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16: Solutions

Last updated

Mar 21, 2025
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- 15.11: Colloids
- 16.1: Solute-Solvent Combinations

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Topic hierarchy

16.1: Solute-Solvent Combinations
This page discusses Chapter 15, which highlights water's role in aqueous solutions and differentiates between solutions, suspensions, and colloids. It explores various solute-solvent combinations, including gas-gas, solid-solid, and liquid-liquid solutions, emphasizing that while water is a key solvent, other combinations are also present. The chapter raises concerns about mercury's toxicity in dental amalgam despite its effectiveness in binding metals.
16.2: Rate of Dissolution
This page explains how sugar dissolves in iced tea, highlighting factors such as sugar amount, temperature, surface area, and agitation. Stirring enhances dissolution by increasing interactions between solvent and solute, while smaller sugar particles dissolve more quickly due to a larger surface area. Higher temperatures improve dissolution rates by increasing the energy and collision frequency of solvent molecules.
16.3: Saturated and Unsaturated Solutions
This page explains recrystallization as a method for purifying compounds by dissolving them in hot solvent and allowing them to precipitate when cooled. It distinguishes between saturated (maximum solute dissolved) and unsaturated (less than maximum) solutions, noting that saturation occurs when the rates of dissolution and recrystallization balance. Undissolved solute indicates a saturated solution, while dissolved excess indicates an unsaturated one.
16.4: How Temperature Influences Solubility
This page discusses the environmental impact of nuclear power plants on aquatic ecosystems due to water usage for cooling and steam generation, which leads to temperature increases and lower oxygen levels. It explains solubility, noting that while solids usually dissolve more in warmer conditions, gases do not. The page also features solubility curves that help determine whether solutions, such as those containing potassium nitrate, are saturated or unsaturated.
16.5: Supersaturated Solutions
This page describes a thermal pack containing a supersaturated sodium acetate solution that can heat or cool. Activating it via a metal trigger causes crystallization, releasing heat. Supersaturation involves solute exceeding usual capacity at a specific temperature, with crystallization initiated by a seed crystal, enabling quick solidification of the excess solute.
16.6: Henry's Law
This page discusses the challenges of maintaining carbonation in soft drinks in outer space due to microgravity. It explains that carbonation dissipates without pressure and highlights the use of a pressurized container to keep gases dissolved. Henry's Law is referenced to show that higher gas pressure increases solubility. Opening a bottle reduces pressure, leading to gas escape and flat beverages over time.
16.7: Percent Solutions
This page discusses the variation in numerical recognition across cultures, noting that some do not count beyond three. It highlights American cultural practices in expressing solution concentrations, defining concentrated and dilute solutions, and explaining how concentrations can be represented in mass percent and volume percent. The page includes examples to illustrate methods for calculating these percentages.
16.8: Molarity
This page explains molarity as a concentration measure in solutions, defined as moles of solute per liter of solution. It contrasts molarity with percent solutions, which measure mass instead of molecular quantity. The text includes examples for calculating molarity and finding the mass of solute required for specific concentrations, underscoring the importance of accurate measurements in laboratory settings.
16.9: Preparing Solutions
This page discusses the shift from intuitive cooking to precise scientific preparation in cooking, highlighting the example of making a 1.00 L solution of 1.00 M sodium chloride. It emphasizes the importance of using a volumetric flask for accurate measurements, detailing the steps of weighing the solute, dissolving it, and adding water to reach the desired volume before mixing thoroughly.
16.10: Dilution
This page explains the safe use of muriatic acid (HCl) for cleaning concrete, emphasizing the need for dilution from concentrations of around 18%. It details dilution concepts, including the relationship between initial and final concentrations and volumes. An illustrative example demonstrates how to dilute concentrated nitric acid to achieve a specific molarity, highlighting the importance of precise measurements with pipettes and micropipettes in laboratory settings.
16.11: Molality
This page discusses the differences between molarity and molality as measures of solution concentration. Molarity is volume-based and influenced by temperature, while molality is mass-based and unaffected by temperature. An example of a one-molal sodium chloride solution is provided. Although both measures are similar in dilute solutions, they differ significantly in concentrated solutions, making molality the preferred choice in temperature-dependent scenarios.
16.12: The Lowering of Vapor Pressure
This page explains colligative properties, which depend on the quantity of solute particles in a solution rather than their type. Nonvolatile solutes decrease vapor pressure by occupying surface space and limiting solvent evaporation. Electrolytes, like sodium chloride, create more dissolved particles by dissociating into ions, enhancing the vapor pressure reduction compared to nonelectrolytes like glucose.
16.13: Freezing Point Depression
This page discusses colligative properties, particularly freezing point depression, and their practical applications like using salts to improve road safety by lowering ice melting points. Common salts include sodium chloride, calcium chloride, and magnesium chloride. The freezing point depression depends on solute concentration and is quantified by the molal freezing-point depression constant, Kf, unique to each solvent.
16.14: Boiling Point Elevation
This page explains that salt enhances the flavor of boiling water while only slightly affecting the boiling point. A significant rise in boiling point requires more than 100 grams of salt. Boiling point elevation is linked to lower vapor pressure in solutions and depends on molality, following the equation where the molal boiling-point elevation constant ($K_b$) for water is $0.512^\text{o} \text{C}/\textit{m}$.
16.15: Electrolytes and Colligative Properties
This page discusses how ionic compounds, as electrolytes, dissociate into ions in solution, affecting colligative properties like freezing and boiling points. Using calcium chloride as an example, it demonstrates this dissociation leads to significant alterations in freezing point (calculated at -10.3°C) and boiling point (calculated at 102.84°C), illustrating the substantial impact of ionization.
16.16: Calculating Molar Mass
This page explains how antifreeze in a radiator prevents engine freezing and discusses determining the molar mass of an unknown solute via freezing point depression. By measuring the freezing point change after dissolving a known mass of solute, one can calculate molality and moles of solute, leading to a calculated molar mass of 59.7 g/mol in an example.
16.17: Molecular and Ionic Equations
This page explains how acid rain forms from industrial activities, damaging limestone via reactions with sulfur dioxide and nitrogen oxides. It distinguishes between molecular equations, which show compounds as molecules, and ionic equations, which show them as dissociated ions. Examples of double-replacement reactions, including precipitate formation, and single-replacement reactions, where one element replaces another in a compound, are provided.
16.18: Net Ionic Equations
This page explains spectator ions in chemical reactions and their significance in net ionic equations. It illustrates their role with a silver chloride precipitation example, showing how to create a net ionic equation by omitting spectator ions. The text also emphasizes the importance of balancing equations for mass and charge, concluding that net ionic equations focus solely on the active participants in a reaction.
16.19: Predicting Precipitates Using Solubility Rules
This page discusses two main topics: weather prediction, which relies on analyzing data such as wind patterns and barometric pressure, and the chemistry of precipitates, highlighting solubility rules and examples of chemical reactions that form precipitates, such as cesium bromide and lead (II) nitrate, which produces lead (II) bromide.

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