Electrochemistry Review - Cell Potential & Notation, Redox Half Reactions, Nernst Equation

This paragraph introduces the topic of electrochemistry, specifically focusing on the voltaic cell. It explains how the voltaic cell works, including the concepts of balancing equations under acidic and basic conditions, identifying oxidizing and reducing agents, and calculating cell potential under standard and non-standard conditions. Additionally, it touches on calculating Gibbs free energy and the equilibrium constant K, as well as discussing conceptual examples, electrolysis problems, and stoichiometry related to current and mass deposition on the cathode.

The second paragraph delves deeper into the workings of a voltaic cell, using the example of a zinc-copper cell. It explains the standard reduction potentials for zinc and copper, and how these values determine whether a cell is spontaneous or not. The paragraph clarifies the difference between a voltaic (galvanic) cell and an electrolytic cell, and how the cell potential is related to the energy produced. It also describes the half-reactions occurring at each electrode, the flow of electrons, and the roles of the anode and cathode in the cell's operation.

This paragraph discusses the purpose of the salt bridge in maintaining charge balance within an electrochemical cell. It explains how ions travel through the salt bridge to compensate for the charge separation caused by the flow of electrons. The paragraph clarifies the direction of cation and anion movement, with cations moving towards the cathode and anions moving towards the anode. It also touches on the choice of electrolyte for the zinc electrode, highlighting why zinc sulfate is the most suitable option.

The fourth paragraph focuses on identifying the oxidizing and reducing agents in a redox reaction. It explains the changes in oxidation states and how they relate to oxidation and reduction processes. The paragraph also clarifies the roles of the reducing agent (which causes reduction in another substance) and the oxidizing agent (which causes oxidation in another substance). It emphasizes that these agents are reactants and can be determined by looking at the changes in oxidation states.

This paragraph explains how to calculate the Gibbs free energy (ΔG) from the cell potential and how this relates to the spontaneity of a reaction. It introduces the concept of the equilibrium constant (K) and its relation to the concentration of products and reactants. The paragraph also discusses how changes in cell potential and ΔG can indicate whether a reaction is spontaneous or non-spontaneous, and how this ties into the values of K. It concludes with a worked example of calculating K for the given reaction.

The sixth paragraph instructs on how to write the cell notation for a given reaction, starting with the oxidative half-reaction and including the salt bridge. It emphasizes the importance of representing the reduction half-reaction on the other side. The paragraph then discusses how to calculate the cell potential under non-standard conditions using the Nernst equation, and how changes in the concentration of reactants and products affect the cell potential. It also explains how the cell potential is related to the spontaneity of the reaction and the direction in which the reaction proceeds.

The seventh paragraph presents stoichiometry problems related to electrochemistry, involving calculations of charge, moles of electrons, and mass of metals deposited or plated onto electrodes. It explains the relationship between current, time, and charge, and how to use Faraday's constant to convert between these quantities. The paragraph provides examples of calculating the mass of copper deposited on the cathode and the average current passing through a solution, highlighting the importance of understanding the molar ratios and the direction of electron flow in redox reactions.

This paragraph focuses on calculating the time required for electrodeposition of a specific mass of a metal, using the given current and solution. It explains the process of converting mass into moles, then into moles of electrons, and finally into the time required for the deposition to occur. The paragraph provides a worked example of calculating the time needed to deposit a certain mass of chromium onto the cathode, emphasizing the importance of understanding the relationships between mass, moles, and time in electrochemical reactions.

The ninth paragraph discusses how to identify the strongest reducing and oxidizing agents from a list of species, based on their standard reduction potentials. It explains that the strongest reducing agent is the metal with the most positive reduction potential, and the strongest oxidizing agent is the metal ion with the least negative (or most positive) reduction potential. The paragraph provides examples of such identifications for both metals and non-metals, highlighting the relationship between the cell potential and the tendency of species to gain or lose electrons.

The eleventh paragraph explains the process of balancing redox reactions and writing cell notation, including how to reverse half-reactions and balance the number of electrons. It provides examples of balancing reactions involving different metal ions and explains how to determine which reactions occur at the anode and cathode based on the electron flow and standard reduction potentials. The paragraph also discusses how to write the cell notation, including the representation of solids, gases, and aqueous phase species.

The final paragraph discusses how to balance redox reactions under acidic and basic conditions, providing examples of how to add water or hydroxide ions to balance the charges. It explains how to determine the predominant reactions occurring at the anode and cathode of an electrochemical cell, based on the standard reduction potentials and the species present in the solution. The paragraph concludes with a discussion on the relative likelihood of different reactions occurring, and how the cell potential influences which reactions are more favorable.