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Chapter 15. Electrochemistry

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1. Basic Theory

2. Electrochemical Cells

3. Galvanic Cells

4. Electrolysis


a. Electrochemistry Problems

b. Electrochemistry Reaction Kinetics and Batteries



1. Basic Theory

⑴ Electrical Work

① n : Mole number of electrons

② (Note) Understanding is not difficult if you know that the unit of E is V (volts) ≡ J / C

③ The change in Gibbs free energy associated with a process is the maximum reversible work the system can perform at constant temperature and pressure.

⑵ Faraday’s Law of Electrolysis

① n : Molar equivalent of electrons involved in half-reaction

② F : Faraday’s constant ≒ 96500 C/mol

③ W : Mass of deposited metal

④ M : Atomic equivalent of deposited metal

⑤ Charge of 1 mole of electrons = N₀e = 1F = 96485.33977 C/mol

○ 1 F = 6.02214179 × 10²³ entities/mol × 1.602176487 × 10⁻¹⁹ C/entity = 96485.33977 C/mol

⑥ Amount of charge transferred = Current (A) × Time (s)

⑶ Metal Ionization Tendencies

① K > Ca > Na > Mg > Al > Zn > Fe > Ni > Sn > Pb > H > Cu > Hg > Ag > Pt > Au

② Helps determine what gets oxidized and what gets reduced

Activity(activity) : Electrochemistry problems are conveniently dealt with using concentration, but strictly speaking, they should be calculated using activity.



**2. Electrochemical Cells **

⑴ Definitions

① Galvanic cell : An electrochemical cell that uses the potential of chemical reactions to provide voltage → spontaneous reaction

○ Characteristics : The oxidation electrode is the cathode (-), and the reduction electrode is the anode (+).

○ Galvanic cell acts as its own power source, supplying electrons from the oxidation electrode and recovering electrons at the reduction electrode.

② Electrolytic cell : Electrons move against the potential due to an external power source → electrolysis

○ Characteristics : The reduction electrode is the cathode (-), and the oxidation electrode is the anode (+).

○ An external power source (a stronger cell) is connected to the cathode to forcibly supply electrons at the reduction electrode (-).

⑵ Four Major Components : Refers to the cathode, anode, electrolyte, and separator

① Cathode : The metal that receives electrons from an external wire and undergoes reduction.

② Anode : The metal that undergoes oxidation and releases electrons to an external wire.

③ Electrolyte : Substance that facilitates the movement of substances to ensure that the electrodes remain electrically neutral.

④ Separator : An insulating barrier that prevents direct physical contact between the cathode and anode.

⑤ Salt bridge : Without a salt bridge, as electrochemical reactions proceed, charges accumulate at the oxidation and reduction electrodes.

○ Accumulation of charge hinders additional reactions.

○ The salt bridge provides suitable ions that connect the reduction and oxidation electrodes to alleviate charge accumulation.

⑶ Cell Diagram

① Example : Zn(s) Zn²⁺(aq)   Cu²⁺(aq) Cu(s), Ecellº
: Phase boundary
  : Salt bridge
④ “Zn( s ) Zn²⁺( aq )” : Oxidation electrode
⑤ “Cu²⁺( aq ) Cu( s )” : Reduction electrode

⑥ Ecellº = Erightº - Eleftº (Where Eº is the standard reduction potential)

⑦ Ecellº > 0 indicates a spontaneous reaction.



3. Galvanic Cells

⑴ Standard Electrodes and Standard Potentials

① Reference electrode

○ Necessity : The absolute value of electrode potential cannot be measured; only the potential difference between two electrodes can be measured.

○ Condition : The reference electrode must exhibit ideally non-polarizable electrode characteristics.

Type 1: Standard hydrogen electrode (SHE)

○ Half-cell composed of 1 atm of hydrogen gas in contact with a hydrogen ion solution with an activity of 1 at 25 ℃.

○ Platinum (Pt) is used as a simple electron conductor without participating in the reaction.

Type 2: Normal hydrogen electrode (NHE)

○ Half-cell composed of 1 atm of hydrogen gas in contact with a 1 M H+ solution at 25 ℃.

Type 3: Saturated calomel electrode (SCE)

○ Maintains a saturated concentration of KCl.

Reason 1: To fix the electrode potential influenced by chloride ion concentration.

Reason 2: To make the potential similar to SHE.

Type 4: Silver/silver chloride electrode (Ag/AgCl electrode)

② Standard Electrode Potential (E°)

○ Standard reduction potential at 25 ℃ and 1 atm, with a solution concentration of 1 M in a half-cell, taking the standard hydrogen electrode as the cathode, and determining the standard reduction potential based on the reduction reaction.

