ELECTROCHEMISTRY

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21 Terms

1
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What is electrochemistry?

Electrochemistry is the area of chemistry concerned with the interconversion of chemical energy and electrical energy. It has broad significance and applications.

2
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What are electrochemical cells and their main types?

An electrochemical cell is a device that converts electrical energy into chemical energy or vice-versa. The two main types are Galvanic (or Voltaic) cells and Electrolytic cells.

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Describe Galvanic (Voltaic) cells.

Definition & Energy Conversion: Galvanic cells produce electricity from a spontaneous reaction, converting chemical energy into electrical energy.

    ◦ Process Nature: The reactions are spontaneous.

    ◦ Electrodes: Oxidation occurs at the anode, and reduction occurs at the cathode. In a Galvanic cell, the anode is typically the negative electrode (negative potential) and the cathode is the positive electrode (positive potential).

    ◦ Mnemonic for Anode: "ALOK" (A-node L-oss O-xidation).

    ◦ Electron and Current Flow: Electrons flow from the anode (negative) to the cathode (positive). Current flows in the opposite direction (from cathode to anode).

    ◦ Overall Reaction: A redox reaction.

    ◦ Cell Representation: Standard convention places the anode on the left and cathode on the right, using single vertical lines for phase boundaries and double vertical lines for the salt bridge . A mnemonic for representation is the "ABC trick" (A=Anode, B=Bridge, C=Cathode).

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Explain the Daniel Cell as an example of a Galvanic cell.

 ◦ Components: A classic example of a Galvanic cell, composed of zinc and copper.

    ◦ Half-Cell Reactions:

         Oxidation (Anode): Zn → Zn²⁺ + 2e⁻

         Reduction (Cathode): Cu²⁺ + 2e⁻ → Cu

    ◦ Overall Reaction: Zn + Cu²⁺ → Zn²⁺ + Cu

    ◦ Standard Cell Potential (E°): 1.1 V.

    ◦ Conditions/External Potential Effects: The cell's function changes based on an external opposing potential relative to 1.1 V: E° > 1.1 V (electrolytic), E° < 1.1 V (Daniel cell discharging), E° = 1.1 V (neutral/equilibrium).

    ◦ Electrode Thickness Changes: The zinc anode decreases in thickness, while the copper cathode increases in thickness.

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Describe Electrolytic Cells.

Definition & Energy Conversion: An electrolytic cell is one where a non-spontaneous reaction is driven by an external current source, converting electrical energy into chemical energy.

    ◦ Process: This process is called electrolysis.

    ◦ Process Nature: Reactions are non-spontaneous and require an external energy source.

    ◦ Contrast with Galvanic Cells: Electrolytic cells operate in the opposite manner to Galvanic cells. In electrolytic cells, the anode is positive and the cathode is negative.

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What is a salt bridge and what are its functions?

  ◦ Definition: A salt bridge is a U-shaped inverted tube containing a gel with an inert electrolyte (e.g., KCl).

    ◦ Functions: Its significance includes:

         Completing the circuit.

         Minimizing liquid junction potential.

         Maintaining electrical neutrality.

    ◦ Note: A salt bridge is not required if both electrodes dip in the same electrolyte solution .

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What is electrode potential, and what factors affect it?

 ◦ Definition: Electrode potential is the potential difference developed between a metal electrode and its ions in solution.

    ◦ Types: It includes oxidation potential and reduction potential.

    ◦ Relationship: Oxidation Potential = - (Reduction Potential).

    ◦ Factors Affecting: Concentration, nature of the metal, pressure/temperature, and the electrolyte.

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Define Standard Electrode Potential (E°) and how is E°_cell calculated?

 ◦ Definition: E° is the potential difference at standard conditions (1 M concentration, 1 bar pressure, and typically 298K).

    ◦ Calculation of E°_cell: The common formula is E°_cell = E°_cathode (Reduction Potential) - E°_anode (Reduction Potential). An alternative method uses process potentials (oxidation for anode, reduction for cathode).

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What is the Standard Hydrogen Electrode (SHE)?

   ◦ Purpose: SHE is the standard reference electrode used because the potential of a single electrode cannot be measured independently.

