Electrochemistry Overview

Electrochemistry Overview

• Electrochemistry is the study of the interrelationship between chemical reactions and the electrical energy that accompanies them.

Learning Objectives

• Define oxidation and reduction reactions
• Balance redox reactions using the change in oxidation number method
• Identify the reaction occurring in different parts of the cell
• Define reduction potential, oxidation potential, and cell potential
• Calculate the standard cell potential
• Relate cell potential to the feasibility of generating an electric current
• Describe electrochemistry in common batteries (Leclanche dry cell, button batteries, fuel cells, lead storage)
• Apply electrochemical principles to explain corrosion
• Explain electrode reactions during electrolysis
• Describe reactions in commercial electrolytic processes

Redox Reactions

• Redox reactions involve electron transfer.
• Oxidation is the loss of electrons.
• Reduction is the gain of electrons.
• The combustion of barium metal is a redox reaction (2Ba + O2 → 2BaO).
• The example shows the oxidation and reduction half-reactions in the reaction.

Oxidation Numbers

• Oxidation number indicates the number of electrons lost or gained by an atom in a molecule or ionic compound.
• Atoms in free elements have a zero oxidation number (e.g., He, N2, P4, S8).
• For a monatomic ion, the oxidation number equals the charge of the ion (e.g., Na+, Al3+, Hg₂²⁺, F⁻, S₂⁻).
• Oxygen's oxidation number is typically -2 in compounds, except in OF₂ (where it's +2) or peroxides (where it's -1).
• Hydrogen's oxidation number is +1 in compounds with nonmetals and -1 when bonded to metals.
• Fluorine in compounds has an oxidation number of -1. Other halogens are -1 in halide form but can have positive oxidation numbers when bonded with oxygen or a more electronegative element.
• In a compound, the sum of the oxidation numbers of all atoms is zero; in a polyatomic ion, the sum of the oxidation numbers equals the ion's charge.

Changes in Oxidation Numbers

• A reaction is considered redox when reactant atoms change their oxidation numbers.
• Oxidizing agents are reactants that cause oxidation, and they are reduced in the reaction.
• Reducing agents cause reduction and are oxidized in the reaction.
• A mnemonic to remember is LEORA-GEROA (Loss of Electrons is Oxidation, Reduction is Gain of Electrons, Oxidizing Agent is Reduced, Reducing Agent is Oxidized).

Types of Redox Reactions

• Redox reactions can be observed in combination (e.g., 2Mg + O₂ → 2MgO), decomposition (e.g., 2KClO₃ → 2KCl + 3O₂), combustion (e.g., CH₄ + 2O₂ → CO₂ + 2H₂O), metal displacement (e.g., CuSO₄ + Zn → Cu + ZnSO₄), hydrogen displacement, and halogen displacement reactions.

Balancing Redox Reactions

• Steps for balancing redox reactions involve:
  1. Write the unbalanced equation.
  2. Assign oxidation numbers to all elements.
  3. Draw lines connecting elements involved in oxidation and reduction.
  4. Determine multipliers that make increases and decreases equal.
  5. Perform balancing by inspection for other elements.

Electrochemical Cells

• Electrochemical cells create electric current from redox reactions and/or supply electrical energy for the reactions.
• Types include galvanic (voltaic) and electrolytic cells.

Galvanic Cells

• Galvanic cells produce electrical energy from a spontaneous redox reaction.
• Electrodes are surfaces for redox reactions, with the anode being where oxidation takes place and electrons originate and the cathode being where reduction occurs and electrons proceed..
• An anolyte is the solution next to the anode, and the catholyte is next to the cathode.
• Voltmeters are used to measure the potential differences.
• Wires, or external circuits, connect electrodes to a power/battery terminal.
• Salt bridges maintain electrical neutrality in the cell.

Half-Cell Reactions

• An electrochemical cell consists of two half-reactions (anode and cathode). The anode involves oxidation, and the cathode involves reduction.

Overall Cell Reactions

• The overall cell reactions can be determined by adding the anode and cathode half-reactions.

Examples of Galvanic Cells

• The Daniell cell utilizes zinc and copper electrodes immersed in ZnSO4 and CuSO4 solutions, respectively.
• The copper-silver cell uses copper and silver electrodes submerged in CuSO4.and AgNO3 solutions.

Electrode Potentials

• Electrode potential (Ered) represents the potential of reduction.
• Standard reduction potential (E^o_red), measured under standard conditions, indicates the tendency of a species to be reduced.
• High E^o_red suggests strong oxidizing potential.
• Conversely, the species with the lowest E^o_red has the highest reducing capacity.
• Standard cell potential (E^o_cell) equals the sum of the standard oxidation potential of the anode and the standard reduction potential of the cathode.
• The sign of E^o_cell determines spontaneity. A positive value indicates a spontaneous reaction.

Batteries

• Batteries are galvanic cells used to provide electrical energy.
  • Leclanché
  • Button
  • Lead Storage
  • Nickel-Cadmium
  • Lithium-ion

Corrosion

• Corrosion is the process where metals are destroyed by oxidation, frequently in presence of water and oxygen.
• Susceptible sites for rusting on iron (like a nail) are those experiencing mechanical stress.
• Factors hastening rust include electrolytes, acids, and lack of protective coatings (like paint).

###Electrolysis

• Electrolysis is a process that uses an electric current to drive a non-spontaneous redox reaction.
• Electrolytic cells are electrochemical cells used in electrolysis.
• Electrolysis of water decomposes water into hydrogen and oxygen gas.
• Electroplating is an electrolytic coating technique for metals.
• Electrorefining is an electrolytic metal purification process.