Chemical Bonding

Definition: Chemical bonding is the attraction that holds atoms together in molecules or extended solids, primarily through ionic bonds (electron transfer), covalent bonds (electron sharing), or metallic bonds (delocalized electron sharing among many atoms).

How It Works

Atoms bond because doing so lowers the total energy of the system, usually by achieving a more stable electron arrangement, often a filled outer shell, the basis of the octet rule.

Ionic bonds form when one atom transfers one or more electrons to another, typically a metal losing electrons to a nonmetal.
The result is oppositely charged ions, cation and anion, held together by electrostatic attraction.
These form an extended crystal lattice, not a single discrete “molecule.”
Covalent bonds form when two atoms share one or more electron pairs, typically between nonmetals.
Sharing lets both atoms count the shared electrons toward their own stable configuration.
Bonds can be single, double, or triple depending on how many electron pairs are shared.
Bonds can be polar (unequal sharing) or nonpolar (equal sharing) depending on electronegativity difference.
Metallic bonds occur in solid metals, where valence electrons are delocalized across the entire lattice of metal cations, a “sea of electrons.”
This delocalization is why metals conduct electricity and heat well.
It’s also why metals can be bent or drawn into wires, malleable and ductile, without breaking, since the electron sea keeps holding the shifted cations together.
Whether a bond leans ionic or covalent is a matter of degree, not a hard switch.
It’s estimated from the electronegativity difference (ΔEN) between the two atoms on the Pauling scale.
Bonding theory also distinguishes sigma bonds from pi bonds.
Sigma bonds form by head-on orbital overlap along the bond axis, present in every single bond.
Pi bonds form by sideways overlap of parallel p orbitals, added on top of a sigma bond in double and triple bonds.
Pi bonds are generally weaker and more reactive than sigma bonds.
This is why double and triple bonds tend to be more chemically reactive sites than single bonds.

Under the Hood

Rough Pauling-scale guidelines for bond character:

ΔEN	Bond Character
0.0-0.4	Nonpolar covalent
0.4-1.7	Polar covalent
>1.7	Ionic

Worked example: predicting bond type.

Given: sodium (EN = 0.93) and chlorine (EN = 3.16)
Step 1: find ΔEN

ΔEN = 3.16 - 0.93 = 2.23

Answer: ionic, consistent with NaCl being a classic ionic solid
Given: hydrogen (EN = 2.20) and chlorine (EN = 3.16)

ΔEN = 3.16 - 2.20 = 0.96

Answer: polar covalent, consistent with HCl being a covalent molecule with a significant dipole
Given: carbon (EN = 2.55) and hydrogen (EN = 2.20)

ΔEN = 2.55 - 2.20 = 0.35

Answer: essentially nonpolar covalent, why C-H bonds are treated as nonpolar in most organic contexts

Lattice energy measures how strongly an ionic solid’s ions attract each other:

E ∝ (Q1 × Q2)/r

Q1, Q2 are ionic charges, r is the distance between ion centers.
MgO (charges +2/-2) has a far higher lattice energy and melting point (2852°C) than NaCl (charges +1/-1, 801°C).
Doubling both charges roughly quadruples the electrostatic attraction.
The smaller ionic radii of Mg²+ and O²- further shrink r and increase attraction even more.

Bond energy (energy needed to break one mole of a given bond) increases with bond order:

C-C single bond ≈ 347 kJ/mol
C=C double bond ≈ 614 kJ/mol
C≡C triple bond ≈ 839 kJ/mol

The increase isn’t linear: a triple bond is stronger than three single bonds’ worth.
But the additional pi bonds are individually weaker than the first sigma bond.

Bond order and length:

C-C single bond ≈ 154 pm
C=C double bond ≈ 134 pm
C≡C triple bond ≈ 120 pm

As bond order increases, bond length decreases and bond strength increases.
More shared electron pairs pull the two nuclei closer together and hold them more tightly.

