CHAPTER 5: Molecules and Compounds

Law of Constant Composition

The law of constant composition states that every sample of a given compound has the same proportion of its constituent elements by mass. This principle is fundamental to understanding chemical compounds and their formulas.

• Definition: A compound always contains the same elements in the same ratio by mass, regardless of the sample size or source.

• Example: Water (H2O) always consists of 2 hydrogen atoms and 1 oxygen atom, with a mass ratio of approximately 1:8.

Chemical Formulas

Chemical formulas represent the types and numbers of atoms in a compound. They can be empirical, molecular, or structural.

• Empirical formula: Shows the simplest whole-number ratio of elements (e.g., CH2O).

• Molecular formula: Shows the actual number of atoms in a molecule (e.g., C6H12O6).

• Structural formula: Shows the arrangement of atoms.

• Counting atoms: Use subscripts and parentheses to determine the total number of each atom. For example, in Mg(NO3)2, there are 6 oxygen atoms.

Types of Elements and Compounds

Elements and compounds can be classified based on their atomic or molecular nature and the types of bonds they form.

• Atomic elements: Exist as single atoms (e.g., He, Ne).

• Molecular elements: Exist as diatomic molecules (e.g., H2, N2, O2, F2, Cl2, Br2, I2).

• Molecular compounds: Formed from two or more nonmetals.

• Ionic compounds: Formed from a metal and a nonmetal, held together by ionic bonds (e.g., NaCl).

Writing Ionic Formulas

Ionic compounds must be electrically neutral, meaning the total positive charge equals the total negative charge.

• Balance charges: Adjust subscripts so that the sum of positive and negative charges is zero.

• Example: Aluminum oxide is Al2O3, balancing Al3+ and O2– ions.

Common vs. Systematic Names

Chemical compounds can have systematic names (following IUPAC rules) or common names (traditional).

• Systematic name: Sodium chloride

• Common name: Table salt

• Example: Water (common) vs. Dihydrogen monoxide (systematic)

Naming Ionic Compounds

Ionic compounds are named based on the type of metal involved.

• Type I: Fixed-charge metals (e.g., Na+, Ca2+): Name metal + base anion + "ide" (e.g., sodium chloride).

• Type II: Variable-charge metals (e.g., Fe2+, Fe3+): Name metal + (Roman numeral charge) + base anion + "ide" (e.g., iron(II) chloride).

• Polyatomic ions: Name the polyatomic ion unchanged (e.g., sodium nitrate).

Naming Molecular Compounds

Molecular compounds use prefixes to indicate the number of each element.

• Prefixes: mono-, di-, tri-, tetra-, penta-, etc.

• Format: Prefix + first element + prefix + second element + "ide" (e.g., dinitrogen pentoxide for N2O5).

Naming Acids

Acids are named based on their composition.

• Binary acids: hydro + base name + "ic" acid (e.g., hydrochloric acid).

• Oxyacids: If the polyatomic ion ends in "-ate", use base name + "ic" acid (e.g., sulfuric acid for H2SO4). If it ends in "-ite", use base name + "ous" acid.

Formula Mass

The formula mass is the sum of the atomic masses of all atoms in a compound.

• Formula:

• Example: CH4 has a formula mass of 16.05 amu.

CHAPTER 6: Chemical Composition

Counting with Moles

The mole is a fundamental unit for counting atoms, molecules, or ions in chemistry.

• Definition: 1 mole = particles (Avogadro's number).

• Example: 2.0 moles of helium contains atoms.

Mole, Atom, and Mass Conversions

Conversions between mass, moles, and number of particles are essential for chemical calculations.

• Grams to moles:

• Moles to atoms:

• Example: 24 g of carbon is 2.0 mol C.

Moles and Mass of Molecules

Use the molar mass of a compound to convert between mass and moles, and Avogadro's number to convert between moles and molecules.

• Example: 0.50 mol of H2O contains molecules.

Using Formulas to Convert

Subscripts in chemical formulas allow conversion between moles of compounds and moles of elements.

