Chemistry Empirical and Molecular Formulas Quiz

Chapter 7 Objectives

• Explain the significance of a chemical formula.
• Determine the formula of an ionic compound formed between two given ions.
• Name an ionic compound given its formula.
• Name a binary molecular compound from its formula using prefixes.
• Write the formula of a binary molecular compound given its name.

Significance of a Chemical Formula

• A chemical formula indicates the relative number of atoms of each kind in a chemical compound.
• For a molecular compound, the chemical formula reveals the number of atoms of each element contained in a single molecule of the compound.
• For example, octane (C₈H₁₈) indicates 8 carbon atoms and 18 hydrogen atoms per molecule.
• For an ionic compound, the chemical formula represents one formula unit—the simplest ratio of the compound's positive ions (cations) and its negative ions (anions).
• For example, aluminum sulfate (Al₂(SO₄)₃) has 2 aluminum ions and 3 sulfate polyatomic ions (polyatomic ion SO₄₂-). The parentheses around SO₄ indicate that it is a single unit.

Monatomic Ions

• Many main-group elements can lose or gain electrons to form ions.
• Ions formed from a single atom are called monatomic ions.
• For example, nitrogen (N) can gain three electrons to form the nitride ion (N³⁻).
• Monatomic cations retain the element's name (e.g., potassium cation K⁺).
• Monatomic anions use the root name of the element followed by the suffix -ide (e.g., fluoride anion F⁻).

Common Monatomic Ions

• A table of common monatomic ions is provided, listing group 1, 2, and 13 cations and corresponding anionic forms. These include, Li+, Na+, K+, Rb+, Cs+, Be²⁺, Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺, Al³⁺, F⁻, O²⁻, S²⁻, N³⁻, P³⁻.

d-block elements

• A table includes common monatomic ions for d-block elements, showing their multiple possible charges and their corresponding names (e.g., copper(I), copper(II), iron(II), iron(III) etc.)

Binary Ionic Compounds

• Compounds composed of two elements are known as binary compounds.
• In a binary ionic compound, the total number of positive charges must equal the total number of negative charges.
• The formula for a binary ionic compound can be written given the identities of the compound's ions (e.g., magnesium bromide MgBr₂––Mg²⁺ + 2Br⁻).
• A general rule for determining binary ionic compound formulas is to 'cross over' the charges (e.g., aluminum oxide Al₂O₃––Al³⁺ + 2O²⁻).

Binary Ionic Compounds, Continued

• Include examples of calculating formulas for binary ionic compounds, checking the combined charges to see if they are equal.

Writing Formulas of an Ionic Compound

• Steps to write the formula of an ionic compound (e.g., iron(III) oxide):
  • Write the symbol and charges of the cation and anion.
  • Write the symbols for the ions side by side, beginning with the cation.
  • Determine the lowest common multiple of the charges on the ions.
  • For the example of iron (III) oxide the formula is Fe₂O₃.

Naming Binary Ionic Compounds

• The nomenclature system for naming binary ionic compounds involves combining the names of the compound's positive (cation) and negative ions (anion), giving the name of the cation first, followed by the name of the anion (e.g. Al₂O₃ – aluminum oxide).
• For simple ionic compounds, the ratio of the ions is not given in the name, as it's implied by the cation and anion charges (e.g. Potassium Iodide).

Naming Binary Ionic Compounds, Continued

• The Stock System is used when elements can form more than one cation with different charges. A Roman numeral in parentheses following the cation name indicates the ion’s charge (e.g., iron(II) oxide vs. iron(III) oxide).

Naming Binary Ionic Compounds, Continued: Sample Problems

• Includes sample problems for writing formulas and names of binary ionic compounds containing polyatomic ions. Illustrates how to determine formulas and names using the correct principles and rules.

