Ionic and Covalent Compounds: Structure, Nomenclature, and Formulas

Compounds: Definitions and Types

Substance and Compound

In chemistry, a substance is a form of matter with uniform properties. Substances can be classified as elements or compounds:

• Element: A substance that cannot be separated into simpler substances by chemical means.

• Compound: A substance composed of two or more elements chemically combined in fixed proportions.

Compounds can be separated into simpler substances by physical processes.

Lewis Electron Dot Formulas

Valence Electrons and Lewis Structures

Lewis electron dot formulas represent the valence electrons of atoms, which are involved in chemical bonding.

• Only valence electrons are shown; core electrons are not involved in bonding.

• Gilbert N. Lewis introduced the dot symbol method for tracking electrons in bonding.

• Each electron is represented as a dot around the element symbol.

• Electrons are placed on four sides of a square (top, bottom, left, right) before pairing.

• Elements in the same group have the same valence electron configuration.

Element	Lewis Dot Symbol
Na	Na•
Cl	Cl•••••••
O	O••••••
N	N•••••
C	C••••

Additional info: Table inferred from periodic table and Lewis dot conventions.

Ionic Bonding

Formation and Properties

An ionic bond is a chemical bond formed by the electrostatic attraction between positive cations and negative anions.

• Ionic bonds result from the complete transfer of one or more electrons from a metal atom to a nonmetal atom.

• Compounds with ionic bonds are typically formed between metals and nonmetals.

• Example: Sodium chloride (NaCl) forms when sodium loses an electron to become Na+ and chlorine gains an electron to become Cl-.

Electron Transfer Example:

Na → Na+ + e- Cl + e- → Cl-

Sodium loses an electron to become a cation; chlorine gains an electron to become an anion.

• Ionic solids form crystalline structures called crystal lattices.

• The arrangement of ions maximizes attraction and minimizes repulsion.

Lattice Energy (LE)

Definition and Calculation

Lattice energy is the energy required to convert a mole of ionic solid into its constituent ions in the gas phase.

• The force of attraction between ions determines the lattice energy.

• The larger the lattice energy, the more stable the ionic compound.

• Lattice energy is also the energy released when gaseous ions combine to form the lattice.

Coulomb's Law:

The magnitude of lattice energy is given by:

$LE = \dfrac{Q_1 Q_2}{r}$

• As the charge on the ions increases, LE increases.

• As the distance between the ions decreases, LE increases.

Naming Simple Compounds

Systematic Nomenclature

Modern chemistry uses systematic naming conventions to identify compounds:

• Monatomic ions: Main group cations keep their element name (e.g., Al3+ is aluminum ion).

• Monatomic anions replace the ending with "-ide" (e.g., Cl- is chloride).

• Monatomic transition metal ions: The charge is indicated in parentheses using Roman numerals (e.g., Fe2+ is iron(II) ion).

• Older names use "-ic" for higher charge and "-ous" for lower charge (e.g., Fe3+ is ferric, Fe2+ is ferrous).

Metal	Archaic Name	Common Charges
Fe	Ferrum	+2, +3
Cu	Cuprum	+1, +2
Sn	Stannum	+2, +4
Au	Aurum	+1, +3
Pb	Plumbum	+2, +4

Writing Ionic Compounds

Formulas and Balancing Charges

• To write the formula, combine ions so that the total charge is zero.

• The cation is written first, followed by the anion.

• Use the smallest whole number ratio of ions.

Example:

K+ + I- → KI (potassium iodide)

Ca2+ + O2- → CaO (calcium oxide)

Balance charges by using the absolute value of the superscript as the subscript for the other ion.

Binary Ionic Compounds

Naming and Examples

• Binary ionic compounds consist of two elements: a metal and a nonmetal.

• The metal (cation) is named first, followed by the nonmetal (anion) with "-ide" ending.

• Example: Na2S is sodium sulfide.

Covalent Bonding and Molecules

Formation and Representation

Covalent bonds are formed when two nonmetals share electrons.

• Each pair of shared electrons is represented as a straight line (single bond).

• The nuclei of atoms attract their own electrons and those of neighboring atoms, pulling atoms together.

• If atoms get too close, their nuclei repel each other.

Molecules and the Law of Definite Proportions

• A molecule is a pure substance containing two or more different elements in a fixed ratio.

• The Law of Definite Proportions states that a compound always contains the same elements in the same ratio by mass.

Example: Water (H2O) always contains two hydrogen atoms and one oxygen atom.

The 2:1 ratio of H:O by atoms corresponds to an 8:1 ratio by mass.

$\frac{2 \times 1.008\,\text{g}}{16.00\,\text{g}} = 0.126$

$\frac{16.00\,\text{g}}{2 \times 1.008\,\text{g}} = 7.94$

Law of Multiple Proportions

• If two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

Example: H2O and H2O2

$\frac{16.00}{1.008} = 15.87$ (H2O) $\frac{32.00}{1.008} = 31.75$ (H2O2)

Ratio: $\frac{31.75}{15.87} = 2$

Types of Molecules

• Diatomic molecules: Contain only two atoms (e.g., H2, N2, O2, F2, Cl2, Br2, I2).

