Compounds, Nomenclature, and Molecular Mass – General Chemistry 1A Chapter 3 Study Notes

Compounds: Basics and Types

Introduction to Compounds

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

• Valence electrons are the outermost electrons of an atom and are involved in chemical bonding.

• Most elements do not naturally have a full set of valence electrons, so they form compounds to achieve stability.

• There is a vast variety of compounds, each with unique properties.

Properties of Compounds

Big Idea 1: Unique Properties

Compounds have properties distinct from the elements that compose them. For example, water (H2O) is formed from hydrogen and oxygen, but its properties are very different from those of either element.

Selected Properties	Hydrogen	Oxygen	Water
Boiling Point	-253°C	-183°C	100°C
State at Room Temperature	Gas	Gas	Liquid
Flammability	Explosive	Necessary for combustion	Used to extinguish flame

Big Idea 2: Diversity of Compounds

With over 100 elements, the number of possible combinations results in an enormous variety of molecules and compounds. Chemistry provides systematic ways to describe and represent this diversity.

Types of Compounds

Ionic Compounds

Ionic compounds are formed when electrons are transferred from one atom to another, typically between metals and nonmetals.

• Metals lose electrons to form cations (positively charged ions).

• Nonmetals gain electrons to form anions (negatively charged ions).

• The electrostatic attraction between cations and anions holds the compound together, forming a crystal lattice.

Covalent (Molecular) Compounds

Covalent compounds are formed when two or more nonmetals share electrons, resulting in discrete, neutral molecules.

• Shared electrons count toward the valence shell of both atoms involved.

• The resulting bond is called a covalent bond.

Ions: Cations and Anions

Definitions and Examples

• Anion: An atom that gains one or more electrons, acquiring a negative charge. Example: An oxygen atom (Z = 8) gains two electrons to become O2−.

• Cation: An atom that loses one or more electrons, acquiring a positive charge. Example: A sodium atom (Z = 11) loses one electron to become Na+.

Predicting Ion Charge

Using the Periodic Table

The periodic table helps predict the charges of ions formed by main-group elements:

• Group 1 metals: lose 1 electron → 1+ charge

• Group 2 metals: lose 2 electrons → 2+ charge

• Group 17 nonmetals: gain 1 electron → 1− charge

• Group 16 nonmetals: gain 2 electrons → 2− charge

Atoms become isoelectronic with the nearest noble gas after gaining or losing electrons.

Polyatomic Ions

Definition and Examples

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying an overall charge.

• Example: Carbonate ion, CO32−

• Oxyanions are polyatomic ions containing oxygen (e.g., sulfate SO42−, nitrate NO3−).

Properties of Ionic Compounds

• Usually solids with high melting and boiling points.

• Conduct electricity when molten or dissolved in water (not as solids).

Covalent Bonds and Molecular Compounds

Formation and Representation

• Covalent bonds form when electrons are shared between nonmetals.

• Molecules are represented by molecular formulas (e.g., H2O), structural formulas, or models (ball-and-stick, space-filling).

Some elements exist naturally as molecules (e.g., H2, N2, O2).

Chemical Formulas

Molecular and Structural Formulas

• Molecular formula: Shows the types and numbers of atoms in a molecule (e.g., C2H6).

• Structural formula: Shows how atoms are connected.

Nomenclature (Naming Compounds)

Ionic Compounds

• Name the cation first, then the anion.

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

• Monatomic anions: use the element root + "-ide" (e.g., Cl− is chloride).

• Polyatomic ions: use the common name (e.g., SO42− is sulfate).

Ionic Compounds with Variable Charge Metals

• Some metals (mainly transition metals) can form more than one cation.

• Specify the charge with a Roman numeral in parentheses (e.g., FeCl2 is iron(II) chloride).

Oxyanions

• Oxyanions are named based on the number of oxygen atoms:

• "-ate" for more oxygens, "-ite" for fewer (e.g., sulfate SO42−, sulfite SO32−).

• Prefixes "per-" (most oxygens) and "hypo-" (least oxygens) are used for series (e.g., perchlorate ClO4−, hypochlorite ClO−).

Acids

• Binary acids: Contain hydrogen and one other element. Named as "hydro-" + root + "-ic acid" (e.g., HCl is hydrochloric acid).

• Oxyacids: Contain hydrogen, oxygen, and another element (usually as a polyatomic ion). Named by modifying the anion name: "-ate" becomes "-ic acid", "-ite" becomes "-ous acid" (e.g., H2SO4 is sulfuric acid, H2SO3 is sulfurous acid).

Hydrates

• Hydrates are ionic compounds containing water molecules in their crystal structure.

• Name: anhydrous compound + Greek prefix for number of waters + "hydrate" (e.g., CuSO4·5H2O is copper(II) sulfate pentahydrate).

Binary Molecular (Covalent) Compounds

• Composed of two nonmetals.

• Name the element farther left (or lower) on the periodic table first.

• Second element gets "-ide" ending.

• Use Greek prefixes to indicate the number of each atom (mono-, di-, tri-, tetra-, penta-, etc.). Omit "mono-" for the first element.

• Example: CO2 is carbon dioxide, N2O5 is dinitrogen pentoxide.

The Mole and Molar Mass

The Mole Concept

• One mole contains $6.022 \times 10^{23}$ entities (Avogadro's number).

• Used to count atoms, molecules, ions, or formula units in a sample.

Molar Mass

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

• For elements, molar mass (g/mol) numerically equals atomic mass (amu).

• For compounds, sum the atomic masses of all atoms in the formula.

Example: Molar mass of H2O = (2 × 1.008) + (1 × 16.00) = 18.02 g/mol

Formula Mass for Ionic Compounds

• Ionic compounds do not exist as discrete molecules; use formula mass instead of molecular mass.

• Formula mass is calculated by summing the atomic masses of the ions in the empirical formula.

Empirical and Molecular Formulas

Definitions

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

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

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

Determining Empirical Formulas

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

2. Convert mass to moles using molar mass.

3. Divide by the smallest number of moles to get the simplest ratio.

4. If necessary, multiply to get whole numbers.

Determining Molecular Formulas

1. Find the empirical formula and its 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.

Summary Table: Greek Prefixes for Molecular Compounds

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-
9	nona-
10	deca-

Key Equations

• Number of moles: $n = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Number of entities: $\text{Number} = n \times N_A$ where $N_A = 6.022 \times 10^{23}$

• Percent composition: $\%\,\text{element} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100\%$

Additional info: Some context and examples have been expanded for clarity and completeness, and tables have been reconstructed for study purposes.