Ionic and Covalent Compounds: Structure, Nomenclature, and Calculations

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Ionic and Covalent Compounds

Compounds

A compound is a substance composed of two or more elements combined in a specific ratio and held together by chemical bonds. Common examples include water (H2O) and sodium chloride (NaCl).

Key Point: Compounds have fixed ratios of elements.
Example: Water always contains two hydrogen atoms and one oxygen atom.

Spring water and Morton salt as examples of compounds

Lewis Dot Symbols

Lewis dot symbols represent the valence electrons of an atom as dots around the element's symbol. These are used to visualize electron interactions during bond formation.

Key Point: Each dot represents a valence electron.
Example: Boron has three valence electrons, shown as three dots.

Lewis dot symbol for boron Lewis dot symbols for main group elements Lewis dot symbols for B, C, N

Lewis Dot Symbols for Ions

Ions are also represented by Lewis dot symbols, showing the gain or loss of electrons and the resulting charge.

Key Point: Cations have fewer dots; anions have eight dots and are shown in brackets with the charge.
Example: Na+ has no dots; O2− has eight dots.

Lewis dot symbols for Na+ and O2- ions

Ionic Compounds and Bonding

Ionic Bonding

Ionic bonding is the electrostatic attraction between oppositely charged ions, forming ionic compounds. Electrons are transferred from metals to nonmetals.

Key Point: Ionic compounds are electrically neutral.
Example: Formation of NaCl from Na and Cl.

Formation of sodium chloride crystal from Na and Cl ions

Lattice Energy

Lattice energy is the energy required to separate one mole of an ionic solid into its constituent ions in the gas phase. It is a measure of the stability of an ionic compound.

Key Point: Lattice energy increases with higher charge and smaller ionic radius.
Formula:
Example: LiI has higher lattice energy than NaI or KI due to smaller ionic radius.

Coulomb's law for lattice energy Lattice energies of selected ionic compounds

Naming Ions and Ionic Compounds

Naming Ions

Monatomic cations: Add 'ion' to the element name (e.g., sodium ion).
Monatomic anions: Change the ending to '-ide' (e.g., chloride).
Variable charge metals: Use Roman numerals (e.g., iron(II) ion).

Formulas of Ionic Compounds

Ionic compounds must be electrically neutral. The sum of charges in the formula must equal zero.

Example: Aluminum oxide is Al2O3.

Formation of aluminum oxide from Al3+ and O2- ions

Naming Ionic Compounds

Name the cation (omit 'ion', use Roman numeral if needed).
Name the anion (omit 'ion').
Example: NaBr is sodium bromide; FeCl2 is iron(II) chloride.

Flowchart for naming ionic compounds

Covalent Bonding and Molecules

Covalent Bonding

In covalent bonding, electrons are shared between atoms to achieve noble gas configurations. This is described by the Lewis theory of bonding.

Key Point: Covalent bonds involve shared pairs of electrons.
Example: H2 molecule forms by sharing electrons.

Molecules and Laws of Proportions

Molecule: Combination of at least two atoms held by chemical bonds.
Law of Definite Proportions: Same compound always has same element ratio.
Law of Multiple Proportions: Elements can combine in different ratios to form different compounds.
Example: CO2 and CO have different O:C ratios.

O to C ratio in CO2 and CO

Types of Molecules

Diatomic molecules: Two atoms (homonuclear or heteronuclear).
Polyatomic molecules: More than two atoms.

Homonucler diatomic, heteronuclear diatomic, and polyatomic molecules

Molecular and Empirical Formulas

A molecular formula shows the exact number of atoms in a molecule. An empirical formula shows the simplest whole-number ratio of elements.

Example: Ethanol's molecular formula is C2H6O; empirical formula is CH3O.
Allotropes: Different forms of the same element (e.g., O2 and O3).

Ball-and-stick model of ethanol Table of molecular and empirical formulas

Naming Molecular Compounds

Binary Molecular Compounds

Name the first element.
Name the second element, changing its ending to '-ide'.
Use Greek prefixes to denote the number of atoms (mono-, di-, tri-, etc.).
Example: N2O5 is dinitrogen pentoxide.

