Covalent bonds: two atoms sharing an electron pair

How to use this simulation

1. Start with “Water (H₂O)” selected: watch the oxygen and two hydrogen atoms drift together, then one shared pair settle into each O–H bond.
2. Keep “Show lone electron pairs” on: two extra pairs appear on oxygen — electrons that never joined a bond.
3. Switch to “Nitrogen gas (N₂)”: three pairs form one after another into a triple bond, and the readings panel names the bond type “Triple bond”.
4. Turn on “Show bond polarity and electronegativity difference” for H₂O: oxygen gets a δ− label, hydrogen a δ+, with an arrow showing the pull.
5. Keeping that toggle on, try “Methane (CH₄)” and “Carbon dioxide (CO₂)”: the bonds still show a pull, but the readings panel calls the molecule “Non-polar” — the next section explains why.

Every breath you take is a bonding story

The air filling your lungs right now is almost entirely nitrogen and oxygen gas — two atoms of the same element, paired up as N₂ and O₂. There is no metal anywhere in that air, so no atom is giving an electron away completely. So how do two atoms of the same element stick together at all?

The answer is inside you too. Water (H₂O), the glucose you eat, the protein your muscles are built from — the bonds holding all of them together work the same way. Neither atom hands an electron to the other; instead, a pair of electrons is used by both atoms at once. That sharing is the whole idea, and its name is the covalent bond.

In the ionic bond page you saw sodium give its one outer electron to chlorine completely, so both end up permanently charged (Na⁺, Cl⁻). Here the story is different: nobody loses anything and nobody becomes charged (when the two atoms are identical). Both atoms become joint owners of one electron pair. That contrast is the whole point of this page.

Starting from zero: what a covalent bond really is

When two or more atoms share one or more pairs of outer-shell electrons to hold themselves together, the bond formed is a covalent bond, and the compound is a covalent (molecular) compound. Every shared pair has exactly two electrons, one contributed by each atom.

Atoms bond for one reason: every atom "wants" the same outer-shell electron count as the nearest noble gas — eight, for every element in this page except hydrogen, which wants two. A full outer shell is the most stable arrangement there is. An atom that cannot reach that number alone pairs up with another, and the shared pair is counted toward both atoms' totals at once. This rule is called the octet rule.

Chlorine, for example, has 7 outer electrons and needs just one more. Two chlorine atoms sharing one pair means that pair counts for both of them, so both reach eight. That is exactly how Cl₂ forms, with no atom losing an electron and no new charge appearing anywhere.

Covalent bonds usually form between two non-metals — both have a strong pull on electrons (a high electronegativity), so neither is willing to give one away. Ionic bonds, by contrast, form between a metal and a non-metal, because the metal's electronegativity is low enough that it gives an electron up easily.

Key terms

Get the vocabulary straight before the diagrams — definition questions in exams come straight from this table.

Single, double and triple bonds: how many pairs are shared

The name of a bond comes straight from how many pairs the two atoms share — this count is the bond order. Hydrogen, chlorine, oxygen and nitrogen gas give four clean examples of four different bond orders, all in molecules made of identical atoms.

The rules for drawing a dot-and-cross diagram

In a dot-and-cross (Lewis) diagram, one atom's electrons are drawn as dots (•) and the other's as crosses (×) — even for two atoms of the same element, so a reader can tell whose electron is whose. The simulation draws the central atom's electrons as dots and the outer atom's as crosses, exactly as a textbook diagram would.

Drawing one is a three-step process. First, write down each atom's total valence electrons. Then place as many shared pairs (one dot plus one cross) between the two atoms as it takes to complete the octet (or duet, for hydrogen) — those are the bonding pairs. Finally, draw whatever electrons are left over in pairs beside their own atom — those are the lone pairs.

Check it against water. Oxygen brings 6 electrons, and each of the two hydrogens brings 1, for a total of 8. The two O–H bonds use 2 shared pairs, leaving oxygen with 2 lone pairs. Counting every electron — shared pairs × 2 plus lone pairs × 2 — gives 8, exactly matching the 8 you started with.

Bond polarity vs molecule polarity

When the two bonded atoms are identical (H₂, Cl₂, N₂), they pull on the shared pair equally hard — the bond is non-polar. When they are different, the more electronegative atom pulls the pair slightly towards itself, leaving it with a partial negative charge (δ−) and the other atom with a partial positive charge (δ+). That is a polar covalent bond.

In H₂O, oxygen's electronegativity is 3.44 against hydrogen's 2.20, a difference of ΔEN = 1.24 — large enough to make the O–H bond strongly polar. In NH₃, nitrogen's electronegativity is 3.04, so the N–H bond's ΔEN is only 0.84, noticeably less polar than O–H.