    ◦ Assumed Potential: Its electrode potential is arbitrarily assumed to be 0.00 V.

    ◦ Components & Reactions: Involves hydrogen gas at 1 bar pressure bubbled over a platinum electrode immersed in a 1 M H⁺ solution. Both oxidation and reduction reactions are described.

    ◦ Measurement: SHE is used to measure the standard electrode potential of other half-cells, with examples provided for Cu and Zn .

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How is Gibbs Free Energy (ΔG) related to cell potential (E_cell)?

 ◦ Formula: The fundamental formula is ΔG = -nFE_cell.

    ◦ Standard Conditions: For standard conditions, ΔG° = -nFE°_cell.

    ◦ Additivity: ΔG values are additive for combined half-reactions, unlike E_cell values.

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What is the Nernst Equation and what does it tell us about equilibrium?

Purpose: The Nernst equation relates cell potential under non-standard conditions to standard cell potential and concentrations (or activities) of reactants and products.

    ◦ Formula (at 298K): E_cell = E°_cell - (0.059 / n) log Q (where Q is the reaction quotient, and n is the number of electrons transferred). A detailed derivation for the Daniel cell is provided in one source .

    ◦ At Equilibrium: At equilibrium, ΔG = 0 and E_cell = 0. The reaction quotient Q becomes the equilibrium constant (K_eq).

    ◦ Relation to K_eq: log K_eq = (nE°_cell) / 0.059 (at 298K).

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Describe the two types of conductors.

Metallic (Electronic) Conductors:

         Charge Carriers: Electrons.

         Chemical Change/Mass Transfer: No chemical change or mass transfer occurs.

         Effect of Temperature: Conductivity decreases with increasing temperature.

    ◦ Electrolytic (Ionic) Conductors:

         Charge Carriers: Ions.

         Chemical Change/Mass Transfer: Decomposition of electrolyte and mass transfer occurs.

         Effect of Temperature: Conductivity increases with increasing temperature.

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Define electrical resistance, resistivity, conductance, and conductivity, including their units and related formulas.

Electrical Resistance (R): The opposition to the flow of electric current. Unit: Ohm (Ω). Formula: R = ρ * (L/A).

    ◦ Resistivity (ρ): The resistance of a conductor of unit length and unit cross-sectional area. Unit: Ohm-meter (Ω·m). Formula: ρ = R * (A/L).

    ◦ Conductance (G): The reciprocal of resistance. Unit: Siemens (S) or Ohm⁻¹. Formula: G = 1/R.

    ◦ Conductivity (κ): The reciprocal of resistivity. Unit: Siemens/meter (S/m) or Siemens/cm. Formula: κ = 1/ρ.

    ◦ or L/A):* The ratio of the distance between the electrodes (L) to the area of cross-section (A) of the electrodes . It relates conductivity to conductance: κ = G G .

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What is Molar Conductivity (Λm)?

Definition: Molar conductivity is the conductivity of a solution containing one mole of electrolyte.

    ◦ Formula: Λm = (κ * 1000) / C (where κ is conductivity in S cm⁻¹, C is concentration in mol L⁻¹).

    ◦ Unit: Siemens cm² mol⁻¹ (S cm² mol⁻¹).

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How does molar conductivity vary with concentration for strong and weak electrolytes?

   ◦ General Trend: Molar conductivity (Λm) increases upon dilution (or decreases with increasing concentration).

    ◦ Strong Electrolytes:

         Reason: Strong electrolytes are completely dissociated. The increase in Λm upon dilution is due to a decrease in inter-ionic interactions, allowing ions to move more freely.

         Trend: The increase is gradual . This behavior is described by the Debye-Hückel-Onsager equation .

    ◦ Weak Electrolytes:

         Reason: Weak electrolytes do not dissociate completely. Upon dilution, the degree of dissociation increases significantly, leading to a dramatic increase in the number of ions and thus molar conductivity.

         Trend: The increase is very steep at lower concentrations .

    ◦ Limiting Molar Conductivity (Λm°): Defined as the molar conductivity at infinite dilution (zero concentration). For strong electrolytes, it can be found by extrapolation, but for weak electrolytes, direct extrapolation is not possible.