Compounds Containing Hydrogen

Acids: Remove '-gen' from hydrogen, change '-ide' to '-ic' (e.g., hydrogen chloride → hydrochloric acid).
Must contain at least one ionizable hydrogen atom.

Organic Compounds

Hydrocarbons and Alkanes

Organic compounds contain carbon and hydrogen, sometimes with other atoms. Hydrocarbons are compounds of only carbon and hydrogen. Alkanes are the simplest hydrocarbons.

Example: Methane (CH4), Ethane (C2H6), Propane (C3H8).

Table of simple alkanes Table of simple alkanes continued

Functional Groups

Functional groups are specific groups of atoms within molecules that determine their chemical reactivity.

Example: Alcohol (–OH), Aldehyde (–CHO), Carboxylic acid (–COOH), Amine (–NH2).

Table of organic functional groups

Covalent Bonding in Ionic Species

Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge. They combine in ratios to form neutral compounds.

Example: Calcium phosphate, Ca3(PO4)2.

Formation of calcium phosphate from Ca2+ and PO43- ions

Common Polyatomic Ions

Name	Formula/Charge
Ammonium	NH4+
Hydronium	H3O+
Acetate	C2H3O2−
Carbonate	CO32−
Nitrate	NO3−
Sulfate	SO42−
Phosphate	PO43−
Hydroxide	OH−

Table of common polyatomic ions

Oxoanions and Oxoacids

Oxoanions: Polyatomic anions containing oxygen and another element.
Naming: -ate (most O), -ite (one less O), per- (one more O), hypo- (two less O).
Oxoacids: Acids formed from oxoanions; -ic for -ate, -ous for -ite.
Example: HClO3 is chloric acid; HClO2 is chlorous acid.

Hydrates

Hydrates are compounds with a specific number of water molecules in their structure. When water is removed, the compound is called anhydrous.

Example: CuSO4·5H2O is copper(II) sulfate pentahydrate; anhydrous CuSO4 is white, hydrated is blue.

Anhydrous and hydrated copper(II) sulfate

Molecular and Formula Mass

Molecular Mass

The molecular mass is the sum of atomic masses of all atoms in a molecule, measured in atomic mass units (amu).

Formula:
Example: Propane (C3H8): amu

Formula Mass

For ionic compounds, the formula mass is calculated similarly, using the empirical formula.

Percent Composition of Compounds

Percent composition is the percent by mass of each element in a compound.

Formula:
Example: Lithium carbonate (Li2CO3): Calculate percent by mass for Li, C, and O.

Molar Mass and Interconversions

Molar Mass

Molar mass (M) is the mass in grams of one mole of a substance. For elements, it is numerically equal to the atomic mass.

Example: 1 mol C = 12.01 g

Interconverting Mass, Moles, and Number of Particles

Conversions between mass, moles, and number of particles use molar mass and Avogadro's number.

Formula:
Formula:

Flowchart for converting mass, moles, and particles

Determination of Empirical and Molecular Formulas from Percent Composition

Empirical formulas are determined from percent composition by converting mass percentages to moles and finding the simplest ratio. Molecular formulas are found by dividing the molar mass by the empirical formula mass and multiplying the empirical formula subscripts by this number.

Example: A compound with 30.45% N and 69.55% O, molar mass 92 g/mol, has empirical formula NO2 and molecular formula N2O4.

Summary of Key Points

Compounds are formed from elements in fixed ratios.
Lewis dot symbols visualize valence electrons and bonding.
Ionic compounds are formed by electron transfer; covalent compounds by electron sharing.
Nomenclature rules differ for ionic and molecular compounds.
Polyatomic ions and hydrates have special naming conventions.
Molecular and formula masses, percent composition, and molar mass are essential for quantitative chemistry.
Empirical and molecular formulas can be determined from percent composition and molar mass.