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19: Equilibrium

Last updated

Mar 21, 2025
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- 18.15: Mechanisms and Potential Energy Diagrams
- 19.1: Reversible Reaction

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Topic hierarchy

19.1: Reversible Reaction
This page discusses the color change of cobalt chloride solutions due to solvated $\ce{Co^{2+}}$ ions, which turn pink, and shift to blue upon the addition of $\ce{HCl}$ forming $\ce{CoCl_4^{2-}}$ ; this change can be reversed by adding water. It further explains reversible chemical reactions, exemplified by the formation and decomposition of hydrogen iodide from hydrogen and iodine, emphasizing that this process continues until a balance is reached between the forward and reverse reactions.
19.2: Chemical Equilibrium
This page explains chemical equilibrium using the metaphor of a tug of war, where forces are equal but no movement occurs. It highlights the dynamic equilibrium in reactions like hydrogen iodide formation, where forward and reverse reactions occur at equal rates, maintaining constant concentrations. Key conditions for equilibrium include a closed system and the presence of both reactants and products. It also mentions other forms of equilibrium, such as phase and solution equilibrium.
19.3: Equilibrium Constant
This page discusses the role of hemoglobin in transporting oxygen in red blood cells and the effects of carbon monoxide poisoning, which impairs oxygen transport by binding to hemoglobin. Treatment involves administering pure oxygen to displace CO. Additionally, it explains the concept of the equilibrium constant ( $K_\text{eq}$ ), which quantifies the ratio of product to reactant concentrations in chemical reactions at equilibrium, a value determined experimentally and dependent on temperature.
19.4: Calculations with Equilibrium Constants
This page discusses the importance of iron in red blood cells and its role in preventing anemia. It explains the use of Ferrozine for assessing serum iron, highlighting the significance of the equilibrium constant $K_\text{eq}$ in indicating product favorability.
19.5: Le Châtelier's Principle
This page explores two themes: the thrill of skydiving for stress relief through adrenaline, and Le Chatelier's Principle in chemistry, which explains how a chemical equilibrium reacts to external stressors by shifting towards either the forward or reverse reaction, impacting the concentrations of reactants and products.
19.6: Effect of Concentration
This page discusses phenolphthalein's color change in relation to hydrogen ion concentration and the principles of equilibrium in chemical reactions. It highlights how altering reactant or product concentrations impacts equilibrium, exemplified by the Haber-Bosch process for ammonia production. Adding reactants shifts equilibrium towards product formation, while adding products reverses this shift.
19.7: Effect of Temperature
This page discusses the dual nature of carbon monoxide (CO) as both a hazardous gas and a key component in acetic acid production. It highlights the effects of temperature on chemical equilibria, emphasizing Le Chatelier's principle, which explains how temperature changes influence reactions such as the Haber-Bosch process and the equilibrium between dinitrogen tetroxide and nitrogen dioxide.
19.8: Effect of Pressure
This page discusses the dual purpose of ammonia storage tanks for high-pressure storage to prevent reverse reactions and deter theft in illicit drug production. It highlights the effects of pressure changes on chemical equilibrium, noting that increased pressure favors the formation of fewer gas molecules, while decreased pressure benefits reactions producing more gas. The content also states that solids and liquids are not affected by pressure changes regarding equilibrium shifts.
19.9: Nonreversible Reactions
This page discusses the dual impact of fires on environments, highlighting their role in rejuvenation and destruction through greenhouse gas emissions. It also covers chemical reactions that proceed to completion, which do not reverse and often create products such as insoluble precipitates or gases. Examples provided include reactions involving silver chloride, hydrogen gas generation, and neutralization resulting in water and sodium chloride.
19.10: Le Châtelier's Principle and the Equilibrium Constant
This page discusses two distinct topics: online banking and Le Chatelier's Principle in chemistry. Online banking enhances personal finance management through features like automatic deposits and payments, ensuring a stable savings-to-checking ratio.
19.11: Solubility Product Constant $\left( K_\text{sp} \right)$
This page discusses gravimetric analysis, a traditional method for determining ion quantities through precipitation and weighing, noting its accuracy but slowness. It contrasts this with modern, faster ion measurement techniques. The solubility of ionic compounds varies, with some being highly soluble and others, like zinc hydroxide, being mostly insoluble.
19.12: Conversion of Solubility to $K_\text{sp}$
This page details the preparation of sodium bicarbonate from ammonia and sodium chloride via a reaction with carbon dioxide, leading to its precipitation. It covers solubility measurements in grams per liter and their conversion to molar solubility. Using lead (II) fluoride as a case study, the page illustrates how to calculate the solubility product constant (K_sp) from molar solubility by determining ion concentrations and applying the K_sp equation.
19.13: Conversion of $K_\text{sp}$ to Solubility
This page outlines the purification of drinking water, emphasizing the removal of heavy metals through the formation of insoluble compounds using carbonates and sulfates. It details the calculation of molar solubility via solubility product constants ( $K_\text{sp}$ ), including the creation of an ICE table to find ion concentrations and deriving molar solubility. The final conversion from molar solubility to solubility is also addressed.
19.14: Predicting Precipitates
This page discusses the impact of the x-ray machine on medical diagnosis, emphasizing barium sulfate's role in x-ray imaging. It also explains predicting precipitate formation in chemistry using the ion product and solubility product constant (K_sp). By calculating ion concentrations, one can determine precipitate formation, illustrated with a practical example of barium sulfate precipitate formation based on the ion product calculation.
19.15: Common Ion Effect
This page discusses lithium carbonate's role in lithium batteries, produced through a reaction with carbon dioxide. It explains the common ion effect, where adding a common ion decreases a compound's solubility, illustrated by adding calcium nitrate to calcium sulfate. This shifts equilibrium, causing precipitation. Additionally, calculations show how common ions, like hydroxide with zinc hydroxide, affect ion concentrations in solutions.

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