Chemical Bonding

Why Atoms Form Chemical Bonds

Atoms do not “like” to exist on their own if they can reach a lower-energy, more stable state by joining with other atoms. Chemical bonding is the general term for all the ways in which atoms (or ions) are held together in aggregates such as molecules, crystals, or extended networks.

A very simple way to think about bonding is:

• Unbonded atoms: relatively high potential energy, often reactive
• Bonded atoms: lower potential energy, often more stable

The drive toward lower energy and greater stability is the underlying reason chemical bonds form.

A key idea that will be used throughout this topic (without yet going into the details you’ll see later) is:

• Atoms tend to form bonds so that their outer electron arrangement resembles that of a noble gas (a particularly stable configuration).

Exactly how this is achieved is different for different bond types and will be discussed in the separate chapters on covalent, ionic, metallic bonds, and intermolecular interactions.

Types of Chemical Bonding – The Big Picture

In this course, you will encounter two broad classes of interactions:

1. Intramolecular bonds – forces that hold atoms together within a single chemical unit (molecule, ion, or extended solid)
2. Intermolecular interactions – forces between separate molecules or ions

Separate chapters will deal with each main type, but it is useful to see how they fit together conceptually.

Intramolecular Bonds (Within Particles)

These define the basic connectivity of a chemical substance.

• Covalent bonding
  Atoms share one or more pairs of electrons.
  Typical in molecules like $ \mathrm{H_2O} $, $ \mathrm{CO_2} $, $ \mathrm{O_2} $.
• Ionic bonding
  Electrons are transferred from one atom to another, producing oppositely charged ions that attract one another.
  Typical in salts like $ \mathrm{NaCl} $ or $ \mathrm{CaCl_2} $.
• Metallic bonding
  Many metal atoms share a “sea” of delocalized electrons, giving metals their typical properties (conductivity, malleability).

These three will be compared in the “Main Types of Chemical Bonds” section in detail; here you only need to recognize that they are fundamental frameworks that hold matter together.

Intermolecular Interactions (Between Particles)

These do not generally change which atoms are connected to which; instead they influence how particles attract each other at close range.

• van der Waals forces (London dispersion, dipole–dipole, etc.)
• Hydrogen bonding

These interactions are usually weaker than covalent, ionic, or metallic bonds, but they strongly affect boiling points, melting points, solubilities, and the physical state (solid, liquid, gas). They will be handled as “Special Intermolecular Interactions” in a separate section.

Bonding and Energy

The formation and breaking of bonds are central to understanding chemical reactions.

• Bond formation releases energy to the surroundings (exothermic at the bond level) because the system moves to a lower-energy state.
• Bond breaking requires energy (endothermic at the bond level) to separate atoms that attract each other.

On an energy diagram:

• The bonded state corresponds to a minimum in potential energy as a function of the distance between two atoms.
• If atoms are too far apart, they mostly do not feel each other.
• If they are too close, positively charged nuclei repel strongly, raising energy.
• At an intermediate distance, attraction and repulsion balance at a minimum: that is the equilibrium bond distance.

The bond energy (or bond enthalpy, in a more thermodynamic sense) is the energy needed to break a bond in a molecule in the gas phase. It is a measure of how strong that bond is.

You will later connect bond energies to broader thermodynamic quantities (such as enthalpy changes) in physical chemistry chapters; here the important idea is just:

• Stronger bonds → larger bond energies → generally more stable arrangements.

Bonding and Structure

The type and arrangement of bonds determine the structure and therefore many properties of a substance.

Some key patterns (details follow in later chapters):

• Discrete molecules:
  Atoms linked by covalent bonds into independent units (e.g. $ \mathrm{H_2O} $, $ \mathrm{CH_4} $).
  These units interact via intermolecular forces.
• Ionic lattices:
  Ions held together in extended three-dimensional networks by ionic bonds (e.g. crystals of $ \mathrm{NaCl} $).
  No separate “molecules” of $ \mathrm{NaCl} $ exist in the solid.
• Metallic crystals:
  Metal atoms in regular arrays with delocalized electrons throughout the solid (e.g. copper metal, $ \mathrm{Cu} $).

