Chemical Quantities and the Composition of Matter

Conversion Factors and Moles

Understanding the relationships between mass, moles, and number of particles is fundamental in chemistry. Conversion factors allow chemists to switch between these quantities using molar mass and Avogadro's number.

• Mole (mol): The SI unit for amount of substance. One mole contains Avogadro's number of entities (atoms, molecules, ions, etc.).

• Avogadro's Number: particles per mole.

• Molar Mass: The mass (in grams) of one mole of a substance, numerically equal to the substance's relative molecular or atomic mass.

Key Conversion Factors:

• Grams ↔ Moles: Use molar mass as the conversion factor.

• Moles ↔ Particles: Use Avogadro's number as the conversion factor.

Example: To convert 12.016 g of H2O to moles:

• Molar mass of H2O = 18.016 g/mol

Example: To find the number of molecules in 2.00 mol H2O:

Empirical and Molecular Formulas

Empirical and molecular formulas represent the composition of compounds. The empirical formula shows the simplest whole-number ratio of atoms, while the molecular formula gives the actual number of atoms in a molecule.

• Empirical Formula: The simplest integer ratio of elements in a compound.

• Molecular Formula: The actual number of each type of atom in a molecule; it is a whole-number multiple of the empirical formula.

Steps to Determine Empirical Formula:

1. Convert the mass of each element to moles using their molar masses.

2. Write a pseudoformula using the mole values as subscripts.

3. Divide all subscripts by the smallest number of moles to get the simplest ratio.

4. If any subscript is not a whole number, multiply all subscripts by the smallest factor to obtain whole numbers.

Example: Given 30.97 g P and 69.03 g O:

• Moles of P:

• Moles of O:

• Pseudoformula: P1.00O4.31

• Divide by smallest: P1O4.31

• If necessary, multiply to get whole numbers (e.g., if 4.31 is close to 4.33, multiply by 3 to get P3O13).

Determining Molecular Formula:

1. Find the empirical formula mass (EFM).

2. Divide the compound's molar mass by the EFM to get a whole number (n).

3. Molecular formula = (Empirical formula) × n

Example: If empirical formula is CH2O (EFM = 30.026 g/mol) and molar mass is 180.156 g/mol:

• Molecular formula: C6H12O6

Solutions and Concentration

Definitions and Types of Solutions

A solution is a homogeneous mixture of two or more substances. The substance being dissolved is the solute, and the substance doing the dissolving is the solvent. When water is the solvent, the solution is called an aqueous solution.

• Solute: The substance present in a lesser amount; it is dissolved in the solvent.

• Solvent: The substance present in a greater amount; it dissolves the solute.

• Aqueous Solution: A solution in which water is the solvent.

Concentration of Solutions

The concentration of a solution describes the amount of solute dissolved in a given quantity of solvent or solution. Two common terms are:

• Dilute Solution: Contains a small amount of solute relative to solvent.

• Concentrated Solution: Contains a large amount of solute relative to solvent.

Molarity (M)

Molarity is the most common unit of concentration in chemistry, defined as the number of moles of solute per liter of solution.

• Formula:

• Molarity is expressed in units of mol/L (M).

Example: To prepare a 0.50 M NaCl solution, dissolve 0.50 mol NaCl in enough water to make 1.00 L of solution.

Sample Table: Types of Solution Concentrations

Additional info: Molality and percent by mass are other common ways to express concentration, though not directly mentioned in the notes.

Summary

• Use conversion factors to relate mass, moles, and number of particles.

• Empirical formulas give the simplest ratio of elements; molecular formulas give the actual number of atoms.

• Solutions are homogeneous mixtures; concentration is most commonly measured in molarity (mol/L).