Structure and Properties of Ionic and Covalent Compounds

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Chemical Bonding

Definition and Nature of Chemical Bonds

Chemical bonding is fundamental to the structure and properties of all compounds. A chemical bond is the force of attraction between any two atoms in a compound. This attractive force must overcome the repulsion between the positively charged nuclei of the two atoms involved. Chemical bonds are primarily the result of interactions involving valence electrons.

Valence electrons are the outermost electrons of an atom and are responsible for chemical bonding.
Only valence electrons participate in bonding, making it easier to apply the octet rule.

Lewis Symbols

Representation of Atoms and Valence Electrons

The Lewis symbol is a way to represent atoms using the element symbol and dots to indicate valence electrons. This method simplifies the visualization of bonding and electron arrangement.

The number of dots equals the number of valence electrons in the atom's outermost shell.
Each of the four sides around the atomic symbol can have up to two dots, for a maximum of eight (the octet).
Unpaired dots (unpaired valence electrons) are available to form chemical bonds.

Example: The Lewis symbol for oxygen (O) with six valence electrons: O with six dots arranged around it.

Principal Types of Chemical Bonds

Ionic and Covalent Bonds

There are two main types of chemical bonds: ionic and covalent.

Ionic bond: Formed by the transfer of one or more electrons from one atom (usually a metal) to another (usually a nonmetal). The resulting ions (cations and anions) are held together by electrostatic attraction.
Covalent bond: Formed by the sharing of electrons between atoms (typically nonmetals).
Some bonds have characteristics of both types and are not easily classified as purely ionic or covalent.

Ionic Bonding

Formation and Properties

Ionic bonding typically occurs between metals and nonmetals. Metals lose electrons to form cations, while nonmetals gain electrons to form anions. Both achieve a noble gas electron configuration.

Electrons are lost by a metal and gained by a nonmetal.
The resulting cation and anion are attracted to each other, forming an ionic bond.
Compounds of opposite charge aggregate into a crystal lattice.

Example: Formation of NaCl

Sodium (Na) loses one electron to become Na+.
Chlorine (Cl) gains one electron to become Cl-.
The reaction:

Covalent Bonding

Formation and Properties

Covalent bonds form when two atoms share one or more pairs of electrons. This type of bonding is common between nonmetals.

Each atom attains a noble gas configuration by sharing electrons.
Compounds with covalent bonds are called molecules or covalent compounds.
Diatomic elements (e.g., H2, O2, N2, F2, Cl2, Br2, I2) have completely covalent bonds with equal sharing.

Example: Formation of H2 molecule:

Polar Covalent Bonding and Electronegativity

Bond Polarity and Electronegativity

Polar covalent bonds occur when electron pairs are shared unequally between atoms due to differences in electronegativity (the ability of an atom to attract electrons in a bond).

The greater the difference in electronegativity, the more polar the bond.
Electronegativity increases across a period and decreases down a group in the periodic table.

Example: The H–F bond is more polar than the H–Cl bond because the electronegativity difference is greater for H–F.

Electronegativity difference for H–F:
Electronegativity difference for H–Cl:

Naming Compounds and Writing Formulas

Nomenclature Systems

Chemical nomenclature assigns unique names to compounds. There are different systems for ionic and covalent compounds.

Ionic compounds: Name the cation (metal) first, then the anion (nonmetal) with the suffix -ide.
If the metal can form more than one ion (e.g., transition metals), use Roman numerals to indicate the charge (Stock system), or use -ic/-ous endings (common system).
Covalent compounds: Use prefixes to indicate the number of each atom (mono-, di-, tri-, etc.), and the suffix -ide for the second element.

Example Table: Prefixes for Covalent Compounds

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

Properties of Ionic and Covalent Compounds

Physical State, Melting and Boiling Points

Ionic compounds are usually solid and crystalline at room temperature, with high melting and boiling points due to strong electrostatic forces.
Covalent compounds can be solids, liquids, or gases, and generally have lower melting and boiling points.

Electrical Conductivity

Ionic compounds conduct electricity when dissolved in water (electrolytes).
Covalent compounds usually do not conduct electricity (nonelectrolytes).

Comparison Table: Ionic vs. Covalent Compounds

Property	Ionic	Covalent
Composition	Metal + nonmetal	2 nonmetals
Electron behavior	Transferred	Shared
Physical state	Solid, crystalline	Any; crystal or amorphous
Conductivity	Electrolytes	Nonelectrolytes
Melting/Boiling point	High	Low

Drawing Lewis Structures

Guidelines for Lewis Structures

Write the skeletal structure, placing the least electronegative atom in the center (except hydrogen, which is always terminal).
Determine the total number of valence electrons (add for anions, subtract for cations).
Connect atoms with single bonds, then complete octets for terminal atoms, and finally for the central atom.
If the central atom lacks an octet, form double or triple bonds as needed.
Check that all atoms have complete octets (or duets for hydrogen) and that the total number of electrons is correct.

Example: Lewis Structure of CO2

Carbon is the central atom, oxygen atoms are terminal.
Total valence electrons:
Structure: O=C=O (each bond is a double bond)

Exceptions to the Octet Rule

Types of Exceptions

Incomplete octet: Some atoms (e.g., Be, B) can have fewer than eight electrons.
Odd electron species: Molecules with an odd number of electrons (e.g., NO) cannot have all atoms with octets.
Expanded octet: Atoms in period 3 or below (e.g., P, S) can have more than eight electrons.

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the shape of molecules based on the repulsion between electron pairs around a central atom. Electron pairs arrange themselves to minimize repulsion, determining the molecular geometry.

2 electron groups: linear (180°)
3 electron groups: trigonal planar (120°)
4 electron groups: tetrahedral (109.5°)
Lone pairs distort bond angles, leading to shapes like trigonal pyramidal (e.g., NH3, 107°) and bent (e.g., H2O, 104.5°).

Molecular Polarity

Determining Polarity

A molecule is polar if it has polar bonds and an asymmetric shape, resulting in a net dipole moment.
Molecules with no lone pairs on the central atom and identical terminal atoms are usually nonpolar.
Molecules with lone pairs on the central atom are usually polar.

Example: H2O is polar; CO2 is nonpolar.

Intermolecular Forces and Physical Properties

Types of Forces

Intramolecular forces: Forces within molecules (chemical bonds).
Intermolecular forces: Forces between molecules, affecting physical properties like melting and boiling points.
Hydrogen bonding is a strong type of intermolecular force, especially in molecules like water and ammonia.

Solubility

"Like dissolves like": Polar molecules dissolve in polar solvents; nonpolar molecules dissolve in nonpolar solvents.
Example: Ammonia (NH3) dissolves in water due to hydrogen bonding; oil does not mix with water because it is nonpolar.

Melting and Boiling Points

Stronger intermolecular forces lead to higher melting and boiling points.
Larger molecular mass also increases melting and boiling points.
Polar molecules have higher melting and boiling points than nonpolar molecules of similar mass.

Example Table: Melting and Boiling Points by Bonding Type

Compound	Bonding Type	Melting Point (°C)	Boiling Point (°C)
Sodium chloride (NaCl)	Ionic	801	1413
Potassium bromide (KBr)	Ionic	730	1435
Water (H2O)	Covalent (polar)	0	100
Carbon dioxide (CO2)	Covalent (nonpolar)	-78	-57
Methane (CH4)	Covalent (nonpolar)	-182	-164