Electrochemistry - Galvanic cell

• Electrochemistry is a branch of chemistry that deals with the study of the conversion between electrical and chemical energy.

• A galvanic cell is an electrochemical cell that converts chemical energy into electrical energy.

• It consists of two half-cells, each containing an electrode and an electrolyte solution.

• The electrode at which oxidation occurs is called the anode, while the electrode at which reduction occurs is called the cathode.

• The anode is negatively charged, while the cathode is positively charged.

• In a galvanic cell, oxidation takes place at the anode.

• The electrons released during oxidation flow through an external circuit towards the cathode.

• Reduction takes place at the cathode, where the electrons from the anode combine with the ions from the electrolyte solution.

• This creates a flow of electrons from the anode to the cathode, which generates an electric current.

• The flow of electrons creates a potential difference between the anode and the cathode.

• This potential difference is commonly referred to as the cell potential or electromotive force (EMF).

• The cell potential is measured in volts (V) and can be determined using the Nernst equation.

• The Nernst equation relates the cell potential to the concentrations of the reactants and products involved in the redox reaction.

• The cell potential can be calculated using the equation:

  • Ecell = Ecathode - Eanode

• Ecell represents the cell potential.

• Ecathode is the reduction potential of the cathode.

• Eanode is the oxidation potential of the anode.

• The cell potential determines the feasibility of a redox reaction.

• The standard electrode potential (E°) is the potential difference between a half-reaction and the standard hydrogen electrode (SHE).

• The SHE has a potential of 0 volts and is used as a reference electrode.

• The standard electrode potential is a measure of the tendency of a species to gain or lose electrons.

• It is commonly used to compare the reactivity of different substances.

• The standard electrode potential can be positive, negative, or zero.

• A positive standard electrode potential indicates that the species has a greater tendency to gain electrons and undergo reduction.

• A negative standard electrode potential indicates that the species has a greater tendency to lose electrons and undergo oxidation.

• Zero standard electrode potential indicates that the species is in equilibrium with the SHE.

• The cell potential and standard electrode potentials can be used to predict the direction of electron flow in a galvanic cell.

• The reaction with the higher reduction potential will occur at the cathode, while the reaction with the lower reduction potential will occur at the anode.

• The electrons flow from the anode to the cathode and the cell potential is positive.

• This flow of electrons generates an electric current.

• The cell potential and the equilibrium constant (K) of the reaction are related by the equation:

• ΔG° = -nFE°cell

• ΔG° represents the change in Gibbs free energy.

• n represents the number of electrons transferred in the balanced redox reaction.

• F is Faraday’s constant (96500 C/mol).

• E°cell is the standard cell potential.

• If ΔG° is negative, the reaction is spontaneous and the cell potential is positive.

• If ΔG° is positive, the reaction is non-spontaneous and the cell potential is negative.

• If ΔG° is zero, the reaction is in equilibrium and the cell potential is zero.

• In summary, a galvanic cell converts chemical energy into electrical energy through redox reactions.

• The anode is the site of oxidation, while the cathode is the site of reduction.

• The cell potential is determined by the difference in reduction potentials between the cathode and anode.

• The standard electrode potential is a measure of the reactivity of a species.

• The cell potential and the equilibrium constant are related by the Gibbs free energy equation.

Oxidation and Reduction Reactions

• Oxidation is the loss of electrons by a substance.
• Reduction is the gain of electrons by a substance.
• In a redox reaction, both oxidation and reduction occur simultaneously.
• The substance that gets oxidized is called the reducing agent.
• The substance that gets reduced is called the oxidizing agent.

Oxidation States

• Oxidation states, or oxidation numbers, are assigned to atoms in a compound to track the transfer of electrons.
• The oxidation state of an element in its elemental form is 0.
• The sum of oxidation numbers in a neutral compound is 0.
• The sum of oxidation numbers in an ion is equal to the charge of the ion.
• Common rules can be followed to assign oxidation states based on electronegativity and electron ownership in a bond. Example: In H₂O, oxygen has an oxidation state of -2 and hydrogen has an oxidation state of +1.

