Electrode Potentials and Cells - Everything You NEED To Know｜AQA A Level Chemistry

The video begins with an introduction to electrode potential, highlighting the challenges associated with understanding the topic. The speaker acknowledges the complexity due to its association with redox reactions and memory-based problem-solving. The speaker also sets expectations, warning viewers with weak foundational knowledge in oxidation and reduction to skip the video. The aim is to cover the bulk of electrode potentials, including half-cells, without delving into commercial applications like rechargeable cells or fuel cells. The video is structured around the AQA A-Level Chemistry specification, emphasizing the importance of grasping electrode potentials for academic success.

This paragraph delves into the components and setup of an electrochemical cell, including the salt bridge, electrodes, and solutions within the beakers. The salt bridge, often a filter paper soaked in a saturated potassium nitrate solution, facilitates the flow of ions between the two beakers. The electrodes are solid metals that can facilitate redox reactions, and the solutions contain metal ions to support these reactions. The speaker also clarifies common misconceptions about the flow of electrons versus ions in the cell and emphasizes the importance of using a high-resistance voltmeter to measure the potential difference without allowing current to flow.

The speaker explains the fundamental role of redox reactions in cells, focusing on the movement of electrons. Oxidation involves the loss of electrons, while reduction involves their gain. These processes occur in half-cells, with electrons flowing from the more negative (oxidized) half-cell to the more positive (reduced) half-cell. The speaker introduces the NOPR rule, a mnemonic for remembering that the more negative half-cell is oxidized, and the more positive half-cell is reduced. This rule is crucial for predicting the direction of simple redox reactions and calculating the electromotive force (EMF) of a cell.

The speaker discusses the conventional cell diagrams used in chemistry, which simplify the representation of electrochemical cells. These diagrams include a salt bridge symbolized by two vertical lines, the most oxidized form of the species in the half-cell, and the electrode represented by a single vertical line indicating a phase change. The speaker also explains the IUPAC convention of placing the more positive half-cell on the right and the more negative half-cell on the left. The explanation is accompanied by an example of a zinc-copper cell diagram, illustrating the placement of species and the electrode.

The speaker clarifies that the electrode potential of a single half-cell cannot be measured in isolation; it is the potential difference between two half-cells that is determined. The standard hydrogen electrode (SHE) serves as a reference point with a potential of zero volts. It consists of a platinum electrode immersed in a solution of hydrogen ions at standard conditions (1 mole per decimeter cubed, 100 kilopascals pressure, and 298 Kelvin temperature). The SHE is used to compare and calculate the potential difference for other half-cells. The speaker also discusses the importance of standard conditions for maintaining constant equilibrium in electrochemical reactions.

The speaker introduces the electrochemical series, a table listing the standard electrode potentials of various half-cells. This series is crucial for predicting the feasibility and direction of redox reactions. The speaker explains that the forward direction of the half-equations in the series represents reduction, and the strength of oxidizing or reducing agents can be deduced from their positions in the series. The speaker emphasizes the importance of understanding the NOPR rule and the direction of electron flow in these reactions. The video aims to equip viewers with the knowledge to handle data tables and perform calculations involving electrode potentials.

The speaker provides a detailed explanation of how to perform calculations involving electrode potentials, using a zinc-copper cell as an example. The process involves identifying the reduced and oxidized half-cells, applying the NOPR rule, and using the equation E_cell = E_reduced - E_oxidized to calculate the cell potential. The speaker also discusses the practical aspects of electrochemical cells, such as the use of platinum electrodes and the conditions under which standard electrode potentials are measured. The emphasis is on understanding the principles behind the calculations and the significance of standard conditions in maintaining consistent results.

The speaker explores how changes in concentration and temperature affect the electrode potential of a cell. According to Le Chatelier's principle, increasing the concentration of a reactant increases the cell's electrode potential, while decreasing it has the opposite effect. Similarly, increasing the temperature of an exothermic cell decreases its electrode potential, as the equilibrium shifts to oppose the change. The speaker reinforces the importance of standard conditions in maintaining equilibrium and ensuring consistent electrode potentials. The video concludes with a series of practice questions to help viewers apply and reinforce their understanding of electrode potentials.

The speaker demonstrates how to construct a full redox reaction from two half-equations, using the reaction between fluorine and water as an example. The process involves identifying the reduced and oxidized species from their electrode potentials, manipulating the half-equations to form a full reaction, and ensuring the conservation of electrons. The speaker emphasizes the importance of balancing the number of electrons in both half-equations and correctly representing the species involved in the reaction. The explanation is designed to help viewers understand the steps involved in creating redox reactions and calculating the associated electrode potentials.

The speaker explains why copper does not react with most acids but does react with nitric acid, using the given electrochemical series data. The explanation relies on comparing the electrode potentials of the copper half-cell and the acid in question. The speaker provides a step-by-step breakdown of how to construct the explanation and write the balanced chemical equation for the reaction between copper and nitric acid. The focus is on understanding the redox processes involved and applying the NOPR rule to predict the outcome of reactions between different species.