Atoms & Elements

Avogadro's Number and the Mole Concept

The mole is a fundamental unit in chemistry used to count particles such as atoms, ions, or molecules. Avogadro's number provides the exact number of particles in one mole.

• Avogadro's Number: particles per mole.

• Application: Used to convert between number of particles and moles.

• Formula:

• Example: How many atoms are in 2.5 moles of iron? atoms

Molar Mass Calculations

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). It is calculated using the atomic masses from the periodic table.

• Definition: The sum of the atomic masses of all atoms in a molecule or formula unit.

• Formula:

• Example: Molar mass of H2O: g/mol

Molecules, Compounds & Chemical Equations

Naming Ionic and Molecular Compounds

Chemical nomenclature is essential for identifying compounds. Ionic compounds consist of metals and nonmetals, while molecular compounds are formed from nonmetals.

• Ionic Compounds: Name the cation first, then the anion. E.g., NaCl is sodium chloride.

• Molecular Compounds: Use prefixes to indicate the number of each element. E.g., CO2 is carbon dioxide.

• Polyatomic Ions: Common ions include nitrate (), sulfate (), and ammonium ().

• Hydrates: Compounds with water molecules attached. E.g., CuSO4·5H2O is copper(II) sulfate pentahydrate.

Mass Percent Calculations

Mass percent expresses the proportion of each element in a compound by mass.

• Formula:

• Example: Mass percent of oxygen in H2O:

Empirical and Molecular Formula Determination

Empirical formulas show the simplest ratio of elements, while molecular formulas show the actual number of atoms in a molecule.

• Empirical Formula: Derived from percent composition or combustion analysis.

• Steps:

  1. Convert mass percent to grams (assume 100 g sample).

  2. Convert grams to moles.

  3. Divide by smallest number of moles to get ratio.

  4. Multiply to get whole numbers if necessary.

• Molecular Formula: where

• Example: If empirical formula is CH2 and molar mass is 28 g/mol, molecular formula is C2H4.

Chemical Reactions & Chemical Quantities

Balancing Chemical Equations

Balanced equations ensure the conservation of mass and atoms. Each side of the equation must have the same number of each type of atom.

• Steps:

  1. Write the unbalanced equation.

  2. Balance atoms one at a time, starting with the most complex molecule.

  3. Adjust coefficients as needed.

  4. Check that all atoms are balanced.

• Example:

Stoichiometry and Chemical Quantities

Stoichiometry uses balanced equations to relate quantities of reactants and products.

• Stoichiometric Relationships: Use mole ratios from the balanced equation.

• Formula:

• Example: In , 4 moles of H2 produce 4 moles of H2O.

Using a Periodic Table

A "clean" periodic table will be provided during the exam. Students should be familiar with using it to find atomic masses, element symbols, and other relevant information.

• Atomic Mass: Used for molar mass calculations.

• Element Symbols: Used in writing formulas and equations.

Summary Table: Key Concepts for Exam 2

Additional info: This guide covers the main topics listed for Exam 2, including representative examples and formulas. Students should review all material from Chapters 2–4, including practice problems and concepts not explicitly listed here.