Chemical Bonds and Compounds: Introduction to Ionic and Covalent Chemistry

Chemical Bonds

Definition and Types

Chemical bonds are the forces that hold atoms and ions together in compounds. There are two main types of chemical compounds: ionic and covalent. The formation of these bonds is primarily due to the behavior of valence electrons, which are the electrons in the outermost shell of an atom.

• Ionic bonds: Formed when atoms transfer electrons; one atom gains electrons (becomes an anion), another loses electrons (becomes a cation).

• Covalent bonds: Formed when atoms share electrons between them.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a stable configuration of 8 valence electrons, similar to noble gases.

Example: Sodium (Na) transfers an electron to chlorine (Cl) to form sodium chloride (NaCl), an ionic compound.

Ionic Compounds

Formation and Naming

Ionic compounds are formed from the electrostatic attraction between cations (positively charged ions) and anions (negatively charged ions). Most often, these compounds consist of a metal and a nonmetal.

• Cations are always listed first in the formula, followed by anions.

• The sum of the charges in an ionic compound must be zero.

• If ions have the same charge magnitude, use one of each. If charges differ, use the crossover method to balance charges (reduce subscripts if needed).

• Naming is based on the ions present. Metals are named as the element; nonmetals often end with '-ide'.

• Some transition metals can have multiple charges; their charge is indicated in the compound name (e.g., iron(III) chloride).

Example: Magnesium chloride: Mg2+ and Cl- combine to form MgCl2.

Polyatomic Ions

Common Polyatomic Ions and Their Properties

Polyatomic ions are ions composed of more than one atom. While a few are cations, most are anions. Their names and formulas do not change when forming compounds.

Name	Formula	Name	Formula
Hydronium ion	H3O+	Nitrate ion	NO3-
Ammonium ion	NH4+	Sulfate ion	SO42-
Hydroxide ion	OH-	Phosphate ion	PO43-

• Names of polyatomic ions remain unchanged in compounds.

• If more than one polyatomic ion is needed in a formula, enclose the ion in parentheses and add the subscript (e.g., Ca(NO3)2).

Example: Sodium sulfate: Na2SO4

Covalent Bonding

Formation and Properties

Covalent bonds are formed when atoms share electrons, typically between two nonmetals. A single unit of a covalent compound is called a molecule. The shared electrons are represented by lines between atoms in structural formulas.

• Each line represents two shared electrons (a single bond).

• Double and triple bonds involve four or six shared electrons, respectively.

• Unshared electrons are called lone pairs.

• Some elements form molecules with themselves, known as diatomic molecules (e.g., O2, N2).

• Electron sharing is not always equal; differences in electronegativity create polarity.

• If atoms are identical, the bond is nonpolar; if not, the bond may be polar covalent.

Example: Water (H2O) has polar covalent bonds due to the difference in electronegativity between hydrogen and oxygen.

Naming Binary Covalent Compounds

Rules and Prefixes

The International Union of Pure and Applied Chemistry (IUPAC) sets the rules for naming covalent compounds. Names are based on the elements and the number of atoms present, using specific prefixes.

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

• First element: name with prefix (except mono- is not used for the first element).

• Second element: name with prefix and '-ide' ending.

• If the prefix ends in 'o' or 'a' and the element name starts with a vowel, drop the final vowel of the prefix.

Example: CO2 is named carbon dioxide.

Lewis Structures

Drawing and Interpreting Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the arrangement of valence electrons in a molecule. They help visualize the octet rule and predict molecular shapes.

• Sum the total number of valence electrons for all atoms in the molecule.

• Draw a line (bond) between each pair of connected atoms.

• The atom that needs the most electrons (usually the least electronegative) goes in the center.

• Place lone pairs around the outer atoms (except hydrogen).

• Place remaining electrons around the central atom.

Example: The Lewis structure for methane (CH4) shows four single bonds between carbon and hydrogen, with no lone pairs on carbon.

Molecule Shape and Polarity

VSEPR Theory and Determining Polarity

Lewis structures show bonds and lone pairs, but not the three-dimensional shape of molecules. The Valence Shell Electron-Pair Repulsion (VSEPR) theory is used to predict molecular geometry based on the repulsion between electron sets around the central atom.

• Count the number of bonded atoms and lone pairs on the central atom.

• Compare to VSEPR tables to identify the molecular shape (e.g., linear, bent, trigonal planar, tetrahedral).

• To determine polarity:

  • If the compound is only one element or only C and H, it is nonpolar.

  • If the central atom has a lone pair, the molecule is polar.

  • If the outer atoms are not all the same, the molecule is polar.

  • If none of the above are true, the molecule is nonpolar.

Example: Ammonia (NH3) is polar due to a lone pair on nitrogen and three hydrogen atoms.