.
• Cathode (positive terminal): Reduction / Copper strip immerses in copper(II) sulphate solution.
  Cu²⁺(aq) + 2e –> Cu(s) / A brown layer formed around copper strip. / Concentration Cu²⁺ ions decreases cause the intensity blue colour of solution decreases.
• Zinc is more electropositive than copper. Electrons are flowed from zinc strip to copper strip through the external circuit. (Note: Conventionally, electrons flow in the opposite direction of electrical current).
• Overall equation:  Zn(s) + Cu²⁺(aq) –> Zn²⁺(aq) + Cu(s)

As a result, the solution containing Zn(s) becomes more positively charged while the solution containing Cu(s) becomes more negatively charged. In order for the voltaic cell to work, the solutions in the two half-cells must remain electrically neutral. Therefore, a salt bridge containing KNO₃ is added to keep the solutions neutral by adding NO₃^–, an anion into the anode solution and K⁺, a cation into the cathode solution. As oxidation and reduction proceed, ions from the salt bridge migrate to neutralize charge in the cell compartments.

The diagram of this electrical cell : Zn_(s) │Zn²⁺_(aq) ║ Cu²⁺_(aq) │ Cu_(s)

Use ║ to separate anode(-) and cathode(+) and so represents the salt bridge in between.

Electrolytic Cells

The redox reaction in an electrolytic cell is non-spontaneous. Electrical energy is required to induce the electrolysis reaction. An example of an electrolytic cell is shown below, in which molten NaCl is electrolyzed to form liquid sodium and chlorine gas. The sodium ions migrate toward the cathode, where they are reduced to sodium metal. Similarly, chloride ions migrate to the anode and are oxidized to form chlorine gas. This type of cell is used to produce sodium and chlorine. The chlorine gas can be collected surrounding the cell. The sodium metal is less dense than the molten salt and is removed as it floats to the top of the reaction container.

Similarities between electrolytic and electrochemical cells

• redox reaction.
• Anode: oxidation
• Cathode: reduction
• Electrons flow from anode to cathode in the external circuit

Differences between electrolytic and electrochemical cells

Differences	Electrolytic Cell (Electrolysis)	Chemical Cell / Voltaic Cell
Structure	With electrical supply.	No electrical supply.
Electrodes	Can be the same or difference metal (graphite or platinum).	Must be two different metals.
Flows of electrons	From anode to cathode through external circuit.	From more electropositive metal to less electropositive metal through external circuit.
Transformation of energy	Electrical energy to chemical energy.	Chemical energy to electrical energy.
At positive terminal	At anode Oxidation occurs. Anions release electrons at the anode.	At cathode Reduction occurs. Oxidizing agent gain electrons.
At negative terminal	At cathode Reduction occurs. Cations gain electrons from the cathode.	At anode Oxidation occurs. Reducing agent releases electrons

Cell Potential

The oxidation of Zn(s) into Zn²⁺ and the reduction of Cu²⁺ to Cu(s) occur spontaneously. In other words, the redox reaction between Zn and Cu²⁺ is spontaneous. This is caused by the difference in potential energy between the two substances. The difference in potential energy between the anode and cathode dictates the direction of electrons movement. Electrons move from area of higher potential energy to area of lower potential energy. In this case, the anode has a higher potential energy so electrons move from anode to cathode. The potential difference between the two electrodes is measured in units of volts. One volt (V) is the potential difference necessary to generate a charge of 1 coulomb (C) from 1 Joule (J) of energy.

For a voltaic cell, this potential difference is called the cell potential, and is denoted E_cell. For a spontaneous reaction, E_cell is positive and ΔG (Gibbs free energy that can be  used to determine if a reaction occurs spontaneously) is negative. Thus, when ΔG  is negative the reaction is spontaneous. By merging electrochemistry with thermodynamics we get this formula: ΔG= −n× F × E_cell(or EMF). EMF= Electromotive Force. Cell potential is different for each voltaic cell, its value depends upon the concentrations of specific reactants and products as well as temperature of the reaction.

