25. Oxidation-Reduction and Electrochemical Cells

The script begins with an introduction to MIT OpenCourseWare and a creative commons license. It then transitions into a chemistry lesson led by Catherine Drennan, focusing on oxidation-reduction (redox) reactions. The lecture includes a demo involving magnesium, which reacts with carbon dioxide to produce graphene and magnesium oxide. Definitions of oxidation, reduction, oxidizing and reducing agents are provided, and the importance of understanding these concepts in chemistry is emphasized.

This paragraph delves into the rules for assigning oxidation numbers to elements in compounds. It explains that free elements have an oxidation number of zero, monatomic ions equal their charge, and metals from Groups 1 and 2 have oxidation numbers of +1 and +2 respectively. The paragraph also covers the common oxidation states of oxygen and hydrogen, with exceptions noted for peroxides and superoxides. The importance of balancing charges in compounds and polyatomic ions is highlighted through examples like NH4+, where nitrogen has an oxidation number of -3.

The script continues with a discussion on oxidation numbers that are not integers, exemplified by oxygen in superoxides having an oxidation number of -1/2. It then introduces disproportionation reactions, where a single element is both oxidized and reduced, using the example of chlorine in various compounds. The audience is engaged in determining oxidation numbers and identifying the type of reactions taking place.

The lecture moves on to balancing redox reactions, especially in the context of electrochemistry. It outlines the steps for balancing half-reactions in acidic conditions, starting with balancing elements that are not oxygen or hydrogen, followed by balancing oxygen with water, hydrogen with H+, and finally balancing the charge with electrons. The process is illustrated with examples involving chromium and iron, highlighting the need to multiply the half-reactions to ensure the electrons cancel out when combined.

This section underscores the significance of redox reactions in energy production, mentioning processes like photosynthesis and fuel cells. It introduces electrochemistry as the study of redox reactions at electrodes and differentiates between galvanic cells, which generate electric current from spontaneous reactions, and electrolytic cells, which use electric current to drive non-spontaneous reactions. The lecture also explains the role of anodes and cathodes in electrochemical cells.

The script provides an overview of electrochemical cells, focusing on the function of anodes and cathodes. It describes an experiment involving a zinc anode and a copper cathode, detailing the oxidation at the anode and the reduction at the cathode. The importance of the salt bridge in maintaining charge balance is also discussed, along with the representation of electrochemical cells in shorthand notation.

The lecture incorporates interactive learning with clicker questions, allowing the audience to participate in identifying anode and cathode reactions and differentiating between oxidation and reduction processes. The use of zinc and tin in an electrochemical cell is explored, with the audience voting on the correct reactions and their types. The correct answers are revealed, and the notation for representing the electrochemical cell is discussed.

This section introduces Faraday's law, which relates the consumption of an anode and the deposition of metal at the cathode to the charge passing through the system. An example calculation is provided to demonstrate how much zinc is consumed and copper is deposited given a current of 1 amp for 1 hour. The lecture also touches on the practical applications of electrochemistry, such as electroplating, where a thin coating of copper is deposited onto an object like a steel spoon.

The script discusses the use of inert electrodes like platinum, which do not participate in the reaction but facilitate it. It explains that in some electrochemical cells, the anode and cathode reactions occur in solution without the deposition or consumption of the electrode material. The concept of a standard hydrogen electrode (SHE) is introduced as a reference point for measuring reduction potentials, and the audience is engaged to determine the reaction occurring at the hydrogen electrode when used as an anode.

The final paragraph wraps up the lecture by transitioning to the topic of cell potentials, setting the stage for the next lecture. It summarizes the covered material, including the reactions at anodes and cathodes and the notation for writing electrochemical cells. The potential of these cells is highlighted as an important aspect to be explored in subsequent lessons.