21. Acid-Base Equilibrium: Is MIT Water Safe to Drink?

The script begins with an introduction to MIT OpenCourseWare, encouraging donations to support free educational resources, with a specific mention of the platform's website. It then transitions into a chemistry lecture, discussing a review of Lewis acid-base definitions where a Lewis base donates a lone pair of electrons to a Lewis acid. The lecture includes an interactive element with the audience, exploring the Bronsted-Lowry definitions of acids and bases and comparing them to the Lewis definitions. The discussion aims to show the compatibility between the two sets of definitions.

This paragraph continues the chemistry lecture, focusing on the relationship between pKw, pH, and pOH, and the strengths of acids and bases. It introduces the concept of water as an amphoteric substance, capable of acting as both an acid and a base, forming H3O+ and OH- ions respectively. The lecturer explains the equilibrium constant (K) in the context of water and how it relates to the Gibbs free energy (delta G0), leading into a discussion about the spontaneity of the forward reaction and the calculation of K from delta G0.

The lecturer delves into the ionization of water, explaining that only a small fraction of water molecules are ionized, with the equilibrium constant for water (Kw) being 1.0 x 10^-14 at room temperature. The importance of Kw in acid-base problems is emphasized, as it represents the product of the hydronium and hydroxide ion concentrations. The explanation includes the exclusion of pure solvents like water from equilibrium expressions and introduces the concepts of pH and pOH, relating them to the concentrations of H3O+ and OH- ions.

This section discusses the relationship between pH, pOH, and pKw, highlighting the mathematical connection that pH + pOH = pKw = 14 at room temperature. The practical application of this relationship in problem-solving is emphasized, as is the significance of pH in determining the acidity or alkalinity of a solution. The lecture also mentions a problem set due on Friday, indicating that it will involve various acid-base problems.

The script describes a practical demonstration of pH measurement using pH paper to test various substances, including ammonia, soda, vinegar, tap water, and a prescription medicine. The audience is engaged by having volunteers measure the pH of these substances, emphasizing the importance of understanding pH in everyday life, especially in the context of medicine and its potential effects on health.

The lecture discusses the corrosive nature of certain substances, with a pH less than 3 or greater than 12.5, as defined by the EPA. It uses the example of a medicine with a pH of 2, which was found to cause sores in the lecturer's daughter's mouth, to illustrate the importance of being aware of the pH of substances we consume or administer to others. The lecturer shares a personal anecdote about finding an alternative way to administer iron supplements to her daughter using Nutella to mask the taste of the medicine, reinforcing the significance of pH in healthcare.

The script explains the concept of the strength of acids, introducing the acid ionization constant (Ka) as a measure of how much an acid ionizes in solution. It contrasts strong acids, which have a high Ka and ionize almost completely, with weak acids that have a low Ka and ionize to a lesser extent. The lecture includes an example of acetic acid, discussing its ionization in water and the resulting formation of hydronium ions and the acetate ion.

This section introduces pKa as the negative logarithm of Ka, emphasizing its importance in various scientific fields, including organic chemistry and biology. The lecture clarifies the relationship between Ka and pKa, noting that a higher pKa indicates a weaker acid, and vice versa. The importance of understanding pKa is highlighted, with a humorous anecdote about students in organic chemistry being unaware of the concept.

The script compares the strengths of different acids using their Ka and pKa values, illustrating the concept with examples from a table of acids. It explains that a strong acid will have a Ka much greater than 1 and a very low pKa, while a weak acid will have a Ka less than 1 and a higher pKa. The lecture also introduces the concept of conjugate acids and bases, noting that the strength of an acid is inversely related to the strength of its conjugate base.

The script summarizes the key concepts discussed in the lecture, including the strength of acids and bases, Ka and Kb values, and the relationship between conjugate acids and bases. It previews the topic of buffers, which will be discussed in a future class, and emphasizes the importance of understanding these concepts for future applications in various scientific fields.

This section describes the process of calculating the pH of a weak acid solution, using vitamin C as an example. The lecture explains the steps for converting mass to moles, calculating molarity, and setting up an equilibrium expression with the given Ka value. It discusses the assumption that the change in concentration (x) is small, allowing for simplification of the quadratic equation, and emphasizes the importance of checking this assumption for accuracy.

The script outlines the detailed steps for calculating the pH of a weak base solution, such as ammonia in water. It explains how to set up an equilibrium table, write the Kb expression, and calculate the change in molarity. The lecture demonstrates the use of the assumption that x is small to simplify calculations and the importance of checking this assumption. It concludes with the calculation of pOH and then pH, ensuring that the final pH value makes sense for a basic solution.