Chemical Formulas and Types of Compounds

Counting Atoms in Chemical Formulas

Understanding chemical formulas is essential for determining the composition of compounds. Each element symbol is followed by a subscript indicating the number of atoms of that element in one formula unit.

• Key Point: If no subscript is present, only one atom of that element is present.

• Key Point: Parentheses indicate groups of atoms that appear more than once; multiply the subscript outside the parentheses by the subscript of each atom inside.

• Example: In Ca(NO3)2, there is 1 Ca, 2 N, and 6 O atoms.

Classifying Compounds: Ionic vs. Molecular

Compounds are classified based on the types of elements they contain and the nature of their bonding.

• Ionic Compounds: Composed of metals and nonmetals; consist of cations and anions held together by electrostatic forces.

• Molecular (Covalent) Compounds: Composed of nonmetals; atoms are held together by shared electrons.

• Example: NaCl (ionic), CO2 (molecular)

Writing Chemical Formulas

Formulas represent the simplest ratio of ions in ionic compounds or the actual number of atoms in molecular compounds.

• Ionic Compounds: Balance charges to ensure neutrality.

• Molecular Compounds: Use prefixes to indicate the number of each atom.

• Acids: Special naming conventions; binary acids (e.g., HCl), oxyacids (e.g., H2SO4).

Naming Compounds

• Ionic Compounds: Name the cation (metal) first, then the anion (nonmetal with -ide ending).

• Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) for both elements; second element ends with -ide.

• Acids: Binary acids: "hydro-" + root + "-ic acid"; Oxyacids: root + "-ic" or "-ous acid" depending on polyatomic ion.

• Example: CO2: carbon dioxide; Na2SO4: sodium sulfate; HNO3: nitric acid

Chemical Quantities and Calculations

Formula Mass and Molar Mass

The formula mass is the sum of the atomic masses of all atoms in a formula unit; molar mass is the mass of one mole of a substance.

• Formula Mass: Add atomic masses (from the periodic table) for all atoms in the formula.

• Molar Mass: Expressed in grams per mole (g/mol).

• Example: For H2O:

Dimensional Analysis: Grams, Moles, and Particles

Dimensional analysis uses conversion factors to relate grams, moles, and number of particles (atoms, molecules, formula units).

• Key Conversions:

  • Grams ↔ Moles: Use molar mass as a conversion factor.

  • Moles ↔ Particles: Use Avogadro's number ( particles/mol).

• Example: To convert 10.0 g H2O to molecules:

• Steps: Identify given and desired units, set up conversion factors, and solve.

Mass Percent Composition

Mass percent expresses the mass of each element in a compound as a percentage of the total mass.

• Formula:

• Example: For H in H2O:

• Application: Use mass percent as a conversion factor in stoichiometric calculations.

Empirical and Molecular Formulas

• Empirical Formula: Simplest whole-number ratio of atoms in a compound.

• Molecular Formula: Actual number of atoms; may be a multiple of the empirical formula.

• Determination: Use experimental data (masses or percentages) to find moles of each element, then the simplest ratio.

• Relationship: , where

Chemical Reactions and Equations

Evidence of Chemical Reactions

Chemical reactions involve the transformation of substances. Evidence includes:

• Color change

• Formation of a precipitate

• Gas evolution (bubbles)

• Temperature change

• Emission of light

Writing and Balancing Chemical Equations

• Balanced Equation: Same number of each type of atom on both sides; obeys the Law of Conservation of Mass.

• Steps: Write correct formulas, balance elements one at a time, adjust coefficients as needed.

• Example:

Solubility and Precipitation Reactions

Solubility rules determine whether a compound dissolves in water. Precipitation reactions form an insoluble product (precipitate).

• Solution: Homogeneous mixture of solute dissolved in solvent.

• Solubility Table: Used to predict if a compound is soluble.

• Spectator Ions: Ions that do not participate in the reaction.

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

Molecular, Complete Ionic, and Net Ionic Equations

• Molecular Equation: Shows all reactants and products as compounds.

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

• Net Ionic Equation: Shows only species that change during the reaction.

• Example:

Molecular: Complete Ionic: Net Ionic:

Acid–Base and Gas Evolution Reactions

• Acid–Base Reaction: Acid reacts with base to form water and a salt.

• Example:

• Gas Evolution Reaction: Reaction produces a gas (e.g., CO2, H2, SO2).

• Example:

Redox (Oxidation–Reduction) Reactions

• Redox Reaction: Involves transfer of electrons between species.

• Oxidation: Loss of electrons; Reduction: Gain of electrons.

• Oxidizing Agent: Causes oxidation (is reduced); Reducing Agent: Causes reduction (is oxidized).

• Example:

Combustion Reactions

• Combustion Reaction: Substance reacts with O2 to produce energy, CO2, and H2O (if hydrocarbon).

• Example:

Classification of Chemical Reactions

Chemical reactions can be classified into several types:

• Synthesis (Combination): Two or more substances combine to form one product.

• Decomposition: One substance breaks down into two or more products.

• Single Replacement (Displacement): One element replaces another in a compound.

• Double Displacement (Metathesis): Exchange of ions between two compounds.

• Combustion: Reaction with oxygen producing heat and light.

Additional info: For more practice, refer to textbook examples and lecture videos for step-by-step solutions to problems involving these concepts.