Copy of Unit 8 Review Guide 2025

The Mole and Avogadro's Number

• Avogadro's Number: There are 6.022 x 10^23 particles in a mole of any substance.

Unit 8 Exam Review Guide

• Exam content focuses on the following learning objectives:

  • Count small particles by mass.

  • Convert properly between mass, moles & particles.

  • Differentiate between atoms, molecules & formula units.

  • Use the triple threat for every step in quantitative calculations.

  • Calculate a compound's molar mass.

  • Calculate a compound's percent composition.

  • Determine a compound's empirical formula when given percent composition.

  • Determine a compound's molecular formula when given its empirical formula and molar mass.

  • Determine the molarity of a solution.

Exam Format

• 20 Multiple Choice Questions

• 4 Free Response Questions:

  1. Molar mass

  2. Percent composition

  3. Mole conversion problem + Molarity

  4. Empirical/Molecular formula problem

Study Resources

• Review the following resources:

  • Unit 8 - The Mole Slides

  • Mole Crossword Puzzle (answer key provided)

  • Murder on Mole Island (answer key provided)

  • Molevelopes (answer key provided)

  • Molecular Formulas and Molarity Practice Set

Types of Particles

1. Atoms: Elements found on the periodic table.

   • Example: Iron (Fe)

2. Molecules: Chemical formulas for covalent compounds.

   • Example: Water (H2O)

3. Formula Units: Chemical formulas for ionic compounds.

   • Example: Sodium Chloride (NaCl)

Calculating Molar Mass

• To calculate molar mass of a compound:

  • Add up the molar masses of each individual atom in the compound.

Example Calculation: Nitric Acid (HNO3)

Percent Composition

• Percent Element = (Mass of element in compound) / (Total mass of compound)

• A pure compound always consists of the same elements combined in the same proportions by weight.

Types of Chemical Formulas

1. Molecular Formula: Shows the actual number of atoms in a molecule or formula unit.

2. Empirical Formula: Shows the smallest whole number mole ratio.

3. Structural Formula: Includes molecular formula information plus bonding arrangements (not on the exam).

Differences between Molecular and Empirical Formulas

• Molecular and empirical formulas represent the same proportions by weight of all their respective elements. However, empirical formulas may not reflect the actual composition of a compound, although they can be identical (e.g., CO2).

Determining Formulas and Structures

• Understanding how to derive empirical and molecular formulas based on percent compositions and given data is crucial.

Practice Problems

Concept Questions:

1. Can empirical and molecular formulas be the same? Give examples (e.g. Methane: CH4).

2. What is Avogadro's number and what does it represent? (6.022 x 10^23 particles).

3. Distinguish between smallest particles of different compounds:

   • Pure elements: atoms

   • Covalent compounds: molecules

   • Ionic compounds: formula units

Molar Mass Calculations:

1. Molar mass of boron: 10.811 g/mol

2. Molar mass of sodium chloride: 58.443 g/mol

3. Molar mass of lithium phosphate: 115.79 g/mol

Example Mole Conversions:

1. Grams in 6.95 mol of CO2: Calculate using molar mass.

2. Particles in 53.65 mol of fluorine: Use Avogadro's number.

3. Moles in 64.43 g of KCl: Calculate using molar mass.

Percent Composition Calculations:

1. Percent by mass of carbon in CCl4.

2. Percent by mass of oxygen in Al2O3.

3. Percent by mass of potassium in potassium nitrate.

Empirical and Molecular Formula Calculations:

1. Calculate empirical formulas (samples provided).

2. Recognize the method: Percent → Mass → Mass → Moles → Divide by Smallest → Return to Whole.

3. Determining molecular formula from empirical formula and molar mass is key (samples provided).