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State Kohlrausch's Law of Independent Migration of Ions and its applications.

   ◦ Statement: At infinite dilution, each ion makes a definite contribution to the total molar conductivity of an electrolyte, irrespective of the nature of the other ion.

    ◦ Formula: Λm° = ν⁺λ°⁺ + ν⁻λ°⁻ (where ν⁺ and ν⁻ are the number of cations and anions, and λ°⁺ and λ°⁻ are their limiting molar conductivities).

    ◦ Applications: It has several applications, including:

         Calculating Λm° for weak electrolytes (by combining the Λm° values of strong electrolytes).

         Determining the degree of dissociation (α) of a weak electrolyte using the formula: α = Λm / Λm°.

         Determining the dissociation constant (Ka) of a weak electrolyte using the formula: Ka = Cα² / (1 - α).

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State Faraday's Laws of Electrolysis.

The mass (w) of a substance produced at an electrode is directly proportional to the quantity of electricity (Q) passed through the electrolyte.

         Formula: w = (Molar Mass / n-factor) (I t) / 96500. (Where n-factor is the number of electrons involved, I is current, t is time, and 96500 C/mol is Faraday's constant).

    ◦ Second Law: When the same quantity of electricity is passed through different electrolytic cells connected in series, the masses of substances produced are directly proportional to their equivalent masses.

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What are batteries, and what are the main types?

  ◦ Definition: A battery is two or more electrochemical cells connected in series.

    ◦ Types: Batteries are broadly categorized into three types:

         Primary Batteries (Non-Rechargeable): These cannot be recharged. Examples include the Dry Cell (Leclanché Cell) (carbon rod cathode, zinc container anode, MnO₂, carbon powder, NH₄Cl, ZnCl₂ electrolyte; approx. 1.5V) and the Mercury Cell (mercury-zinc amalgam anode, HgO-carbon paste cathode, KOH-ZnO electrolyte; approx. 1.35V).

         Secondary Batteries (Rechargeable): These can be recharged. Examples include the Lead Storage Cell (lead anode, lead dioxide cathode, H₂SO₄ electrolyte; approx. 2V per cell) and the Nickel-Cadmium Cell (cadmium anode, nickel dioxide cathode, KOH electrolyte; approx. 1.4V).

         Fuel Cells: These are Galvanic cells that convert the energy of combustion of fuels directly into electrical energy. They emphasize continuous operation and have high efficiency compared to thermal plants. A detailed example is the Hydrogen-Oxygen Fuel Cell (H₂ anode, O₂ cathode, porous carbon electrodes, hot KOH electrolyte; reactions yield water; used in applications like Apollo space missions).

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What is corrosion, and how does rusting occur?

 ◦ Definition: Corrosion is the deterioration of a metal due to its reaction with the environment.

    ◦ Nature: It is an electrochemical phenomenon.

    ◦ Rusting of Iron (Primary Example): Rusting is the formation of hydrated ferric oxide (Fe₂O₃·xH₂O).

         Anodic Reaction (Oxidation of Iron): Fe(s) → Fe²⁺(aq) + 2e⁻.

         Cathodic Reaction (Reduction of Oxygen): O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) (occurring in the presence of H⁺, often from CO₂ dissolving in water).

         The Fe²⁺ ions are then further oxidized by atmospheric oxygen to Fe³⁺, which precipitates as hydrated ferric oxide (rust).

    ◦ Other Examples: Tarnishing of silver and the green coating on copper/bronze are also forms of corrosion .

    ◦ Prevention: Methods briefly discussed include painting, covering with other metals, and using sacrificial electrodes .

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What is a "Concentration Cell"?

   ◦ A concentration cell is a specific type of electrochemical cell where the half-cells have identical electrodes but different electrolyte concentrations.

    ◦ For such cells, the standard cell potential (E°_cell) is zero.

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Briefly explain "The Hydrogen Economy".

 ◦ "The Hydrogen Economy" is a concept discussing hydrogen as a clean and renewable energy source .

    ◦ Electrochemistry plays a crucial role in both the production of hydrogen (e.g., through electrolysis of water) and its subsequent use in fuel cells for energy generation .