• Example: In CO2, 1 mole contains 1 mole of C and 2 moles of O.

Mass Percent Composition

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

• Formula:

• Example: Carbon in CO2 is 27.3% by mass.

Empirical and Molecular Formulas

The empirical formula shows the simplest ratio of elements, while the molecular formula shows the actual number of atoms.

• Empirical formula: Convert grams to moles, then find the simplest whole-number ratio (e.g., CH2O).

• Molecular formula: ; multiply empirical subscripts by n (e.g., C6H12O6).

CHAPTER 7: Chemical Reactions

Evidence of Chemical Reaction

Chemical reactions can be identified by observable changes.

• Signs: Color change, formation of a solid (precipitate), formation of a gas (bubbles), heat/light produced or absorbed.

• Example: Formation of bubbles indicates a chemical reaction; melting ice does not.

Chemical Equations and Balancing

Chemical equations represent reactants and products, and must be balanced to obey the law of conservation of mass.

• Balanced equation: Each element has the same number of atoms on both sides.

• Steps: Write correct formulas, balance elements appearing once first, balance free elements last, remove fractional coefficients.

• Example: Combustion of propane:

Solubility and Precipitation Reactions

Solubility rules help predict whether a substance will dissolve in water or form a precipitate.

• Soluble: Dissolves in water.

• Insoluble: Forms a solid (precipitate).

• Example: Mixing AgNO3(aq) and NaCl(aq) forms AgCl(s).

Molecular, Ionic, and Net Ionic Equations

Reactions in solution can be represented in three ways.

• Molecular equation: Shows compounds as whole units.

• Complete ionic equation: Shows all ions present.

• Net ionic equation: Shows only ions that participate in the reaction.

• Example:

  • Molecular: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

  • Complete ionic: Ag+ + NO3– + Na+ + Cl– → AgCl(s) + Na+ + NO3–

  • Net ionic: Ag+ + Cl– → AgCl(s)

Acid–Base and Gas-Evolution Reactions

Acid–base reactions produce water and a salt; gas-evolution reactions produce a gas.

• Acid–base: Acid + base → salt + water (e.g., HBr + KOH → KBr + H2O)

• Gas-evolution: Produces a gaseous product.

Redox Reactions

Redox (oxidation-reduction) reactions involve the transfer of electrons.

• Oxidation: Loss of electrons (increase in oxidation state).

• Reduction: Gain of electrons (decrease in oxidation state).

• Example: In 2Mg + O2 → 2MgO, Mg is oxidized (0 → +2), O is reduced (0 → –2).

Types of Chemical Reactions

Chemical reactions are classified by the changes occurring in reactants and products.

• Synthesis: A + B → AB

• Decomposition: AB → A + B

• Single displacement: A + BC → B + AC

• Double displacement: AB + CD → AD + CB

CHAPTER 8: Stoichiometry

Stoichiometric Calculations

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

• Mole ratios: Use coefficients from balanced equations to relate moles of reactants and products.

• Example: 2 H2 + O2 → 2 H2O; 4.0 mol H2 produces 4.0 mol H2O.

Mass Calculations

Convert between mass and moles using molar mass, then use stoichiometry to find product masses.

• Example: CH4 + 2 O2 → CO2 + 2 H2O; 50.0 g CH4 produces 137.5 g CO2.

Limiting Reactant and Yield

The limiting reactant determines the maximum amount of product formed. Percent yield measures efficiency.

• Limiting reactant: The reactant that runs out first and limits product formation.

• Example: N2 + 3 H2 → 2 NH3; with 3.0 mol N2 and 5.0 mol H2, H2 is limiting.

• Percent yield:

• Example: If theoretical yield is 3.33 mol and actual yield is 2.50 mol, percent yield is 75%.

Summary Table: Types of Chemical Formulas

Summary Table: Types of Chemical Reactions

Additional info: Academic context was added to expand brief points into full explanations, provide formulas, and clarify examples for self-contained study notes.