○ If the standard reduction potential is (+), it is easier to reduce than hydrogen ions; if (-), it is harder to reduce than hydrogen ions.

○ The standard oxidation potential has the same absolute value as the standard reduction potential but with the opposite sign.

Table 1. Standard Reduction Potentials Table

③ Calculation of Standard EMF : Given two half-reactions, E1° for the oxidation half-reaction and E2° for the reduction half-reaction,

○ Standard EMF = Reduction half-reaction’s reduction potential + Oxidation half-reaction’s oxidation potential = E2° - E1°

○ If the standard EMF is positive, the forward reaction is spontaneous; if negative, the reverse reaction is spontaneous.

Tip: Complex electrode problems can be solved by considering adjacent oxidation and reduction half-reactions as a single cell.

Nernst Equation : Formula concerning actual potential

① Cell Voltage and Thermodynamics : For reaction quotient Q,

② Nernst Equation

③ Measurement of Equilibrium Constants using Electrochemical Devices

Definition of Equilibrium Constants

○ ΔG° < 0 ⇔ E° > 0 ⇔ K > 1

○ ΔG° < 0 ⇔ E° < 0 ⇔ K < 1

④ Eeq and E°

○ Nernst equation can also be applied to half-cells

○ Considering the form An+ + ne- → A, where the degree of freedom of electrons is 1 and Q = [A] / [An+],

○ Prediction of oxidation/reduction tendencies for two half-cells based on Eeq and E° is different: Eeq must be followed.

○ Example: For [Fe2+] = [Fe3+] = 0.1 M and [Ag+] = 10⁻⁵ M, the Fe²⁺/Fe³⁺ half-cell undergoes reduction, and the Ag/Ag⁺ half-cell undergoes oxidation.

Application: Nernst Equation and Resting Membrane Potential

Electrode Reactions

① Reaction between Electrodes

Example 1: Construction of Reaction Index Formula

Example 2: Voltage Difference as Intensive Property: Subtract oxidation electrode potential from reduction electrode potential

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s), ΔG = -212 kJ, Ecellº = +1.10 eV

2Zn(s) + 2Cu²⁺(aq) → 2Zn²⁺(aq) + 2Cu(s), ΔG = -424 kJ, Ecellº = +1.10 eV

② Reaction within Electrodes

Example 1. Latimer Diagram : Calculation of reduction potentials in electrode reactions similar to the internal division formula in mathematics.

Example 2. When ligands form complexes

Phenomenon 1. As complex formation between metal ions and ligands increases (larger Kf), E°complex becomes more negative.

Phenomenon 2. As complex formation between metal ions and ligands increases, the potential decreases.

Phenomenon 2 and its relation to the decrease in [Au3+] leads to a decrease in the value of E = E° - (0.0592 / n) × log (1 / [Au3+]).

○ Note that the situation with E° = 1.50 V and Ecomplex° = 1.00 V is not the same.

Battery Impedance : Also known as EI (Electrochemical Impedance).

① Overview : Reaction kinetics in electrochemistry. In other words, the theory regarding how much current flows.

Elements of Battery Impedance

○ Electrolyte resistance (uncompensated resistance)

○ Double layer capacitance (Cdl)

○ Coating capacitance

○ Warburg impedance

○ Charge transfer resistance (Rct)

○ Constant phase element

○ Virtual inductor

Practical Batteries

Example 1. Fuel Cells : Production of electrical energy from continuously supplied fuel.

○ Hydrogen fuel cell

H2(g) + ½ O2(g) → H2O(ℓ), ΔH = -286 kJ

○ Methane fuel cell

CH4(g) + 2O2(g) → CO2(g) + 2H2O(ℓ), ΔH = 561 kJ

② Capacity of Practical Batteries (ampere-hour rating)

○ Since the voltage of each chemical cell is constant, it is expressed in terms of current (unit: Ah, mAh).

○ Battery lifespan

○ The capacity of chemical cells decreases with higher current and decreases at temperatures higher or lower than room temperature (24 ℃ = 75.2 ℉).

Figure 1. Battery capacity based on temperature, discharge time, and discharge current

⑹ Concentration Cell

① Definition : When the same chemical reaction is used for the oxidation and reduction electrodes.

② Entropy Interpretation : Electrodes with high metal ion concentration tend to reduce, while electrodes with low concentration tend to oxidize, increasing disorder.

③ Voltage Calculation : Usually takes the following form



4. Electrolysis

⑴ Electrolysis of Water : Standard conditions (pH 0)

① Anode Reaction (Oxidation Reaction)

○ Acidic conditions

○ Alkaline conditions

② Cathode Reaction (Reduction Reaction)

○ Acidic conditions

○ Alkaline conditions

③ Memorizing the electrolysis reactions of water is unnecessary; understanding them is important Reference

④ In reality, there’s a potential drop across the two electrodes, considering overpotential is necessary.

