What Is a Chemical Bond? Why Atoms Stick Together

Almost nothing around you is made of single, separate atoms. Water is hydrogen and oxygen held together. Salt is sodium and chlorine locked in a crystal. A steel spoon is iron atoms bound in a lattice. Even the air is mostly pairs of atoms — N₂ and O₂. Only the noble gases, such as helium and argon, are happy on their own. So what holds atoms together, and why do they bother? This guide answers those questions from the ground up.

A simple definition

A chemical bond is a lasting attraction between atoms (or ions) that holds them together in a substance. In every case, the attraction is electrical: it comes from positively charged nuclei and negatively charged electrons pulling on each other.

That’s the key idea to hold on to. Bonds aren’t little sticks or hooks, even though we draw them as lines. They’re the result of electrons arranging themselves so that they’re attracted to more than one nucleus at once, or of oppositely charged ions attracting each other.

Why do atoms bond? The energy picture

Atoms bond because the bonded arrangement has lower energy than the separate atoms. Systems in nature tend to move towards lower energy, releasing the difference, usually as heat. A ball rolls downhill for the same reason.

Imagine two hydrogen atoms approaching each other:

1. Far apart: no interaction. We call this energy zero.
2. Getting closer: each electron starts to be attracted to both nuclei. The energy falls.
3. At the bond length (0.074 nm for H₂): attraction and repulsion balance. The energy is at its lowest — about 436 kJ lower per mole than for two separate atoms.
4. Too close: the two positive nuclei repel each other strongly, and so do the electrons. The energy shoots back up.

If you plot energy against the distance between the nuclei, you get a curve with a dip. The bottom of the dip is the stable bond. Its depth is the bond energy; its position is the bond length (see bond length and bond strength).

This explains two things at once:

• Forming a bond releases energy (exothermic).
• Breaking a bond needs energy (endothermic).

A very common mistake is to say “breaking bonds releases energy”. It never does. When fuels burn and release energy, it’s because the bonds formed in the products (CO₂ and H₂O) are stronger overall than the bonds broken in the reactants (see exothermic vs endothermic).

What about the “full outer shell” rule?

You’ve probably learned that atoms bond “to get a full outer shell” of eight electrons (the octet rule), like the noble gases. That’s a useful pattern for predicting formulas, especially for elements in periods 2 and 3. For example, sodium loses one electron and chlorine gains one; carbon shares four pairs of electrons in methane.

But atoms don’t “want” anything. The octet pattern works because of the energy: electron arrangements like those of noble gases happen to be particularly low in energy. And there are plenty of exceptions — boron in BF₃ has only six outer electrons, while sulfur in SF₆ has twelve around it. So think of the octet rule as a handy guide, and energy as the real reason.

The three main types of bond

Which kind of bond forms depends mainly on how strongly each atom attracts electrons — its electronegativity (see the electronegativity trend).

1. Ionic bonding: metal + non-metal

When a metal with a weak hold on its outer electrons meets a non-metal that attracts electrons strongly, electrons are transferred. The metal becomes a positive ion; the non-metal a negative ion.

Sodium (electronegativity 0.93) and chlorine (3.16) are the classic example:

Na → Na⁺ + e⁻ Cl + e⁻ → Cl⁻

The ions attract each other strongly in all directions, forming a giant ionic lattice — each Na⁺ surrounded by six Cl⁻ and each Cl⁻ by six Na⁺. That’s what a grain of salt is. Ionic compounds usually have high melting points and conduct electricity when molten or dissolved, because the ions can then move. More in ionic bonding explained.

2. Covalent bonding: non-metal + non-metal

When two atoms both attract electrons fairly strongly, neither can take an electron from the other. Instead, they share a pair of electrons. The shared pair sits between the nuclei and is attracted to both, holding them together.

In H₂, each hydrogen contributes one electron to a shared pair. In water, oxygen shares one pair with each hydrogen. Atoms can share two pairs (a double bond, as in O₂) or three (a triple bond, as in N₂). More in covalent bonding explained.

