Exam 2 Study Guide: Chemical Formulas, Composition, Reactions, and Stoichiometry (Chapters 5-8)

Chapter 5: Molecules and Compounds

Chemical Formulas

Chemical formulas represent the types and numbers of atoms in a compound. They are essential for understanding the composition and properties of substances.

• Types of Formulas: - Empirical formula: Shows the simplest whole-number ratio of atoms. - Molecular formula: Shows the actual number of atoms of each element in a molecule. - Structural formula: Shows how atoms are connected.

• Polyatomic Ions: - Groups of atoms with a net charge that act as a single unit in compounds. - Examples: NO3- (nitrate), SO42- (sulfate).

• Molecular vs. Atomic Elements: - Atomic elements: Exist as single atoms (e.g., Na, Fe). - Molecular elements: Exist as molecules (e.g., O2, N2).

Ionic Compounds

Ionic compounds are formed from the electrostatic attraction between cations and anions.

• Writing Formulas: - Balance charges to ensure neutrality. - Example: $\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$

• Transition Metal Cations: - May have multiple charges; use Roman numerals to indicate charge (e.g., Fe2+ is iron(II)).

• Determining Charge: - Use periodic table and compound formula to deduce charge.

Molecular Compounds

Molecular compounds consist of nonmetals bonded together. Naming uses prefixes to indicate the number of atoms.

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

• Naming: - First element keeps its name; second element ends in "-ide". - Example: CO2 is carbon dioxide.

Formula Mass

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

• Calculation: - $\text{Formula mass} = \sum \text{(number of atoms)} \times \text{atomic mass}$

Chapter 6: Chemical Composition

The Mole

The mole is a fundamental unit in chemistry representing $6.022 \times 10^{23}$ particles (Avogadro's number).

• Conversion Factor: - Used to convert between atoms/molecules and moles.

• Calculations: - $\text{Number of moles} = \frac{\text{Number of particles}}{6.022 \times 10^{23}}$

Molar Mass

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol).

• Calculation: - Add atomic masses from the periodic table for all atoms in the formula.

• Conversions: - $\text{Mass (g)} = \text{Moles} \times \text{Molar mass (g/mol)}$

Percent Composition

Percent composition shows the percentage by mass of each element in a compound.

• Formula: - $\text{Percent composition} = \frac{\text{Mass of element in 1 mole}}{\text{Molar mass of compound}} \times 100\%$

Empirical and Molecular Formulas

Empirical formulas show the simplest ratio of elements; molecular formulas show the actual number of atoms.

• Determination: - Use percent composition and molar mass to find empirical and molecular formulas.

Chapter 7: Chemical Reactions

Types of Reactions

Chemical reactions can be classified into several types based on the reactants and products.

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

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

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

• Double Replacement: Exchange of ions between two compounds.

• Combustion: Reaction with oxygen producing heat and light.

Evidence of a Chemical Reaction

• Color change

• Formation of a precipitate

• Gas production

• Energy change (heat/light)

Chemical Equations

Chemical equations represent reactions using formulas and symbols.

• Balancing Equations: - Ensure the same number of each atom on both sides. - Example: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

• Physical States: - (s) solid, (l) liquid, (g) gas, (aq) aqueous

Solubility and Solutions

Solubility describes how well a substance dissolves in water. Solutions can be electrolytic or non-electrolytic.

• Electrolyte Solutions: Conduct electricity due to dissolved ions.

• Solubility Rules: Used to predict if an ionic compound will dissolve.

Precipitation Reactions

Occur when two solutions are mixed and an insoluble product (precipitate) forms.

• Predicting Products: Use solubility rules to determine if a precipitate forms.

• Net Ionic Equations: Show only the species that change during the reaction.

Chapter 8: Quantities in Chemical Reactions

Reaction Stoichiometry

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

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

• Mass-to-Mass Conversions: Convert mass to moles, use stoichiometry, then convert back to mass.

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Percent Yield: - $\text{Percent yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$

Calculating Reactant Left Over

After a reaction, some reactant may remain if it was in excess.

• Calculate how much of the excess reactant remains after the limiting reactant is used up.

Additional info:

• Greek prefixes for molecular compounds: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), deca- (10).

• Common polyatomic ions: NH4+ (ammonium), CO32- (carbonate), PO43- (phosphate).

• Solubility rules: Most nitrates, acetates, and group 1 salts are soluble; most carbonates, phosphates, and sulfides are insoluble except with group 1 or ammonium.