• Polyatomic molecules: Contain more than two atoms.

• Homoelement molecules: Composed of atoms of the same element.

• Heteroelement molecules: Composed of atoms of different elements.

Molecular and Empirical Formulas

Definitions and Models

• Molecular formula: Shows the exact number of atoms of each element in a molecule (e.g., C6H12O6).

• Empirical formula: Shows the simplest whole number ratio of atoms in a compound (e.g., CH2O for glucose).

• Ball-and-stick models show connections between atoms; space-filling models show relative sizes.

Naming Binary Molecular Compounds

Rules and Prefixes

• Name the first element as it is; the second element ends with "-ide".

• Use Greek prefixes to indicate the number of atoms:

Prefix	Number
Mono-	1
Di-	2
Tri-	3
Tetra-	4
Penta-	5
Hexa-	6
Hepta-	7
Octa-	8
Nona-	9
Deca-	10

Examples: BrCl3 is bromine trichloride; N2O5 is dinitrogen pentoxide.

Binary Molecular Compounds with Hydrogen

• Binary molecular compounds with hydrogen do not always follow the standard rules (e.g., NH3 is ammonia, H2O is water).

• Binary compounds between hydrogen and group 17 elements are named as acids when dissolved in water (e.g., HCl is hydrochloric acid).

Binary Compound	Acid Name
HF	Hydrofluoric acid
HCl	Hydrochloric acid
HBr	Hydrobromic acid
HI	Hydroiodic acid

Polyatomic Ions

Definition and Nomenclature

• Polyatomic ions are ions composed of two or more atoms covalently bonded, acting as a single charged unit.

• Compounds with polyatomic ions may involve both ionic and covalent bonding.

• Negatively charged polyatomic ions often contain oxygen (oxyanions).

• If more than one polyatomic ion is present, use parentheses and a subscript (e.g., Al2(CO3)3).

Name	Formula
Carbonate	CO32-
Nitrate	NO3-
Sulfate	SO42-
Phosphate	PO43-

Naming Oxyanions

• "-ate" ending for the ion with more oxygen; "-ite" for less oxygen.

• "Per-" prefix for one more oxygen than "-ate"; "hypo-" for one less than "-ite".

Example: ClO4- is perchlorate, ClO3- is chlorate, ClO2- is chlorite, ClO- is hypochlorite.

Acid Nomenclature

• Compounds between hydrogen and anions are acids.

• For monatomic anions, "-ide" changes to "-ic" and add "hydro-" prefix (e.g., HCl is hydrochloric acid).

• For oxyanions, "-ate" changes to "-ic" and "-ite" changes to "-ous" (e.g., HNO3 is nitric acid, HNO2 is nitrous acid).

Oxyanion	Acid
CO32-	H2CO3 (carbonic acid)
NO3-	HNO3 (nitric acid)
SO42-	H2SO4 (sulfuric acid)

Hydrates

• Some compounds contain water molecules within their solid structure, called hydrates.

• Hydrates are named using prefixes to indicate the number of water molecules (e.g., CuSO4·5H2O is copper(II) sulfate pentahydrate).

• When water is removed, the compound is called anhydrous.

Chemical Formulas and Molar Mass

Calculating Molar Mass

• The subscript in a formula gives the relative number of each atom in a compound.

• Molar mass is the sum of the atomic masses of all atoms in a molecule or formula unit.

Example:

Molar mass of CH4: $M = 12.01\,\text{g/mol} + 4 \times 1.008\,\text{g/mol} = 16.04\,\text{g/mol}$

Molar mass of Ca(NO3)2: $M = 40.08\,\text{g/mol} + 2 \times [14.01\,\text{g/mol} + 3 \times 16.00\,\text{g/mol}] = 164.10\,\text{g/mol}$

Percent Composition from Chemical Formula

Calculating Percent by Mass

• Percent by mass of each element in a compound can be determined from its formula.

• Divide the mass of each element by the total molar mass and multiply by 100%:

$\%\,\text{Composition} = \frac{\text{mass of element in 1 mole}}{\text{molar mass of compound}} \times 100\%$

Chemical Formula from Percent Composition

Steps to Determine Empirical and Molecular Formulas

1. Determine the mass of each element present in the sample.

2. Convert the mass of each element to moles.

3. Divide all moles by the smallest number of moles.

4. If necessary, multiply by a factor to obtain whole numbers.

5. To determine the molecular formula, divide the molar mass by the empirical formula mass.

Example: If a compound contains 70.00% C and 3.36% H by mass, and the molar mass is 240.21 g/mol, calculate the empirical and molecular formula.

Additional info: All formulas, tables, and examples have been expanded and clarified for academic completeness.