Compound Basics

Ionic and Covalent Compounds

Atoms combine to form compounds in order to achieve a stable number of valence electrons, often matching the electron configuration of noble gases. This stability is commonly referred to as the "happy" state for atoms.

• Ionic Compounds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in the formation of ions. The electrostatic attraction between oppositely charged ions (cations and anions) holds the compound together.

• Covalent Compounds: Formed by the sharing of electrons between nonmetals, resulting in discrete, neutral molecules.

Example: Sodium chloride (NaCl) is an ionic compound, while water (H2O) is a covalent compound.

Properties of Compounds

Big Idea 1: Unique Properties

Compounds have properties that are distinct from the elements that compose them. For example, hydrogen and oxygen are both gases at room temperature, but when combined to form water, the resulting compound is a liquid with very different properties.

Big Idea 2: Diversity of Compounds

With over 100 elements, countless combinations are possible, resulting in a vast array of molecules and compounds. Chemistry provides systematic ways to describe and represent these compounds.

Ionic Compounds

Formation and Structure

Ionic compounds are formed when a metal reacts with a nonmetal, transferring electrons from the metal to the nonmetal. Metals lose electrons to form cations (positively charged ions), while nonmetals gain electrons to form anions (negatively charged ions).

• Cation: An atom that loses one or more electrons, resulting in a positive charge.

• Anion: An atom that gains one or more electrons, resulting in a negative charge.

Example: A neutral sodium atom (Z = 11) loses one electron to become Na+. A neutral oxygen atom (Z = 8) gains two electrons to become O2-.

The electrostatic attraction between cations and anions forms a three-dimensional extended structure, characteristic of ionic solids.

Predicting Ion Charge

The periodic table helps predict the ionic charge of main-group elements:

• Group 1 metals: Lose one electron, form 1+ cations

• Group 2 metals: Lose two electrons, form 2+ cations

• Group 17 nonmetals: Gain one electron, form 1- anions

• Group 16 nonmetals: Gain two electrons, form 2- anions

Example: Calcium (Group 2) loses two electrons to form Ca2+, becoming isoelectronic with argon. Bromine (Group 17) gains one electron to form Br-, becoming isoelectronic with krypton.

Covalent Compounds

Covalent Bonds

Covalent bonds are formed when two nonmetals share electrons, resulting in discrete molecules. The shared electrons count toward the valence shell of both atoms, allowing each to achieve a stable configuration.

Molecular formula: Indicates the types and numbers of atoms in a molecule using chemical symbols and subscripts.

Structural formula: Shows how atoms are connected within the molecule.

Example: Methane (CH4) can be represented as a molecular formula, structural formula, ball-and-stick model, or space-filling model.

Nomenclature

Naming Ionic Compounds

Nomenclature is the system of rules for naming chemical compounds. Ionic compounds must be electrically neutral, with the total positive charge from cations balancing the total negative charge from anions.

• Name the cation first, then the anion.

• Monatomic cations: Use the element name (e.g., Na+ is sodium).

• Monatomic anions: Use the element name with the ending replaced by "-ide" (e.g., Cl- is chloride).

• Polyatomic ions: Use the established name (e.g., SO42- is sulfate).

Naming Ionic Compounds with Variable Charge Metals

Some metals, especially transition metals, can form cations with different charges. The charge is specified using Roman numerals in parentheses after the metal name (e.g., Fe2+ is iron(II)).

Naming Oxyanions

Oxyanions are polyatomic ions containing oxygen and another element. When a nonmetal forms two oxyanions:

• -ate: Ion with more oxygen atoms (e.g., SO42- is sulfate)

• -ite: Ion with fewer oxygen atoms (e.g., SO32- is sulfite)

If more than two oxyanions exist, prefixes are used:

• per-: Most oxygen atoms

• hypo-: Fewest oxygen atoms

Naming Acids

• Binary acids: Contain hydrogen and one other element. Named with the prefix "hydro-", the root of the nonmetal, the suffix "-ic", and the word "acid" (e.g., HCl is hydrochloric acid).

• Oxyacids: Contain hydrogen, oxygen, and another element. Named based on the oxyanion: replace "-ate" with "-ic" and "-ite" with "-ous" (e.g., H2SO4 is sulfuric acid, H2SO3 is sulfurous acid).

Naming Hydrates

Hydrates are compounds that contain water molecules within their crystal structure. The name includes the anhydrous compound followed by a Greek prefix indicating the number of water molecules and the word "hydrate" (e.g., CuSO4·5H2O is copper(II) sulfate pentahydrate).

Naming Binary Molecular (Covalent) Compounds

Binary molecular compounds consist of two nonmetals. The element farther left or lower in the periodic table is named first. The second element's name ends with "-ide". Greek prefixes indicate the number of atoms (mono-, di-, tri-, tetra-, penta-, etc.). The prefix "mono-" is omitted for the first element if only one atom is present.

Example: CO2 is carbon dioxide, N2O4 is dinitrogen tetroxide.

Molecular Mass and the Mole

The Mole Concept

The mole is a counting unit in chemistry, representing 6.022 × 1023 entities (Avogadro's number). It allows chemists to relate the mass of substances to the number of atoms, molecules, or ions present.

• 1 mole = 6.022 × 1023 particles

• Used for atoms, molecules, compounds, or ions

Molar Mass

The molar mass of a substance is the mass of one mole of that substance, expressed in grams per mole (g/mol). For elements, the molar mass is numerically equal to the atomic mass in amu.

• For molecules: Add the atomic masses of all atoms in the formula.

• For ionic compounds: Use the term formula mass instead of molecular mass, as ionic compounds do not exist as discrete molecules.

Example:

• Water (H2O): g/mol

• Carbon dioxide (CO2): g/mol

Empirical and Molecular Formulas

Definitions

• Empirical formula: Shows the simplest whole-number ratio of elements in a compound.

• Molecular formula: Shows the actual number of atoms of each element in a molecule.

Example: Hydrogen peroxide has the empirical formula HO and the molecular formula H2O2.

Determining Empirical Formulas

To determine the empirical formula from mass or percent composition:

1. Convert percent composition to mass (assume 100 g sample).

2. Convert masses to moles using molar masses.

3. Divide each mole value by the smallest number of moles.

4. If necessary, multiply by an integer to obtain whole-number subscripts.

Example: A compound contains 1.75 g H and 8.25 g O in a 10.0 g sample. Calculate moles and determine the ratio for the empirical formula.

Determining Molecular Formulas

To determine the molecular formula:

1. Find the empirical formula and its molar mass.

2. Divide the compound's molar mass by the empirical formula mass to get a whole number.

3. Multiply all subscripts in the empirical formula by this number.

Example: Butyric acid has a molar mass of 88.11 g/mol and an empirical formula determined from percent composition. Use the ratio to find the molecular formula.

Summary Table: Key Terms and Concepts

Additional info: These notes are based on the provided slides and summary text. Some details, such as specific examples and stepwise procedures, have been expanded for clarity and completeness.