Topic: Cell potential (NON-standard conditions) / Nernst Equation

This episode covers Le Chatelier's Principle, explaining how a chemical system at equilibrium will always shift its position to counteract and undo an external "stress." It explores three primary stressors: concentration changes, where the system shifts away from what is added and toward what is removed; pressure or volume changes, which force a shift toward the side with fewer gas moles when pressure increases (and vice versa); and temperature changes, which treat heat as a reactant or product depending on whether the reaction is endothermic or exothermic. Crucially, the episode highlight that while concentration and pressure shifts occur to keep the equilibrium constant (K) the same, temperature is the only stressor that actually changes the value of K. Finally, it points out that adding a catalyst or an inert gas has no effect on the equilibrium position, concluding that when analyzing these systems, the chemical balance always votes to restore order.

This episode breaks down the fundamentals of the pH scale, explaining its concept, mathematics, and real-world applications. It defines pH as the "power of hydrogen," a logarithmic measurement of hydrogen ion concentration ([H+]) where each unit represents a tenfold change in acidity or basicity. Through real-world examples ranging from battery acid to bleach, the episode demonstrates how the inverse scale functions at 25°C, where a pH below 7 is acidic, 7 is neutral, and above 7 is basic or alkaline. The discussion covers essential calculations, illustrating how to toggle between pH and hydrogen ion concentration using the formulas pH = -log[H+] and [H+] = 10^-pH. Additionally, it introduces the complementary concept of pOH (pH + pOH = 14) and conclude with practical study tips to help students intuitively master logarithmic shifts and calculator functions.

This episode introduce the Reaction Quotient (Q) and explain how it serves as a "GPS" to predict whether a chemical reaction will shift toward products or reactants to reach its destination, the Equilibrium Constant (K). While both share the identical mathematical expression of products over reactants raised to their coefficients, the episode highlights that K is a fixed constant at a specific temperature calculated only at equilibrium, whereas Q is a variable calculated using non-equilibrium concentrations at any random point in time. There are three potential scenarios when comparing the two: if Q = K, the system is already at equilibrium; if K > Q, the reaction shifts to the right (forward) to create more products; and if K < Q, it shifts to the left (reverse) to form more reactants. Concluding with helpful memory tricks, a practical sample problem, and a warning that Q never alters the fixed value of K, this episode provides listeners with a foundational framework for mastering chemical shift predictions.

This episode introduce the ICE table (Initial, Change, Equilibrium) as a foundational tool for tracking concentration changes and calculating the equilibrium constant (K) in chemical reactions. It explain how to construct the table by using initial concentrations, applying stoichiometry and a variable (x) to determine the change, and combining them to find the equilibrium values. The episode walks through an example problem together. The episode does not go into more complex problems that might require the quadratic formula or the 5% rule.

In this episode, we explore the fundamentals of differential rate laws — the mathematical expressions that describe how the speed of a chemical reaction depends on the concentration of its reactants. We break down the general rate law equation, Rate = k[A]^m[B]^n, explaining the role of the rate constant k, reaction orders, and how to calculate the overall reaction order. Through two fully worked examples, you will learn how to use the method of initial rates to experimentally determine reaction orders and solve for the rate constant, including how to derive its units through dimensional analysis.

This episode explores molecular polarity, explaining how electronegativity differences create polar bonds and why molecular geometry determines overall polarity. The hosts distinguish between bond polarity and molecular polarity, showing how symmetrical molecules like CO₂ and CH₄ are nonpolar despite polar bonds, while asymmetrical molecules like H₂O and NH₃ are polar. They conclude with a step-by-step approach for determining molecular polarity using Lewis structures and VSEPR theory.

Faraday's Law relates the amount of chemical change in an electrolytic cell to the total electric charge passed through it. The amount of product is directly proportional to the charge. Using Faraday's constant (96,485 C/mol/e-), the charge converts to moles of electrons, which then relates to moles of product via the half-reaction stoichiometry. This allows calculation of the mass of substance produced.

Electrolysis uses electrical energy to drive a nonspontaneous redox reaction, the opposite of a voltaic cell. An electrolytic cell requires a power source, where the anode is positive and the cathode is negative (reversed polarity from voltaic). The process is used to split molten salts or aqueous solutions. For aqueous solutions, water's reduction/oxidation potential must be considered, as it often reacts instead of the solute ions.

This episode explains how to calculate non-standard cell potentials (voltage) using the Nernst equation. Real-life batteries operate under non-standard conditions where concentrations change, causing voltage to drop. The equation shows that as the reaction quotient Q increases (more product), Ecell decreases, and vice-versa, connecting electrochemistry to chemical equilibrium.

The standard cell potential (E°cell) for a voltaic cell is calculated using E°cell = E°cathode − E°anode, where E° values come from standard reduction potential tables. A positive E°cell indicates a spontaneous reaction capable of producing electricity. Reduction occurs at the cathode (higher potential), and oxidation occurs at the anode (lower potential). Cell notation is written as: Anode|Anode Ion||Cathode Ion|Cathode.