Half-Reactions

• In an electrochemical cell, the oxidation and reduction reactions occur in separate half-cells.
• A half-reaction is the oxidation or reduction part of a redox reaction.
• The half-reactions are balanced by adding electrons to balance the charges.
• The balanced half-reactions can be combined to form the overall redox reaction. Example: Oxidation half-reaction:

2Cl⁻ → Cl₂ + 2e⁻ Reduction half-reaction:

2H⁺ + 2e⁻ → H₂ Overall redox reaction:

2Cl⁻ + 2H⁺ → Cl₂ + H₂

Standard Hydrogen Electrode

• The standard hydrogen electrode (SHE) is used as a reference electrode in electrochemical measurements.
• It consists of a platinum electrode immersed in a solution with a hydrogen ion concentration of 1 M.
• The potential of the SHE is defined as 0 volts.
• The standard electrode potential of other half-reactions is measured relative to the SHE.

Standard Cell Potential

• The standard cell potential (E°cell) is the potential difference between the cathode and anode in a galvanic cell under standard conditions.
• It is a measure of the driving force behind the redox reaction.
• The standard cell potential is determined by the difference in the standard electrode potentials of the cathode and anode half-reactions.
• The cell potential can be positive, negative, or zero depending on the relative reactivity of the half-reactions.

Nernst Equation

• The Nernst equation calculates the cell potential under non-standard conditions.
• It relates the cell potential (Ecell) to the standard cell potential (E°cell), the concentration of reactants and products, and the gas constant (R) and temperature (T). Nernst equation: Ecell = E°cell - (RT/nF) * ln(Q) where Q is the reaction quotient, n is the number of electrons transferred, R is the gas constant (8.314 J/(mol·K)), and T is the temperature in Kelvin. Example: For the reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) The Nernst equation would be used to calculate the cell potential if the ionic concentrations are not standard.

Spontaneity of Redox Reactions

• The spontaneity of a redox reaction can be predicted using the standard cell potential (E°cell).
• If the cell potential is positive (E°cell > 0), the reaction is spontaneous and will occur as written.
• If the cell potential is negative (E°cell < 0), the reaction is non-spontaneous and will not occur as written.
• If the cell potential is zero (E°cell = 0), the reaction is at equilibrium and there is no net movement of electrons.

Relationship between Equilibrium Constant and Cell Potential

• The Gibbs free energy change (ΔG°) of a redox reaction is related to the cell potential (E°cell) and the equilibrium constant (K) by the equation: ΔG° = -nFE°cell where n is the number of electrons transferred and F is Faraday’s constant (96500 C/mol).
• If ΔG° is negative, the reaction is spontaneous and the cell potential is positive.
• If ΔG° is positive, the reaction is non-spontaneous and the cell potential is negative.
• If ΔG° is zero, the reaction is at equilibrium and the cell potential is zero.

Factors Affecting Cell Potential

• The cell potential can be affected by various factors, including temperature and concentration.
• Temperature: An increase in temperature generally increases the cell potential. The relationship between cell potential and temperature is determined by the temperature coefficient.
• Concentration: Changes in the concentration of reactants and products can affect the cell potential. The Nernst equation can be used to calculate the cell potential under non-standard concentrations. Example: For the reaction:

2Fe³⁺(aq) + 3Sn²⁺(aq) → 2Fe²⁺(aq) + 3Sn⁴⁺(aq) Increasing the concentration of Sn²⁺(aq) will increase the cell potential.

Applications of Galvanic Cells

• Galvanic cells have various practical applications:
  • Batteries: Galvanic cells are commonly used in batteries to provide electrical energy.
  • Corrosion prevention: Galvanic cells can be used to prevent corrosion by intentionally sacrificing a less reactive metal.
  • Electroplating: Galvanic cells can be used in the electroplating process to deposit a layer of metal onto a surface.
• Understanding galvanic cells and their applications is crucial in various fields, including electrochemistry, materials science, and energy storage.

