From Atomic Structure to Bonding: Connecting the Ideas

Chemistry courses often teach atomic structure in one unit and bonding in the next, as if they were separate subjects. They aren’t. Every bond is a consequence of how electrons are arranged in the atoms that form it. Once you see the connection, bonding stops being a list of rules and becomes something you can predict.

This article walks through the chain of reasoning, from the nucleus outwards.

Link 1: the nucleus sets the pull

Every atom has a positively charged nucleus holding negatively charged electrons. How strongly an outer electron is held depends on two things:

• Effective nuclear charge: the nuclear charge an outer electron actually “feels” after inner electrons shield part of it (see effective nuclear charge).
• Distance: outer electrons in higher shells sit further away and feel a weaker pull.

These two factors drive the key periodic trends:

• Atomic radius shrinks across a period and grows down a group.
• Ionization energy rises across a period and falls down a group.
• Electronegativity, an atom’s pull on shared electrons in a bond, rises across a period and falls down a group.

Everything that follows about bonding comes back to these trends.

Link 2: the electron configuration decides what an atom “has to offer”

The electron configuration tells you how many valence electrons an atom has and which orbitals they occupy.

• An atom with one or two valence electrons in an s orbital, far from the nucleus (like sodium, 3s¹), holds them loosely. It can give them away cheaply.
• An atom one or two electrons short of a full p subshell (like chlorine, 3s² 3p⁵), close to the nucleus, holds its electrons tightly and attracts more.
• An atom with a half-filled set of valence orbitals (like carbon, with four valence electrons) is neither a good giver nor a good taker. It tends to share.

So the configuration already hints at the kind of bond an atom will form. The valence electrons and reactivity article looks at this in more detail.

Link 3: electronegativity difference sorts bonds into types

When two atoms meet, what matters is how their pulls on electrons compare. A useful rough guide uses the difference in Pauling electronegativity, ΔEN:

ΔEN (approx.)	Bond type	What happens to the electrons
Less than 0.4	Non-polar covalent	Shared almost equally
0.4 to about 1.8	Polar covalent	Shared unequally
Greater than about 1.8	Mainly ionic	Effectively transferred

Some examples, using values from this site’s element data:

• Cl–Cl: ΔEN = 0 → non-polar covalent.
• C–H (2.55 − 2.20): ΔEN = 0.35 → essentially non-polar.
• H–Cl (3.16 − 2.20): ΔEN = 0.96 → polar covalent.
• O–H (3.44 − 2.20): ΔEN = 1.24 → strongly polar covalent.
• Na–Cl (3.16 − 0.93): ΔEN = 2.23 → ionic.

The boundaries are not sharp. Bonding is a continuum, and every “ionic” bond has a little covalent character, and vice versa. See ionic vs covalent bonds for the full comparison.

Link 4: three kinds of bonding from one set of ideas

Ionic bonding: transfer, then attraction

When a low-ionization-energy metal meets a high-electronegativity non-metal, electrons move from metal to non-metal. Sodium becomes Na⁺ and chlorine becomes Cl⁻, each with a noble-gas configuration. The oppositely charged ions attract in every direction, so they pack into a giant three-dimensional lattice, not separate molecules.

Atomic structure predicts the charges: group 1 metals form 1+ ions, group 2 form 2+, group 16 non-metals form 2−, group 17 form 1− (see predicting ion charges). Ion size, again from atomic structure, affects how tightly the lattice holds together: small, highly charged ions give very strong lattices and high melting points. That’s why magnesium oxide (Mg²⁺ and O²⁻) melts at over 2,800 °C, while sodium chloride melts at 801 °C.

Covalent bonding: sharing through overlapping orbitals

When two atoms with similar, fairly high electronegativities meet, neither can take an electron from the other. Instead, a half-filled orbital on each atom overlaps, and the pair of electrons in the overlap region is attracted to both nuclei at once. That shared pair is the covalent bond.

Atomic structure predicts how many bonds form: an atom typically forms as many covalent bonds as it has unpaired electrons available (or can make available). Carbon forms four, nitrogen three, oxygen two, and hydrogen and the halogens one. This is exactly what Lewis symbols show with their unpaired dots.

It also predicts shape. Lone pairs and bonding pairs around a central atom repel each other, giving the bent shape of water and the pyramidal shape of ammonia (VSEPR theory). Combine shape with bond polarity and you can tell whether a whole molecule is polar (polar vs non-polar molecules).

Metallic bonding: electrons shared by everyone

In a metal, every atom has a few loosely held valence electrons, and there’s no electronegative partner to take them. Instead, the valence electrons spread out through the whole solid as a “sea” of delocalised electrons, holding a lattice of positive metal ions together.

Atomic structure explains metallic properties directly:

• Electrical conductivity: delocalised electrons move freely when a voltage is applied.
• Malleability: layers of ions can slide without breaking the bonding, because the electron sea moves with them.
• Melting point trends: more valence electrons per atom and smaller ions give stronger metallic bonding, so magnesium (two valence electrons) melts at 650 °C, much higher than sodium (one) at about 98 °C.

More in metallic bonding.

Link 5: from bonds to intermolecular forces

Covalent molecules are held together internally by strong bonds, but separate molecules attract each other only through weaker intermolecular forces. Atomic structure matters here too:

• London dispersion forces grow with the number of electrons, so bigger atoms and molecules attract more strongly. That’s why the halogens go from gases (F₂, Cl₂) to a liquid (Br₂) to a solid (I₂) down the group.
• Dipole–dipole forces arise from polar bonds in polar molecules.
• Hydrogen bonds form when hydrogen is bonded to a small, very electronegative atom (N, O or F), and they explain water’s surprisingly high boiling point (hydrogen bonding).

See intermolecular forces for the full picture.

Putting it all together: a worked chain

Why is table salt a hard, high-melting solid that conducts electricity when molten, while chlorine is a gas?

1. Atomic structure: Na has one loosely held 3s electron; Cl has a nearly full 3p subshell and high electronegativity.
2. Bond type: ΔEN = 2.23, so Na gives its electron to Cl → ionic bonding.
3. Structure: Na⁺ and Cl⁻ pack into a giant lattice with strong attractions in all directions.
4. Properties: lots of energy is needed to break the lattice → high melting point (801 °C). In the melt, ions move freely → it conducts.

Meanwhile, in Cl₂, two identical atoms share a pair of electrons (ΔEN = 0), forming small non-polar molecules held to each other only by weak dispersion forces. They separate easily, so chlorine is a gas at room temperature.

The same nucleus-to-properties chain works for almost any substance.

Common misconceptions

• “Ionic and covalent are completely different.” They are two ends of a continuum set by electronegativity difference.
• “Atoms bond to get a full shell because they want to.” Bonds form because the bonded arrangement has lower energy than separate atoms.
• “Metals and non-metals always form ionic bonds.” Many don’t: aluminium chloride, for example, has considerable covalent character.
• “Intermolecular forces are bonds inside molecules.” They act between molecules and are much weaker than covalent bonds.

Key takeaways

• Nuclear charge, shielding and distance set atomic size, ionization energy and electronegativity.
• The electron configuration shows how many valence electrons an atom has and how tightly they’re held.
• The electronegativity difference between atoms decides whether electrons are shared equally, shared unequally or transferred.
• Ionic, covalent and metallic bonding all follow from these atomic properties, and bonding in turn explains melting points, conductivity and states of matter.
• To see the underlying trends at a glance, open the electronegativity heatmap.