Covalent bonds can make:

• Simple molecules — small groups of atoms, like H₂O, CO₂ and CH₄. These usually have low melting and boiling points.
• Giant covalent structures — huge networks, like diamond, graphite and silicon dioxide (sand). These have very high melting points because melting means breaking many strong covalent bonds.

3. Metallic bonding: metal + metal

In a metal, atoms release some of their outer electrons into a shared “sea” of delocalised electrons that moves freely throughout the whole structure. The positive metal ions are held together by their attraction to this sea. This explains why metals conduct electricity and heat, and why they can be bent and hammered into shape without shattering (see metallic bonding).

Summary table

Bond type	Between	What happens to electrons	Typical properties
Ionic	Metal + non-metal	Transferred, forming ions	High melting point; conducts when molten or dissolved; brittle
Covalent (simple molecules)	Non-metals	Shared in pairs	Low melting point; usually doesn’t conduct
Covalent (giant)	Non-metals (e.g. C, Si)	Shared, in a network	Very high melting point; hard (except graphite)
Metallic	Metals	Delocalised in a “sea”	Conducts; malleable; shiny

It’s a spectrum, not three boxes

In reality, bonding isn’t always purely one type. When two different non-metals share electrons, the more electronegative atom pulls the pair closer, making a polar covalent bond with partial charges, δ+ and δ− (see bond polarity). As the electronegativity difference grows, bonds become more and more ionic in character. Hydrogen–hydrogen is purely covalent; hydrogen–chlorine is polar covalent; sodium–chlorine is mostly ionic. A rough guide: an electronegativity difference above about 1.7–2.0 usually indicates mainly ionic bonding (see ionic vs covalent bonds).

Bonds vs intermolecular forces

There’s one more distinction students often mix up. Bonds hold atoms together within a molecule or lattice. Intermolecular forces are much weaker attractions between separate molecules — such as hydrogen bonds between water molecules, or London forces between iodine molecules (see intermolecular forces).

When water boils, the O–H covalent bonds inside each molecule don’t break. The molecules simply separate, overcoming the weaker hydrogen bonds between them. Breaking the O–H bonds themselves would need far more energy — around 460 kJ per mole of bonds, compared with about 41 kJ per mole to boil water at 100 °C.

How strong are bonds?

Bond or attraction	Typical energy (kJ mol⁻¹)
Covalent bonds	about 150–950
Ionic lattice (per mole of compound)	about 700–4,000
Hydrogen bonds	about 10–40
London forces between small molecules	a few kJ mol⁻¹ or less, but increasing with size

The strongest known covalent bond between two atoms is the triple bond in carbon monoxide, at about 1,070 kJ mol⁻¹. The triple bond in N₂ (945 kJ mol⁻¹) is so strong that nitrogen gas is very unreactive — which is why making fertilisers from atmospheric nitrogen needs an industrial process with high temperature, high pressure and a catalyst.

Common misconceptions

• “Breaking bonds releases energy.” Breaking always needs energy; forming bonds releases it.
• “Atoms want a full outer shell.” Atoms have no wants; the arrangement is simply lower in energy.
• “Ionic compounds are made of molecules.” They form giant lattices; the formula NaCl gives the ratio of ions.
• “Boiling breaks covalent bonds.” For simple molecules, boiling overcomes intermolecular forces.
• “Bonds are physical sticks.” Lines in diagrams represent shared electron pairs or attractions, not solid links.

Key takeaways

• A chemical bond is an electrical attraction that holds atoms or ions together.
• Atoms bond because the bonded state has lower energy; forming bonds releases energy.
• Ionic: electron transfer between metals and non-metals; covalent: electron sharing between non-metals; metallic: delocalised electrons in metals.
• Bonding is a spectrum set by electronegativity differences.
• Bonds act within substances; intermolecular forces act between molecules and are much weaker.

To see the particles involved, start with atomic structure and bonding.