Table of Contents
Position of this Chapter within the Course
This chapter gives a compact overview of the main types of chemical bonds that hold atoms together in substances. Later chapters will treat each type (covalent, ionic, metallic) and special intermolecular interactions (van der Waals forces, hydrogen bonding) in detail. Here, the goal is to:
- Name and distinguish the principal bonding types at a basic level.
- Connect types of bonds with the kinds of substances in which they occur.
- Indicate how bonding type relates qualitatively to properties such as hardness, melting point, and electrical conductivity, without going into the deeper physical explanations that belong to later chapters.
Why Different Types of Bonds Exist
Atoms interact because energetically favorable arrangements can be reached by:
- Sharing electrons between atoms.
- Transferring electrons from one atom to another.
- Delocalizing electrons over many atoms.
Which of these patterns dominates depends mainly on:
- The nature of the atoms involved (especially whether they are metals or nonmetals).
- The difference in electronegativity between the atoms.
- The number of valence electrons and how filled the outer shells are.
From these patterns emerge three main bond types within solids and molecules:
- Covalent bonds
- Ionic bonds
- Metallic bonds
and, in addition, weaker interactions between molecules (intermolecular interactions), which are treated in a separate part of the course.
The Three Main Types of Chemical Bonds (Internal Bonds)
Covalent Bonding – Localized Electron Sharing
Typical situation:
Nonmetal + nonmetal (e.g. H and O, C and H, N and O).
Core idea:
Two (or more) atoms share one or more pairs of electrons so that each atom attains a more stable electron configuration (often reminiscent of noble gases).
Key features (only at overview level)
- Electrons are localized between specific pairs of atoms.
- Can form:
- Discrete molecules (e.g. $ \mathrm{H_2O} $, $ \mathrm{CO_2} $, $ \mathrm{CH_4} $).
- Giant covalent networks (e.g. diamond, quartz), where each atom is covalently bonded to many neighbors in a 3D network.
- Bonding is often directional: the orientation of bonds in space is important for structure and properties.
Typical qualitative properties
- Molecules with covalent bonds:
- Often low to moderate melting and boiling points (many exist as gases or liquids at room temperature).
- Usually do not conduct electricity in pure form (no freely moving charged particles).
- Covalent network solids:
- Often very hard and have very high melting points (e.g. diamond).
- Typically insulators (again, due to lack of mobile charges).
(The detailed origin of these properties, including orbital overlap and molecular geometry, is covered elsewhere.)
Ionic Bonding – Electrostatic Attraction between Ions
Typical situation:
Metal + nonmetal (e.g. Na and Cl, Ca and O).
Core idea:
- Electrons are transferred from one atom (usually a metal) to another (usually a nonmetal).
- This forms positively charged cations (e.g. $ \mathrm{Na^+} $) and negatively charged anions (e.g. $ \mathrm{Cl^-} $).
- Oppositely charged ions attract each other via electrostatic forces, forming an extended ionic lattice.
Key features (overview)
- Bonding arises mainly from Coulomb attraction between ions:
$$ F \propto \frac{q_1 q_2}{r^2} $$ - The structure is typically a regular crystal lattice of alternating cations and anions, not discrete molecules.
- Bond is non-directional: the attraction is spherically around each ion, depending only on distance and charge.
Typical qualitative properties
- Often high melting and boiling points (strong electrostatic forces).
- Solid ionic compounds:
- Generally hard and brittle.
- Usually do not conduct electricity (ions are fixed in the crystal lattice).
- Molten ionic compounds and aqueous solutions of ionic compounds:
- Often good electrical conductors (ions can move).
(The detailed treatment of lattice energy, solubility, and conductivity will appear in other chapters.)
Metallic Bonding – Delocalized Electrons in a Metal Lattice
Typical situation:
Metal + metal, or a pure elemental metal (e.g. Na, Cu, Fe, alloys like brass).
Core idea:
- Metal atoms in a solid contribute their valence electrons to a kind of “electron sea” that is delocalized throughout the entire crystal.
- The solid can be viewed (conceptually) as an array of positive metal ions immersed in a shared cloud of mobile electrons.
Key features (overview)
- Electrons are not localized between two atoms but can move relatively freely through the lattice.
- This leads to a non-directional bond connecting each metal atom to many neighbors via the collective electron cloud.
- Metallic bonding naturally supports formation of alloys, where different metal atoms share the same electron sea.
Typical qualitative properties
- Usually good electrical and thermal conductivity (due to mobile electrons).
- Typically malleable and ductile: metals can be hammered into sheets or drawn into wires without shattering (layers of atoms can slide while metallic bonding persists).
- Often moderate to high melting points (varies widely among metals).
(The modern band theory explanation and details of metallic structure are covered in later chapters.)
Intermolecular vs. Intramolecular Interactions
The bond types above (covalent, ionic, metallic) are mainly about intramolecular or intra‑crystalline forces:
- They hold atoms together within molecules or within extended solids.
In addition, there are intermolecular interactions that act between molecules or particles, such as:
- van der Waals forces
- Hydrogen bonds
These interactions are generally weaker than the main bond types, but they critically influence:
- Melting and boiling points,
- Solubility,
- Viscosity,
- Structures of biological macromolecules (e.g. DNA, proteins).
They are treated separately in the chapter on Special Intermolecular Interactions.
Comparing Bond Types at a Glance
The following table summarizes the main bonding categories discussed in this part of the course. (Details and exceptions are handled in dedicated chapters.)
| Aspect | Covalent Bonding | Ionic Bonding | Metallic Bonding |
|---|---|---|---|
| Typical elements involved | Nonmetal–nonmetal | Metal–nonmetal | Metal–metal (or pure metal) |
| Electron behavior | Shared electron pairs, localized | Electrons transferred; ions formed | Valence electrons delocalized over whole lattice |
| Representative structures | Molecules or covalent networks | Extended ionic lattices | Crystalline metal lattices |
| Directionality | Directional (specific bond angles) | Mostly non-directional electrostatics | Non-directional (overall electron cloud) |
| Usual electrical behavior | Mostly insulators; some exceptions | Insulators as solids; conductive when molten/in solution | Good conductors in solid and liquid states |
| Typical mechanical behavior | Molecules: soft; networks: very hard | Hard but brittle | Malleable, ductile |
| Typical melting/boiling point | Low–high (depends strongly on structure) | Generally high | Moderate to high (wide range) |
This overview provides the conceptual map for the more detailed discussions of covalent, ionic, and metallic bonding, as well as the separate treatment of intermolecular forces, that follow in subsequent chapters.