Chapter 15. Electrochemistry
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1. Basic Theory
4. Electrolysis
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
⑷ 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
② : Phase boundary
③ : 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.
⑵ 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
○ 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.
⑤ Application: Nernst Equation and Resting Membrane Potential
⑶ 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.
⑷ Battery Impedance : Also known as EI (Electrochemical Impedance).
① Overview : Reaction kinetics in electrochemistry. In other words, the theory regarding how much current flows.
○ Electrolyte resistance (uncompensated resistance)
○ Double layer capacitance (Cdl)
○ Coating capacitance
○ Warburg impedance
○ Charge transfer resistance (Rct)
○ Constant phase element
○ Virtual inductor
① 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