Structure and Properties of Ionic and Covalent Compounds

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Chemical Bonding

Definition and Nature of Chemical Bonds

Chemical bonding is fundamental to the structure and properties of matter. A chemical bond is the force of attraction between any two atoms in a compound. This attractive force overcomes the repulsion of the positively charged nuclei of the two atoms participating in the bond. Interactions involving valence electrons are responsible for the formation of chemical bonds.

Valence electrons are the outermost electrons of an atom and are primarily involved in bonding.
Only valence electrons participate in chemical 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 corresponds directly to the number of valence electrons in the 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: Attractive force due to the transfer of one or more electrons from one atom to another, resulting in oppositely charged ions.
Covalent bond: Attractive force due to the sharing of electrons between atoms.
Some bonds have characteristics of both types and are not easily classified as purely ionic or covalent.

Ionic Bonding

Formation and Properties

Representative elements form ions that obey the octet rule. Electrons are lost by a metal and gained by a nonmetal, resulting in the formation of cations and anions, which are attracted to each other by electrostatic forces.

Each atom achieves a noble gas configuration.
Oppositely charged ions attract, creating the ionic bond.

Example: Formation of NaCl

Sodium (Na) has a low ionization energy and readily loses its electron:
Chlorine (Cl) has a high electron affinity and gains an electron:
The resulting ions aggregate into a crystal lattice.

Covalent Bonding

Formation and Properties

Covalent bonds form between atoms with similar tendencies to gain or lose electrons. Compounds containing covalent bonds are called covalent compounds or molecules.

Diatomic elements (e.g., H2, O2, N2, F2, Cl2, Br2, I2) have completely covalent bonds (equal sharing).
Each shared pair of electrons constitutes a covalent bond.

Example: Formation of H2:

Polar Covalent Bonding and Electronegativity

Bond Polarity

Polar covalent bonds are formed when electron pairs are shared unequally between atoms due to differences in electronegativity.

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond.
The greater the difference in electronegativity between two atoms, the more polar the bond.

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

Naming Compounds and Writing Formulas

Nomenclature Systems

Ionic compounds: Name the cation (metal) first, followed by the anion (nonmetal) with the suffix -ide.
Covalent compounds: Use prefixes to indicate the number of each type of atom (mono-, di-, tri-, etc.). The last element ends with -ide.

Example: CO is carbon monoxide; N2O4 is dinitrogen tetroxide.

Properties of Ionic and Covalent Compounds

Physical State and Melting/Boiling Points

Ionic compounds are usually solids at room temperature and have high melting and boiling points.
Covalent compounds can be solids, liquids, or gases and generally have lower melting and boiling points.

Electrical Conductivity

Ionic compounds often dissolve in water to form electrolytes (conduct electricity).
Covalent compounds usually do not conduct electricity (nonelectrolytes).

Comparison Table: Ionic vs. Covalent Compounds

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

Drawing Lewis Structures

Guidelines

Write the skeletal structure with 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, complete octets for terminal atoms, then for the central atom.
If necessary, form double or triple bonds to satisfy the octet rule.

Example: CO2 has a linear structure with double bonds between C and each O.

Exceptions to the Octet Rule

Incomplete octet: Some atoms (e.g., Be, B) have fewer than eight electrons.
Odd electron species: Molecules with an odd number of electrons (e.g., NO).
Expanded octet: Elements in period 3 or below can have more than eight electrons (e.g., PF5).

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.

2 bonded atoms: linear (180°)
3 bonded atoms: trigonal planar (120°)
4 bonded atoms: tetrahedral (109.5°)
3 bonded atoms + 1 lone pair: trigonal pyramidal (107°)
2 bonded atoms + 2 lone pairs: bent (104.5°)

Molecular Polarity

Determining Polarity

Draw the Lewis structure and determine the geometry.
If the molecule has no lone pairs on the central atom and all terminal atoms are the same, it is nonpolar.
Molecules with lone pairs on the central atom are usually polar.

Example: H2O is polar due to its bent shape and lone pairs on oxygen.

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 important in 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 because both are polar.
Oil does not dissolve in water because oil is nonpolar and water is polar.

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.

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

Additional info: Some content, such as specific examples and table values, has been inferred or expanded for completeness and clarity.