For standard cell potential, temperature of the reaction is assumed to be at 25^o Celsius, the concentration of the reactants and products is 1M and reaction occurs at 1 atm pressure. The standard cell potential is denoted E^o_cell. For Voltaic cells it will be E^o_cell=E^o(cathode)-E^o(Anode)

Example 1: Calculate E^o_cell for the following redox reaction:

  2Al_(s) + 3Sn²⁺_(aq) →  2Al³⁺_(aq) + 3Sn(s)

Solution:

Oxidation:{Al_(s) → Al³⁺_(aq) +3e^–} x 2                           -E^o = +1.676V

Reduction:{Sn²⁺_(aq) +2e^– → Sn_(s)} x 3                         E^o = -0.137V

 Net:2Al_(s) + 3Sn²⁺_(aq) →  2Al³⁺_(aq) + 3Sn_(s)                E^o_cell = -0.137V – (-1.676V)

 E^o_cell= +1.539 V.

Example 2: Given the redox reaction

Fe³⁺ + V²⁺ → Fe²⁺ + V³⁺

Calculate  Fe³⁺(aq)+e^– → Fe²⁺(aq) =0.771V,  V²⁺(aq)+ e^– → V³⁺(aq)=-0.255V

Solution:  E^o_cell = E_cathode-E_anode

              =0.771-(-0.255)=1.026V

Gibbs Energy and Cell potential

Relating the electrochemistry ideas with the thermodynamic parameter called Gibbs’ Free Energy.

w = -QE                       (work = – charge times potential),

where Q is the charge and can be defined as Q = nF where F is the Faraday constant (96500 C/mol), named after Michael Faraday.

We’re interested in the maximum work since this can be related to the thermodynamic parameter ΔG.

Thus, for the case where the work is done infinitely slowly (chemical system is always at equilibrium with the electrodes, electrical resistance is zero since current is essentially zero, etc…) we have
w_max = – QE_max where E_max for standard conditions is simply E°

Example 3: A cell with a maximum cell potential of 2.50 V. If 1.33 mol of e^– passes through the cell at an average potential E = 2.10 V. What is the efficiency?

Solution:

w = – Q E = – nFE. = 1.33 mol × 96500 C/mol × 2.10 V              (V = J/C)

w = – 2.69 × 10⁵ J = – 269 kJ

w_max = – nFE_max = 1.33 mol × 96500 C/mol × 2.50 V

        = – 321 kJ

Efficiency = w/w_max × 100% = -269/-321 × 100% = 83.8 %

 Since w_max is only achievable if the work is done reversibly (infinitely slowly), we can never reach 100% efficiency in any system in the real world.

We already have seen that ΔG is a measure of the maximum work obtainable from a system. Thus,

ΔG = w_max

ΔG = – Q E

ΔG = – nFE. In this case, the potential is the cell potential E_cell.

ΔG° = – RT lnK = – nFE°_cell.

Example 4: Is Fe²⁺ spontaneously oxidized by the oxygen of the air in acidic solution? Calculate ΔG and K.

EXERCISES

Lets see how much you’ve learnt, attach the following answers to the comment below

1. Potential difference set up between metal and its solution is called (a) metal voltage (b) back e.m.f. (c) electrode potential (d) electrode potential of the metal
2. Two half cells which are capable of converting chemical energy to electrical energy is called (a) a cell (b) electrochemical cell (c) chemical potential (d) metallic potential
3. Potential difference set up when a metal is in contact with one molar solution of its ions at 25⁰C is called (a) inert standard potential (b) standard electrode potential (c) electrochemical cell (d) galvanic cell
4. These are factors affecting standard electrode potential except (a) pressure (b) overall energy change (c) the concentration of ions in the solution (d) temperature
5. The overall redox reactions occurring at the electrodes is represented as Zn_(s) │Zn²⁺_(aq) ║ Cu²⁺_(aq) │ Cu_(s), the double represents (a) capacitor (b) battery (c) salt bridge (d) inert conductor

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