○ Reaction Example

○ Electrolysis occurs when slightly more than 1.23 V is applied, exceeding the potential.

⑵ Electrolysis of Solutions

① Electrolysis of Sodium Chloride (NaCl) Solution

○ Decomposition reaction of sodium chloride and water

○ Anode Reaction : Oxidation potential is high for water or hydroxide ions, but chlorine ions oxidize due to practical reasons.

○ Practical reasons : Reaction rate factors, experimental factors (e.g., overpotential)

○ Cathode Reaction : Reduction potential for hydrogen ions is higher than that of sodium ions, so hydrogen ions are reduced.

② Electrolysis of Copper Sulfate (CuSO4) Solution

○ Decomposition reaction of copper sulfate and water

○ Anode Reaction : Sulfate ions are difficult to oxidize, so hydroxide ions are oxidized.

○ Cathode Reaction : Copper ions readily accept electrons compared to hydrogen ions.

③ Corrosion : Undesired oxidation of metals

Figure 2. Corrosion of iron

○ Reaction equations

○ Reduction of water 1.

2H2O(ℓ) + 2e- → H2(g) + 2OH-(aq), Eº = -0.83 V

○ Reduction of water 2.

O2(g) + 4H+(aq) + 4e- → 2H2O(ℓ), PO2 = 0.2 bar, Eº = 1.23 V

○ Oxidation of iron 1.

Fe2+(aq) + 2e- → Fe, Eº = -0.44 V

○ Oxidation of iron 2.

Fe3+ + e- → Fe2+, Eº = 0.77 V

○ Anaerobic conditions

○ Reduction of water 2 doesn’t occur

○ At pH = 7, reduction of water 1 is E = -0.42 V

○ The potential difference between -0.42 V and -0.44 V is only 0.02 V, so corrosion of iron doesn’t occur vigorously.

○ Aerobic conditions

○ Reduction of water 2 occurs more readily than reduction of water 1

○ 1.23 V is greater than -0.44 V and 0.77 V, so both oxidation reactions of iron 1 and 2 occur

○ Corrosion of iron occurs vigorously, resulting in Fe3+ oxidation

Application 1. Corrosion generally takes a long time

Application 2. Measures to prevent corrosion

○ Coating: Applying paint as a coating

○ Galvanization: Plating with a metal that ionizes easily (the plated metal oxidizes instead, extending the time)

○ Sacrificial anode (cathodic protection)

⑶ Selecting Electrolysis Reactions

Figure 3. Examples of electrolysis reactions

① Reactants : Ni2+, Zn2+, H2O, SO42-

② Oxidation Reactants : H2O (-1.23 V), SO42- (-2.01 V)

○ Oxidation priority : H2O (-1.23 V) > SO42- (-2.01 V)

○ Oxidation reaction chosen : H2O

③ Reduction Reactants : Ni2+ (-0.25 V), Zn2+ (-0.76 V), H2O (-0.83 V)

○ Reduction priority : Ni2+ (-0.25 V) > Zn2+ (-0.76 V) > H2O (-0.83 V)

○ Reduction reaction chosen : Ni2+

⑷ Applications of Electrolysis

① Electroplating : Applying a thin layer of another metal to the surface of a metal to prevent corrosion.

○ Anode (+) (Oxidation electrode): Connected to a solution of salt, including the object to be plated.

○ Cathode (-) (Reduction electrode): Connected to the material to be electroplated.

② Copper Refining : Purifying copper metal containing small amounts of impurities like Zn, Fe, Ag, Au, Pt to obtain pure copper metal.

○ Anode (+): Impure copper

○ Cathode (-): Pure copper electrode

○ Electrolyte: Solution containing copper ions

○ Zn and Fe precipitate at the cathode due to their high ionization tendency, while Ag, Au, Pt do not dissolve and settle to the bottom.

③ Sodium Hydroxide Production : Electrolysis of sodium chloride solution produces sodium hydroxide near the cathode.

○ Chlorine produced at the anode reacts with sodium hydroxide and requires separation.

④ Electrowinning of Noble Metals

○ Overview

○ Light noble metals (density less than 4), such as sodium, magnesium, and aluminum, have strong chemical bonds → Cannot be separated directly.

○ To obtain metals from their compounds, they’re first melted, then separated by electrolysis.

Example 1. Hall-Héroult Process: Magnesium refining

Example 2. Dau Process: Commercial process for separating magnesium from molten MgCl2

○ Anode (+) (Oxidation electrode)

2Cl-(melt) → Cl2(g) + 2e-

○ Cathode (-) (Reduction electrode)

Mg2+(melt) + 2e- → Mg(ℓ)



Input: 2018.12.28 15:01

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