But a polar bond does not automatically make a polar molecule — this is where the biggest misconception in this topic breaks down. In CH₄, the C–H bond's ΔEN is 0.35, a small but real polarity. Yet methane's four bonds point to the four corners of a perfectly symmetric tetrahedron, so their pulls cancel exactly, and the molecule as a whole is non-polar. CO₂ shows the same trick from a different shape: each C=O bond has a sizeable ΔEN of 0.89, but the molecule is a straight line, so the two equal, opposite pulls cancel and CO₂ is non-polar too.

Water is different. Oxygen's two lone pairs bend the molecule into an angle, so the two O–H pulls cannot cancel — they add up instead. That is why H₂O is a genuinely polar molecule. The rule to keep: whether a molecule is polar depends on both the bond's ΔEN and the molecule's shape, never on ΔEN alone.

Try this in the simulation

Guess the answer before you run each one, then check the readings panel.

• Run H₂ and time how long one shared pair takes to form, then run N₂ and check that three pairs take roughly three times as long.
• With lone pairs on, compare NH₃ and CH₄ side by side: nitrogen keeps one extra pair that carbon does not — that single pair is the whole reason their shapes differ.
• Turn polarity on and compare H₂O and CH₄: both show δ+/δ− labels, but the readings panel calls H₂O a "Polar molecule" and CH₄ a "Non-polar molecule".
• Select CO₂ and read the bond type: "Double bond", bonds formed = 2 — that is two separate C=O double bonds, each with two shared pairs.
• Drop the speed to 0.25× on any molecule and watch a single pair travel from each atom's edge to the midpoint — that is what "sharing" actually looks like, frame by frame.

Worked problems

Every answer starts by counting valence electrons, then checks the octet (or duet) to find the shared and lone pairs.

Common mistakes

Clear these up and covalent-bond questions stop costing marks on conceptual and drawing questions alike.

• Thinking electrons are "transferred" in a covalent bond too. In reality both atoms keep joint ownership of the shared pair; nobody loses anything.
• Assuming a polar bond always makes a polar molecule. CH₄ and CO₂ prove otherwise — a symmetric shape cancels the pull.
• Applying the octet rule to hydrogen. Hydrogen follows the duet rule (2 electrons), because it has only one shell, which holds at most two.
• Confusing bond order with the number of bonds. A double bond has two shared pairs, but it is still one single connection between the two atoms.
• Forgetting lone pairs in a diagram. Once the bonding electrons are drawn, whatever is left must still be shown in pairs.
• Assuming ΔEN = 0 means "no bond". H₂ and Cl₂ bond perfectly well — the bond is simply non-polar, not absent.

Covalent bonds in real life

Covalent bonding is not just an exam topic — water, air, food and your own body are all held together by exactly this kind of bond.

• Water (H₂O): its polar bonds let molecules pull on each other (hydrogen bonding), which is why water has such an unusually high boiling point for such a small molecule.
• Nitrogen and oxygen in the air: N₂'s triple bond is so strong that it barely reacts, which is why the air keeps the same composition for millennia.
• Methane (CH₄): the main component of natural gas, a perfectly symmetric molecule of four single C–H bonds.
• Carbon dioxide (CO₂): the raw material plants use in photosynthesis, a linear molecule with two C=O double bonds.
• Plastics and the molecules of life: carbon's ability to form several covalent bonds at once is exactly what makes long chains — plastics, sugars, proteins, DNA — possible.
• Ammonia (NH₃): the key raw material for fertiliser; its lone pair on nitrogen makes it chemically very reactive.

Exam tips

Covalent bonding shows up in every syllabus that covers chemical bonding at all, from Grade 9 through Grade 11 chemistry. Definition questions ask for the bond itself and the octet/duet rule; short-answer questions ask you to draw a dot-and-cross diagram or to explain why a particular molecule is polar or non-polar; longer questions ask you to count shared and lone pairs for a given formula.

Revision: the last-minute summary

The night before an exam, this list plus the polarity table above is all you need to revisit.

• Covalent bond: two atoms sharing a pair of electrons, usually between two non-metals.
• Octet rule: every atom aims for eight outer electrons; hydrogen is the exception, aiming for two (the duet rule).
• Single bond = 1 shared pair, double bond = 2, triple bond = 3.
• Dot-and-cross diagram: one atom's electrons as dots, the other's as crosses; shared pairs in the middle, lone pairs beside their own atom.
• ΔEN = 0 gives a non-polar bond (H₂, Cl₂, N₂); ΔEN > 0 gives a polar bond, with δ− on the more electronegative atom.
• A polar bond does not guarantee a polar molecule — a symmetric shape (CH₄, CO₂) cancels the pull; a bent or asymmetric shape (H₂O, NH₃) does not.

Frequently asked questions