Chapter 3: Chemical Reactions and Reaction Stoichiometry

Balancing Chemical Equations

Balancing chemical equations is essential to ensure the conservation of mass and atoms in a chemical reaction. Each side of the equation must have the same number of atoms for each element.

• Key Point: Adjust coefficients (not subscripts) to balance atoms.

• Key Point: Start with the most complex molecule and balance elements one at a time.

• Example:

Calculating Moles from Mass

To convert grams of an element or compound to moles, use the molar mass as a conversion factor.

• Key Point: Molar mass is the mass (in grams) of one mole of a substance.

• Formula:

• Example: 18 g of water () is mole.

Calculating Atoms or Molecules from Moles

Avogadro’s number () is used to convert between moles and the number of particles (atoms or molecules).

• Key Point: 1 mole contains particles.

• Formula:

• Example: 2 moles of atoms contain atoms.

Empirical Formula Determination

The empirical formula represents the simplest whole-number ratio of elements in a compound. It is determined from the mass composition of each element.

• Key Point: Convert mass to moles for each element, then divide by the smallest number of moles.

• Example: If a compound contains 40 g C and 6.7 g H:

  • C: mol

  • H: mol

  • Ratio: C:H = 1:2 → Empirical formula

Molecular Formula Determination

The molecular formula shows the actual number of atoms in a molecule. It is found by comparing the empirical formula mass to the molar mass.

• Key Point:

• Formula: Multiply the empirical formula by n to get the molecular formula.

• Example: Empirical formula , molar mass 28 g/mol. Empirical mass = 14 g/mol, so n = 2. Molecular formula = .

Stoichiometric Calculations

Stoichiometry involves quantitative relationships between reactants and products in a chemical reaction.

• Key Point: Use balanced equations to relate moles of substances.

• Steps:

  1. Convert mass to moles.

  2. Use mole ratios from the balanced equation.

  3. Convert moles back to mass if needed.

• Example: : 4 mol produces 4 mol .

Chapter 4: Reactions in Aqueous Solution

Electrolytes: Strong, Weak, and Non-Electrolytes

Electrolytes are substances that produce ions in solution and conduct electricity. They are classified based on their ionization in water.

• Strong Electrolytes: Completely dissociate into ions (e.g., NaCl, HCl).

• Weak Electrolytes: Partially dissociate (e.g., acetic acid, ammonia).

• Non-Electrolytes: Do not produce ions (e.g., sugar, ethanol).

• Example: NaCl in water is a strong electrolyte; CH3COOH is a weak electrolyte.

Precipitation Reactions and Ionic Equations

Precipitation reactions occur when two solutions mix and form an insoluble product (precipitate). Equations can be written in three forms.

• Molecular Equation: Shows compounds as intact molecules.

• Complete Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the ions and molecules directly involved in the reaction.

• Example:

  • Molecular:

  • Complete Ionic:

  • Net Ionic:

• Solubility Rules: Used to predict if a precipitate will form. (Will be provided during the exam.)

Acid-Base Neutralization Reactions

Strong acid and strong base neutralization reactions produce water and a salt. Equations can be written in molecular, complete ionic, and net ionic forms.

• Example:

  • Molecular:

  • Complete Ionic:

  • Net Ionic:

Oxidation-Reduction (Redox) Reactions and Oxidation Numbers

Redox reactions involve the transfer of electrons. Oxidation numbers help track electron movement.

• Key Point: Oxidation is loss of electrons; reduction is gain of electrons.

• Rules for Assigning Oxidation Numbers:

  1. Elemental form: 0

  2. Monatomic ion: charge of ion

  3. Oxygen: usually -2

  4. Hydrogen: +1 (except in metal hydrides: -1)

  5. Fluorine: always -1

  6. Sum of oxidation numbers equals charge of molecule/ion

• Example: In , H is +1, O is -2.

Calculating Molarity of Solutions

Molarity (M) is a measure of concentration, defined as moles of solute per liter of solution.

• Formula:

• Example: 0.5 moles NaCl in 1.0 L solution:

Exam Format and Allowed Materials

• Short Answer and Calculation Problems: Show all work for calculations.

• Closed Book, No Notes: Only a Periodic Table will be provided.

• Calculator: Stand-alone scientific calculator allowed; no phones or other electronics.