Why It Matters

Bond type determines a substance’s macroscopic behavior.
Ionic solids tend to have high melting points, be brittle, and conduct electricity only when molten or dissolved.
Covalent molecular substances tend to have lower melting points and don’t conduct as solids.
Covalent network solids like diamond are extremely hard and high-melting, the entire crystal is one continuous bonded structure.
Metallic bonding explains why metals are used for wiring, conductivity, and structural applications, malleability under stress rather than fracture.

Common Pitfalls

Assuming “metal + nonmetal always means ionic.” Some metal-nonmetal compounds, like AlCl3, are significantly covalent.
Aluminum’s small size and higher charge polarize the bond in cases like this.
Confusing bond polarity, a property of one bond, with molecular polarity, a property of the whole molecule that depends on geometry too.
CO2 has polar C=O bonds but is a nonpolar molecule because the two bond dipoles cancel in its linear geometry.
Thinking a double bond is simply twice the strength of a single bond in a linear sense.
The pi bond that makes up part of a double or triple bond is weaker than the original sigma bond.
Treating the octet rule as universal; boron compounds like BF3 are stable with only 6 valence electrons.
Expanded octets occur for period 3+ elements, like sulfur in SF6.
Forgetting that ionic compounds don’t exist as discrete molecules.
NaCl is a formula unit representing a 1:1 ratio within an extended lattice, not a bonded pair floating alone.
Mixing up sigma and pi bonds when counting bond order; a single bond is always one sigma bond.
A double bond is one sigma plus one pi; a triple bond is one sigma plus two pi.

Comparison

Property	Ionic	Covalent	Metallic
Formed between	Metal + nonmetal	Nonmetal + nonmetal	Metal + metal
Electron behavior	Transferred	Shared (localized)	Delocalized (“sea”)
Conductivity (solid)	Poor	Poor	Good
Conductivity (molten/dissolved)	Good	Poor (molecular)	Good
Melting point	Generally high	Variable, often lower	Variable, often high
Malleability	Brittle	N/A (molecular) or hard (network)	Malleable, ductile
Structural unit	Extended crystal lattice	Discrete molecule or network	Extended lattice of cations

Real-World Application

Choosing wiring versus insulation for an electrical cable is a direct, practical application of bond type.

Copper wiring works because metallic bonding’s delocalized electron sea lets electrons flow freely under an applied voltage.
The rubber or PVC insulation wrapped around it is a covalent molecular material.
Its electrons are localized in individual bonds with no free carriers, making it an effective electrical insulator.
Choosing the wrong bond-type material for either role, a covalent polymer as the conductor or a metal as the insulator, would make the cable simply not function.

Example

Diamond and graphite are both pure carbon held together by covalent bonds, but arranged differently:

Diamond: each carbon bonds to four neighbors in a rigid 3D network, extremely hard, an electrical insulator.
Graphite: carbons bond to three neighbors in flat sheets with delocalized electrons between layers, soft, slippery, and conductive along the sheets.
Same bond type, radically different structure, radically different properties.

FAQ

Can a single molecule contain more than one type of bond?

Yes, commonly.
Sodium acetate (NaCH3COO) has an ionic bond between Na+ and the acetate ion.
It also has covalent bonds holding the acetate ion itself together.

Why do ionic compounds dissolve in water but not in oil?

Water’s polar molecules can surround and stabilize individual ions, hydration, overcoming the lattice energy holding the crystal together.
Nonpolar oil molecules have no comparable way to stabilize a charged ion.
Ionic solids generally don’t dissolve in nonpolar solvents.

Is there a bond type stronger than a covalent bond?

Within the same category, network covalent bonding, as in diamond or quartz, creates some of the hardest, highest-melting materials known.
Not because any single bond is fundamentally stronger.
But because breaking the solid means breaking many primary bonds simultaneously throughout the whole structure.

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