Slide 21

• Electrolytic Cells
  • Electrolytic cells are electrochemical cells that use electrical energy to drive non-spontaneous redox reactions.
  • Unlike galvanic cells, electrolytic cells do not generate electrical energy. Instead, they require an external power source.
  • Electrolysis is the process of using an electrolytic cell to decompose a compound into its constituent elements or ions.

Slide 22

• Electrolysis Process
  • In an electrolytic cell, the anode is positively charged, while the cathode is negatively charged.
  • The anode attracts negatively charged ions (anions), which undergo oxidation and lose electrons.
  • The cathode attracts positively charged ions (cations), which undergo reduction and gain electrons.
  • The electrons flow from the external power source to the cathode and back to the anode, completing the circuit.

Slide 23

• Faraday’s Laws of Electrolysis
  • Faraday’s First Law: The amount of substance produced or consumed during electrolysis is directly proportional to the quantity of electricity passed through the cell.
    • Example: 1 mole of electrons corresponds to the production of 1 mole of copper metal during the electrolysis of copper(II) sulfate.
  • Faraday’s Second Law: The amount of different substances produced or consumed during electrolysis is directly proportional to the ratio of their stoichiometric coefficients in the balanced chemical equation.
    • Example: In the electrolysis of water, for every 2 moles of hydrogen gas produced, 1 mole of oxygen gas is produced.

Slide 24

• Electroplating
  • Electroplating is the process of depositing a layer of metal onto a substrate using electrolysis.
  • The object to be electroplated is used as the cathode in the electrolytic cell.
  • The metal to be deposited is used as the anode in the electrolytic cell.
  • The electrolyte solution contains ions of the metal to be deposited.

Slide 25

• Electroplating Process
  1. The object to be electroplated is cleaned and made the cathode in the electrolytic cell.
  2. The metal to be deposited is made the anode in the electrolytic cell.
  3. The electrolyte solution contains ions of the metal to be deposited.
  4. When a current is passed through the cell, metal ions from the electrolyte are reduced at the cathode and deposited onto the object to be electroplated.

Slide 26

• Applications of Electroplating
  • Decorative purposes: Electroplating is commonly used to give a shiny and attractive appearance to objects such as jewelry, silverware, and car parts.
  • Protection against corrosion: Electroplating can be used to provide a protective layer of metal onto objects to prevent corrosion.
  • Enhancement of properties: Electroplating can be used to improve the wear resistance, hardness, or conductivity of objects.

Slide 27

• Electrochemical Cells and Batteries
  • Electrochemical cells are devices that convert chemical energy into electrical energy by redox reactions.
  • Batteries are a type of electrochemical cell that store electrical energy and supply it on demand.
  • Common types of batteries include the alkaline battery, lead-acid battery, and lithium-ion battery.

Slide 28

• Alkaline Battery
  • An alkaline battery consists of a zinc anode, a manganese dioxide cathode, and an alkaline electrolyte such as potassium hydroxide.
  • The anode undergoes oxidation, while the cathode undergoes reduction.
  • The chemical reaction between the anode and cathode produces an electric potential difference (voltage).
  • Alkaline batteries are commonly used in portable electronic devices and toys.

Slide 29

• Lead-Acid Battery
  • A lead-acid battery consists of lead anode, lead dioxide cathode, and a mixture of sulfuric acid and water as the electrolyte.
  • The anode and cathode undergo redox reactions to produce electrical energy.
  • Lead-acid batteries are commonly used in automobiles, uninterruptible power supplies (UPS), and forklifts.

Slide 30

• Lithium-Ion Battery
  • A lithium-ion battery consists of a lithium cobalt oxide or lithium iron phosphate cathode, a graphite anode, and a lithium salt electrolyte.
  • Lithium ions migrate between the anode and cathode during charge and discharge cycles.
  • Lithium-ion batteries are widely used in portable electronic devices, electric vehicles, and renewable energy systems.