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10: The Mole

Last updated

Mar 21, 2025
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- 9.24: Sigma and Pi Bonds
- 10.1: Avogadro's Number

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The mole is the unit of measurement in the International System of Units (SI) for amount of substance. It is defined as the amount of a chemical substance that contains as many elementary entities (e.g., atoms, molecules, ions, electrons, or photons). This number is expressed by the Avogadro constant, which has a value of $6.022140857 \times 10^{23} \;mol^{-1}$ . The mole is one of the base units of the SI, and has the unit symbol mol.

10.1: Avogadro's Number
This page explains Avogadro's number, $6.02 \times 10^{23}$ , which quantifies the number of representative particles in a mole, allowing chemists to measure atoms and molecules. It discusses the mole as the SI unit for substance amount, with representative particles being atoms for elements and molecules for compounds. Additionally, it mentions National Mole Day, celebrating this key chemistry concept.
10.2: Conversions Between Moles and Atoms
This page explains conversion methods between moles, atoms, and molecules, emphasizing the convenience of moles for simplifying calculations. It provides examples on converting carbon atoms to moles and determining hydrogen atoms in water and sulfuric acid. The importance of knowing chemical formulas for accurate calculations is highlighted, accompanied by step-by-step calculation processes. The document concludes with review questions on the discussed concepts.
10.3: Molar Mass
This page explains how to calculate the amount of a substance needed for a solution based on its molar mass, using carbon dioxide and calcium nitrate as examples. It highlights the relationship between moles and grams and underscores the importance of molar mass for accurate measurements in solution preparation. The page also includes review questions to help reinforce understanding of molar mass and its calculations.
10.4: Conversions Between Moles and Mass
This page discusses the importance of measuring product yield in chemical manufacturing, highlighting the need for accurate conversions between moles and mass. It emphasizes the link between molar mass and these conversions, providing practical examples such as calculating the mass needed for specific moles of calcium chloride. The content stresses the significance of precision in calculations and significant figures in laboratory work, concluding with a review of mole and mass conversions.
10.5: Conversions Between Mass and Number of Particles
This page outlines Avogadro's contributions to gas laws, focusing on the connections between gas volume, particle count, and mass. It details the conversion of mass to the number of particles through moles, featuring an example with 20 grams of chlorine gas ( $\ce{Cl_2}$ ). The summary highlights the calculation methods and includes review questions to reinforce understanding of mass and particle conversions.
10.6: Avogadro's Hypothesis and Molar Volume
This page discusses the importance of gauges for scuba divers to monitor gas supply for safety. It explains Avogadro's Hypothesis, which asserts that equal volumes of gas at the same temperature and pressure contain the same number of particles, regardless of mass. The effects of pressure and temperature on gas volume are highlighted, with standard temperature and pressure (STP) defined at 0°C and 1 atm, where one mole of gas occupies 22.4 liters.
10.7: Conversions Between Moles and Gas Volume
This page discusses the measurement of gas volume in chemistry, focusing on the calculation of moles for optimal reactions. It highlights the concept of molar volume at standard temperature and pressure (STP), where $1 \, \text{mol} = 22.4 \, \text{L}$ , and provides examples for converting gas volumes to moles and vice versa. The importance of maintaining STP conditions during these calculations is emphasized, along with practice questions for further comprehension.
10.8: Gas Density
This page explains carbon dioxide's sinking behavior in air due to its higher density compared to lighter gases. It defines gas density (mass per unit volume) and relates it to molar mass, providing examples of gas density calculations at standard temperature and pressure (STP). The content highlights conversions between molar mass and gas density and includes review questions to enhance comprehension.
10.8: Mole Road Map
This page explains the "mole road map" in chemistry, which aids in understanding moles for calculations related to mass, particle number, and gas volume at STP. It presents a structured method for conversions and includes an example of calculating the volume of 79.3 g of neon gas, resulting in 88.0 L. The mole road map is emphasized as a valuable tool for systematic problem-solving in chemistry.
10.10: Percent Composition
This page explains how to calculate the percent composition of compounds using mass data and chemical formulas, highlighting the mass percentage of each element. It provides examples, such as peanut butter and dichlorine heptoxide, to demonstrate the calculation method. Additionally, it discusses the conservation of stained glass and techniques for preserving historical artifacts.
10.11: Percent of Water in a Hydrate
This page explains how the presence of water molecules in hydrates affects the color of copper sulfate and cobalt (II) chloride. Hydrated forms, like cobalt (II) chloride hexahydrate, are magenta due to their water content, while anhydrous forms are blue. It also describes how to calculate the percentage of water in a hydrate, noting that cobalt (II) chloride hexahydrate has about 45.44% water.
10.12: Determining Empirical Formulas
This page explains how to determine empirical formulas in chemistry, especially for organic compounds. It defines empirical formulas as the simplest whole-number ratios of elements and distinguishes them from molecular formulas. The process involves elemental analysis, including converting percentages to grams, calculating moles, and finding whole-number ratios. An example of iron and oxygen is given, resulting in the formula $\ce{Fe_2O_3}$ .
10.13: Determining Molecular Formulas
This page discusses how to differentiate between molecules such as glucose and sucrose by analyzing their molecular weights, despite their similar empirical formulas. It explains molecular and empirical formulas with examples, outlines the steps to calculate a molecular formula from an empirical formula, and includes review questions related to the topic.

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