Atoms, Bonds and Molecular Shapes (VSEPR) Explained

What is inside an atom?

Every atom has a tiny, dense nucleus made of protons (positive) and neutrons (no charge). Around it is a cloud of electrons (negative). The nucleus holds almost all of the mass, yet it takes up a minuscule fraction of the atom's volume. Most of an atom is empty space.

Three numbers organise everything:

• Atomic number: the count of protons. This alone decides which element you have. Six protons is carbon, always.
• Electron count: equal to the protons in a neutral atom.
• Mass number: protons plus neutrons.

Change the protons and you get a different element. Change the neutrons or electrons and you get something subtler, which the next two sections cover.

Isotopes and ions: same element, small changes

Isotopes change the neutrons

Atoms of one element can carry different numbers of neutrons. These versions are called isotopes. Carbon-12 and carbon-14 both have six protons, so both are carbon, but carbon-14 has two extra neutrons. Chemically they behave almost identically, because chemistry is run by electrons. Their masses differ, and some isotopes are unstable and radioactive.

To find the neutrons, subtract: neutrons = mass number − atomic number.

Ions change the electrons

If an atom gains or loses electrons, the protons stay put but the overall charge changes. Lose electrons and you get a positive cation. Gain electrons and you get a negative anion. An oxygen nucleus with 8 protons and 10 electrons carries a net charge of 2−, and it is still oxygen.

A handy check: charge = protons − electrons. Say it out loud each time you meet an ion and the sign rarely goes wrong.

Shells and valence electrons

Electrons are arranged in shells around the nucleus. In the simple Bohr picture used for early chemistry, the first shell holds up to 2 electrons and the second up to 8. Lithium has 3 electrons, so its pattern is 2, 1. Neon has 10, so it is 2, 8, a full outer shell.

The outermost electrons are the valence electrons, and they decide how an atom bonds. The periodic table is organised so that elements in one column have the same number of valence electrons and therefore similar behaviour. Atoms with nearly full or nearly empty outer shells are the most eager to react, because a small change gives them the stable pattern of a noble gas.

Ionic bonds: giving and taking

When a metal meets a non-metal, an electron is often handed over completely. Sodium has one valence electron to spare and chlorine is one short of a full shell. After the transfer you have Na⁺ and Cl⁻, and opposite charges attract strongly. That attraction is the ionic bond.

Ionic compounds like table salt do not form separate molecules. Instead, ions pack into a repeating lattice, which is why they are typically hard, brittle solids with high melting points. If this packing idea interests you, the Crystal Lab shows how repeating cells build solids.

Covalent bonds: sharing pairs

Between two non-metals, neither atom can easily win the tug-of-war for electrons, so they share a pair. One shared pair is a single bond, two make a double bond, three make a triple bond. More shared pairs mean a shorter, stronger link.

Each hydrogen in water shares one pair with oxygen. In carbon dioxide, carbon shares two pairs with each oxygen. Counting bonds is a quick way to check a formula: hydrogen usually forms one bond, oxygen two, nitrogen three and carbon four.

Sharing is not always fair. When one atom pulls harder, the bond is polar, with a slightly negative end and a slightly positive end. Whether a whole molecule is polar then depends on its shape, which brings us to VSEPR.

VSEPR: electron groups push each other apart

VSEPR stands for valence shell electron pair repulsion, and the rule fits in one sentence: regions of electrons around a central atom repel each other, so they spread as far apart as possible.

Count the electron domains on the central atom. A single, double or triple bond to one neighbour counts as one domain, and so does each lone pair.

• 2 domains: linear, 180°
• 3 domains: trigonal planar, 120°
• 4 domains: tetrahedral, about 109.5°

Lone pairs are domains too, but you name the molecule's shape only from where the atoms sit. Carbon dioxide has two domains and no lone pairs, so it is linear. Ammonia has four domains with one lone pair, so the atoms form a trigonal pyramid. Water has four domains with two lone pairs, leaving a bent shape. Lone pairs take up more room than bonding pairs, so the real angles are a little squeezed compared with the ideal numbers.

Larger central atoms can hold five or six domains, giving trigonal bipyramidal and octahedral arrangements.

Common mistakes to avoid

• Treating a double bond as two domains. For VSEPR it is one region of electrons.
• Forgetting lone pairs when counting. Always draw them before predicting a shape.
• Mixing up electron-domain geometry and molecular shape. Ammonia's domains form a tetrahedron, but its atoms form a pyramid.
• Thinking isotopes are different elements. Same protons, same element.
• Assuming electrons orbit like planets. Real electron behaviour is described by probability clouds, a theme in the Quantum Lab.

How to explore this in the lab

The Chemistry Lab simulation is a simplified teaching model, but it lets you build and rotate real 3D molecules. A good order for the free lessons:

1. Inside the atom to fix the vocabulary.
2. Same element, different isotope and Electrons and ions to see what each particle changes.
3. Ionic bonds and Covalent bonds.
4. Electron domains set the shape to predict geometry before you build it.

Try predicting each shape on paper first, then check it in 3D. Being wrong once and seeing why sticks better than reading the answer.