Chapter 4_The Mole and Chemical Reactions_released

Chapter 4: The Mole and Chemical Reactions

Chapter Objectives

• 4.1 Convert between moles, grams, and atoms of elements.

• 4.2 Convert between moles, grams, compounds, and atoms in compounds.

• 4.3 Calculate the percent by mass of atoms in compounds.

• 4.4 Write a balanced chemical equation from a given chemical reaction.

• 4.5 Use the mole concept for stoichiometric calculations in balanced chemical equations.

• 4.6 Use the mole concept to calculate percent yields.

• 4.7 Identify oxidizing and reducing agents in redox reactions given oxidation states.

• 4.8 Classify reactions as redox or nonredox, and as decomposition, combination, single replacement, double replacement, or combustion.

Section 4.1: Avogadro’s Number and The Mole

Definitions

• Mole: Amount of substance containing as many elementary entities (atoms, molecules) as there are atoms in 12 g of carbon-12. One mole corresponds to 6.022 x 10²³ particles (Avogadro’s number).

• Avogadro’s Number: 6.022 × 10²³ particles per mole, the number of atoms or molecules in one mole of a substance.

Molar Mass

• Molar mass is defined as the mass of one mole of a substance in grams, equivalent to its atomic or molecular weight expressed in atomic mass units (amu).

• Example for Carbon: 1 mol = 12.01 g.

Section 4.2: The Mole and Chemical Formulas

Calculating Molecular Weight

• Molecular weight (MW) is calculated by summing the atomic weights of all atoms in a chemical formula.

• Example for Carbon Dioxide (CO₂): MW = (1 x 12.0 amu) + (2 x 16.0 amu) = 44.0 amu.

Examples and Practice

• Example calculations using dimensional analysis based on molar mass and counting molecules in various scenarios (e.g., sulfur, carbon dioxide).

Section 4.3: Chemical Equations

Law of Conservation of Matter

• Atoms cannot be created or destroyed in a reaction; they merely rearrange.

• Example: 2H₂ + O₂ → 2H₂O.

Writing Chemical Equations

• A balanced equation has the same number of atoms of each element on both sides.

• Coefficients are adjusted to achieve balance. Symbols indicate the state of matter: (g) gas, (l) liquid, (s) solid, (aq) aqueous.

Section 4.4: Stoichiometry and Percent Yield

Stoichiometry

• The study of quantitative relationships in chemical reactions.

• Example: In the combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O), 1 mole of methane reacts with 2 moles of oxygen.

Percent Yield

• Theoretical yield: The maximum amount of product that can be predicted based on the balanced equation.

• Actual yield: The measured amount of product obtained from an experiment.

• Percent yield formula: (actual yield / theoretical yield) × 100%.

Section 4.5: Reaction Types

Types of Chemical Reactions

• Redox Reactions: Involves oxidation (loss of electrons) and reduction (gain of electrons).

• Decomposition: A single substance breaks down into two or more products (e.g., A → B + C).

• Combination: Two or more substances form a single product (e.g., A + B → C).

• Single Replacement: One element replaces another in a compound (A + BX → B + AX).

• Double Replacement: Two compounds exchange components (AX + BY → BX + AY).

• Combustion: Hydrocarbons react with oxygen to produce carbon dioxide and water (CxHy + z O₂ → x CO₂ + y/2 H₂O).

Section 4.6: Identifying Redox Reactions

Oxidizing and Reducing Agents

• Oxidizing Agent: Accepts electrons, gets reduced.

• Reducing Agent: Donates electrons, gets oxidized.

• Oxidation Numbers: Indicate the degree of oxidation of an atom in a compound.

Examples of Redox Processes

• Typical examples and classification of specific reactions as either redox or nonredox based on oxidation state changes.