Different bonding patterns lead to very different physical properties, for example:

• melting and boiling points
• hardness and brittleness
• electrical and thermal conductivity
• solubility in water or in other solvents

These property differences will be explained in more detail when individual bond types and intermolecular interactions are discussed.

Bond Polarity and Electronegativity (Conceptual Overview)

Even before learning each specific bond type, it is useful to have a basic idea of bond polarity.

Atoms differ in how strongly they attract shared electrons. This tendency is called electronegativity. When two atoms bond:

• If their electronegativities are similar, the shared electrons are roughly in between; the bond is nonpolar covalent.
• If one atom is more electronegative, it pulls the shared electrons closer; the bond becomes polar covalent (one end slightly negative, the other slightly positive).
• If the difference is very large, electrons are transferred almost completely, leading to ionic character.

Thus, covalent and ionic bonding are not completely separate categories, but rather the ends of a continuum of how electrons are distributed between atoms. Metallic bonding involves a different picture (delocalization over many atoms), but still reflects how electrons can be shared or moved in a system.

Full, quantitative treatment of electronegativity and bond polarity will appear in more specialized bonding chapters; at this stage, remember:

• Bonding involves electrons.
• How equally or unequally electrons are shared affects the charge distribution, which in turn influences both bond type and intermolecular interactions.

Bonding and Chemical Formulas

Chemical bonding concepts are closely tied to how we write and interpret chemical formulas:

• A molecular formula (e.g. $ \mathrm{C_2H_6O} $) tells you the number of each type of atom in a molecule, but not how they are connected.
• A structural formula shows which atoms are bonded to which.
• For ionic compounds, we write a formula unit (e.g. $ \mathrm{NaCl} $, $ \mathrm{CaCO_3} $) that reflects the simplest whole-number ratio of ions in the crystal lattice, not discrete molecules.
• For metals, we write the element symbol ($ \mathrm{Cu} $, $ \mathrm{Fe} $); the bonding pattern is an extended metallic structure.

Understanding bond types is therefore essential for:

• Interpreting formulas and structures
• Predicting how a substance might behave in reactions and in different physical conditions

Bonding and Macroscopic Properties – A Qualitative Map

Here is a qualitative map linking bond types with some typical bulk properties. The detailed “why” belongs in the later chapters; for now, this gives you a sense of how bonding connects to what you observe.

• Ionic compounds (ionic bonding)
• Often crystalline solids at room temperature
• High melting and boiling points
• Usually brittle
• Conduct electricity when melted or dissolved in water (ions can move), but not as solids
• Molecular substances (mostly covalent bonding within molecules, weaker interactions between them)
• Can be gases, liquids, or low-melting solids at room temperature
• Often poor conductors of electricity
• Properties depend strongly on intermolecular forces (e.g. hydrogen bonding → higher boiling point)
• Metals (metallic bonding)
• Usually solid (except mercury) with moderate to high melting points
• Good conductors of heat and electricity
• Malleable and ductile
• Network covalent solids (extended covalent bonding, like diamond or quartz)
• Very high melting points
• Often very hard
• Typically poor electrical conductors (with notable exceptions)

Later chapters will use specific examples to show how these patterns arise from bond type and arrangement.

How Bonding Fits into the Bigger Picture of Chemistry

Chemical bonding is central to almost every topic you will meet later:

• In atomic structure and the periodic table, you learn why different elements tend to form particular bonds.
• In stoichiometry, you use formulas (which reflect bonding) to perform quantitative calculations.
• In thermodynamics and kinetics, you analyze why bonds form or break and how fast this happens.
• In organic, inorganic, and coordination chemistry, you study how varied bonding patterns create the enormous diversity of compounds.
• In materials, biochemistry, and environmental chemistry, you see how bonding determines material properties, biological function, and environmental behavior.

For now, the key ideas for this introductory chapter on chemical bonding are:

• Bonds form because they lower the energy and increase the stability of atoms.
• There are several main bonding types (covalent, ionic, metallic) and a variety of weaker intermolecular interactions.
• The behavior and properties of substances are intimately tied to the kinds of bonds and interactions present.