Naming Binary Ionic Compounds Containing Polyatomic Ions

• Many common polyatomic ions are oxyanions, meaning they contain oxygen.
• Oxyanions with a greater number of oxygen atoms end in -ate (e.g., nitrate NO₃⁻).
• Oxyanions with a smaller number of oxygen atoms end in -ite (e.g., nitrite NO₂⁻).
• Some elements can form more than two types of oxyanions, and then the prefix hypo- or per- is used (e.g., hypochlorite ClO⁻, chlorite ClO₂⁻, chlorate ClO₃⁻, and perchlorate ClO₄⁻).
• Includes a table of polyatomic ions.

Naming Binary Ionic Compounds, Continued: Sample Problems

• Includes a variety of problems to illustrate and use the principles to name and write formulas in an ionic context.

Naming Binary Molecular Compounds

• Molecular compounds are composed of individual covalently bonded units (molecules).
• The naming system for binary molecular compounds uses prefixes to indicate the number of each type of atom (e.g., carbon monoxide CO, carbon dioxide CO₂).
• Includes a table of prefixes indicating the numbers of atoms.

Naming Binary Molecular Compounds, Continued: Solved Examples

• Includes problems in the same format as the naming portion above.

Covalent-Network Compounds

• Some covalent compounds form extended three-dimensional networks; their formulas represent the smallest whole-number ratio of atoms present (e.g., silicon dioxide SiO₂).

Acids and Salts

• An acid is a molecular compound that usually contains hydrogen and releases hydrogen ions (e.g., hydrochloric acid HCI).
• Oxyacids contain hydrogen, oxygen, and another element (usually a nonmetal), such as sulfuric acid (H₂SO₄), and nitric acid (HNO₃).
• An ionic compound composed of a cation and the anion from an acid is known as a salt (e.g., sodium chloride NaCl).

Oxidation Numbers

• Oxidation numbers are assigned to atoms to track electron distribution.
• In a neutral compound, the sum of the oxidation numbers is zero. In a polyatomic ion, the sum is equal to the ion's charge.
• Rules for assigning oxidation numbers are provided, including elements in their natural state having an oxidation number of zero, fluorine having an oxidation number of -1, oxygen normally having an oxidation number of -2 (except in peroxides), and hydrogen having an oxidation number of +1 in all compounds except when bonded to a metal, where it is -1.
• Solved examples for determining oxidation numbers are provided through calculation methods, including examples utilizing the rules.

Using Oxidation Numbers for Formulas and Names

• The oxidation numbers of nonmetals can be used as if they were ionic charges to determine formulas (e.g. sulfur dioxide SO₂).
• Includes a table of common oxidation states of various nonmetals.

Molar Mass as a Conversion Factor

• Molar mass is the mass in grams of one mole of a substance.
• It's used as a conversion factor between mass in grams and amount in moles.

Mole-Mass Calculations

• Illustrates how to convert mass to moles and moles to mass, and relating that to number of molecules, formula units, or ions.
• Examples of calculating are given.

Percentage Composition

• Percentage composition is the percentage by mass of each element in a compound.
• To obtain percentage composition, the mass of each element in one mole of the compound is divided by the molar mass of the compound, multiplied by 100.

Calculation of Empirical Formulas

• The empirical formula represents the smallest whole-number mole ratio of the elements in a compound.
• Examples are provided illustrating the process of calculating empirical formulas from mass composition or percentage compositions (e.g., obtaining the empirical formula of a compound containing a percentage composition for elements).

Calculation of Molecular Formulas

• The molecular formula is the actual formula of a molecular compound.
• It's related to the empirical formula: x(empirical formula) = molecular formula.
• The relationship through formula masses can be observed in the given examples.

Comparing Empirical and Molecular Formulas

• Discusses the relationship and differences between empirical and molecular formulas, and examples across various compounds demonstrating the concept.

Additional Notes (Chapter 9 Stoichiometry)

• Explains Stoichiometry definition, and mole ratios
• Includes solved examples on mole/mass conversions, limiting reactants, calculating masses of products

Test your knowledge on empirical formulas and percentage compositions in chemistry. This quiz covers concepts such as mole ratios, oxidation numbers, and the percentage of elements in compounds. Perfect for students studying chemistry at the high school level.