1.1.1 Why the Mole? From Atoms to Grams

• Macroscopic vs. microscopic quantities:
• A single sodium (Na) atom has a mass of approximately 3.82 × 10⁻²³ g. Measuring individual atoms directly in the laboratory is impossible.
• Chemists define the mole to bridge these scales: one mole of any substance contains a fixed, very large number of "entities" (atoms, molecules, ions, electrons, etc.), allowing us to weigh out amounts in grams.
• Definition of the mole:
• By definition, 1 mol of a substance is the amount of that substance that contains exactly 6.022 140 76 × 10²³ elementary entities. This number is called Avogadro’s constant (NA).
  • Avogadro’s constant, NA = 6.022 140 76 × 10²³ entities per mole.
• Historical basis:
• The mole was originally defined so that exactly 12 g of pure carbon-12 (¹²C) contains 6.022 140 76 × 10²³ carbon atoms. This definition ties the atomic mass unit (1 u ≈ 1.660 539 × 10⁻²⁷ kg) to a macroscopic mass.
• Detailed Explanation: This chunk introduces the mole concept, which is essential in chemistry for counting and measuring atoms and molecules. It explains the mass of a single sodium atom and why measuring single atoms directly is impractical. Instead, chemists use the mole as a unit that corresponds to a large number of particles, enabling them to convert between mass and number of particles.
  Avogadro's constant, which quantifies how many particles are in a mole, is central to this concept, bridging the gap between the microscopic scale (individual atoms) and macroscopic quantities (grams).
• Real-Life Example or Analogy: Think of a dozen eggs. Just like a dozen refers to 12 eggs, a mole refers to 6.022 x 10²³ particles. If you want to buy eggs, you won’t count them one by one; you’ll ask for a dozen. In chemistry, when working with atoms or molecules, we use moles in the same way to handle very large numbers.

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• Chunk Title: Molar Mass and Mass ⇄ Mole Conversions
• Chunk Text: #### 1.1.2 Molar Mass and Mass ⇄ Mole Conversions
• Atomic mass and molar mass:
• The number under each element symbol on the periodic table (relative atomic mass) tells us how many atomic mass units a single atom has.
• By convention, a relative atomic mass in atomic mass units (u) is numerically identical to the molar mass in grams per mole (g/mol).
  • Example: Carbon (C) has a relative atomic mass of 12.011 u; thus, its molar mass is 12.011 g/mol.
  • Chlorine (Cl) appears as 35.45 u because it is a mixture of ³⁵Cl and ³⁷Cl isotopes. Therefore, 1 mol of natural chlorine atoms has a mass of 35.45 g.
• Molecular and formula mass:
• For molecules, add the atomic masses of each constituent atom to find the molecular mass (in u). For ionic compounds or empirical formulas, the sum is called the formula mass (in u).
  • Example: Water, H₂O.
  • Mass of 2 H = 2 × 1.008 u = 2.016 u
  • Mass of 1 O = 16.00 u
  • Molecular mass of H₂O = 2.016 u + 16.00 u = 18.016 u
  • Therefore, the molar mass of H₂O is 18.016 g/mol.
• Detailed Explanation: This chunk explains how to convert between mass and moles using molar mass. Molar mass is the mass in grams of one mole of a substance and is equal to the atomic mass in atomic mass units. It also covers how to calculate the molecular mass of a compound by summing the atomic masses of its constituent atoms. Understanding these concepts is critical for stoichiometric calculations in chemical reactions.
• Real-Life Example or Analogy: Imagine making a fruit salad with different types of fruit. The molar mass is like the weight of a single serving of each type of fruit. Just as you would measure the weight of bananas, apples, and grapefruits separately and then combine them for a salad, in chemistry, you weigh out mass amounts of elements based on their respective molar masses before combining them in a reaction.

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• Chunk Title: General Mass–Mole Relationships
• Chunk Text: 1. General mass–mole relationships:
• Mass to moles:

\[ \text{number of moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}} \]
- Moles to mass:

\[ \text{mass (g)} = \text{number of moles (mol)} \times \text{molar mass (g/mol)} \]
- Moles to number of entities:

\[ \text{number of entities} = \text{number of moles (mol)} \times N_A \]
- Number of entities to moles:

\[ \text{number of moles (mol)} = \frac{\text{number of entities}}{N_A} \]
- Detailed Explanation: This chunk presents the foundational equations for converting between mass, moles, and the number of entities (atoms, molecules). These relationships are vital for performing quantitative analysis in chemistry, allowing chemists to calculate how much product will form from given amounts of reactants or how many particles are present in a specific mass.
- Real-Life Example or Analogy: Think of baking cookies. If a recipe calls for 2 cups of flour (the mass), you might check how many cookies (the entities) you can make. By knowing how many cookies are made per cup (essentially moles of cookies per cup of flour), you can easily calculate how many cookies you can make with your flour, similar to how chemists convert mass to moles to find out how many particles of substance they have.

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• Chunk Title: Practice Problem Examples
• Chunk Text: ### Practice Example 1: Converting mass of substance to number of atoms
  Problem: How many oxygen atoms are present in 5.00 g of O₂ gas?
• Compute the molar mass of O₂:
• Mass of one O atom = 16.00 u = 16.00 g/mol.
• Therefore, molar mass of O₂ = 2 × 16.00 g/mol = 32.00 g/mol.
• Convert mass of O₂ to moles:
• Number of moles of O₂ = 5.00 g ÷ 32.00 g/mol = 0.15625 mol.
• Convert moles of O₂ to number of O₂ molecules:
• Number of O₂ molecules = 0.15625 mol × 6.022 × 10²³ molecules/mol = 9.41 × 10²² molecules.
• Each O₂ molecule contains 2 oxygen atoms. Therefore:
• Number of oxygen atoms = 9.41 × 10²² molecules × 2 atoms/molecule = 1.882 × 10²³ atoms.
  Answer: 1.88 × 10²³ oxygen atoms.
• Detailed Explanation: This example walks through the calculations required to find out how many oxygen atoms are in a given mass of O₂. It demonstrates the steps involved in molar mass computation, mole conversion, and scaling up to find the number of entities (atoms). Practicing problems like this helps reinforce the mass-mole concept in a practical context.
• Real-Life Example or Analogy: If you had a bag of candies, and wanted to know how many candies you have based on the weight, you'd first need to know how much a single candy weighs (analogous to molar mass). After that, you could divide the total weight of the candies you have by the weight of one candy to determine how many candies are in the bag.

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