How to Draw Lewis Structures Step by Step

Start by counting valence electrons for every atom in the formula. Main-group elements in the same periodic-table group often have the same number of valence electrons. Add the totals together. For an ion, add one electron for each negative charge or subtract one for each positive charge.

Next choose a central atom and connect surrounding atoms with single bonds. Hydrogen is never the central atom because it forms only one bond in ordinary Lewis structures. The least electronegative suitable atom is often central, although there are exceptions. Each single bond represents two shared electrons, so subtract those bonding electrons from your total.

Place the remaining electrons as lone pairs around the outer atoms first, usually aiming to give them full outer shells. Hydrogen needs only two electrons. Many common main-group atoms are most stable with eight electrons around them, which is the octet rule. This is the basic routine for drawing Lewis structures.

Sometimes you run out of electrons before the central atom reaches an octet. In that case, a lone pair from a neighbouring atom can sometimes be turned into an additional shared pair, creating a double or triple bond. Carbon dioxide is a familiar example: O=C=O gives carbon and both oxygens complete octets with the available electrons.

When drawing Lewis structures, never invent extra electrons just to make the picture look tidy. The total number of dots and bonding electrons must match the valence-electron count you found at the start. That count is one of your strongest error checks.

The octet rule is useful but not universal. Boron can appear in stable compounds with fewer than eight electrons, while elements in the third period and below can sometimes be shown with expanded valence shells in common Lewis models. Odd-electron species also exist. Learn the basic octet method first, then treat exceptions as a separate layer.

More than one Lewis structure may satisfy the basic electron count. Formal charge can help compare them. For a particular atom, formal charge is its valence-electron number minus its nonbonding electrons minus half of its bonding electrons. The formal charges across a neutral molecule should add to zero, while those in an ion should add to the ion's overall charge.

Structures with smaller formal charges are often preferred, and negative formal charge is generally more reasonable on more electronegative atoms. These are guidelines rather than a substitute for real bonding theory, but they make Lewis diagrams more useful for introductory chemistry.

A final routine for drawing Lewis structures is simple: count electrons, draw the skeleton, complete outer atoms, place leftovers on the centre, consider multiple bonds, and check formal charges. Work slowly enough that every dot has a reason to be there.

The carbonate ion, CO3 with a charge of minus two, shows several ideas together. Carbon brings four valence electrons, the three oxygens bring six each, and the charge adds two more, giving 24. Place carbon in the centre with single bonds to each oxygen, using six electrons. The remaining 18 go on the oxygens as three lone pairs each, completing their octets. Carbon now has only six electrons, so turn one lone pair from an oxygen into a second bond, giving one double bond and two single bonds.

Check the formal charges. The double-bonded oxygen has six valence electrons, four nonbonding and four bonding, so its formal charge is zero. Each single-bonded oxygen has six nonbonding and two bonding electrons, giving minus one. Carbon comes out at zero, so the total is minus two, matching the ion. But the double bond could be drawn to any of the three oxygens. These are called resonance structures, and measurements show the real ion has three identical carbon to oxygen bonds, each somewhere between a single and a double bond.

A Lewis structure is drawn flat on paper, but molecules are three-dimensional, and a common mistake is to assume the drawing shows the real shape. It does not, yet it contains the information you need to predict the shape. The idea is called valence shell electron pair repulsion, usually shortened to VSEPR. Each bond and each lone pair around the central atom is a region of negative charge, and these regions push each other as far apart as possible. Count them from your Lewis structure, and the arrangement follows.

Methane, CH4, has four bonding pairs around carbon, which spread into a tetrahedron with angles of about 109.5 degrees, not the 90 degrees a square drawing suggests. Ammonia, NH3, also has four regions around nitrogen, but one is a lone pair, so the three hydrogens form a low pyramid with angles of about 107 degrees. Water has two bonds and two lone pairs, giving a bent shape of about 104.5 degrees, because lone pairs push slightly harder than bonding pairs. That bent shape is why water molecules are polar, which explains many of water's unusual properties.

The dot diagrams are named after the American chemist Gilbert N. Lewis, who set out the idea of a covalent bond as a shared pair of electrons in a 1916 paper called The Atom and the Molecule. His early sketches pictured electrons sitting at the corners of a cube around each atom, so that a full shell meant all eight corners were filled, an early form of the octet idea. The cube model was soon abandoned, but the shared pair survived. The chemist Irving Langmuir popularised the approach in the following years and helped give the word octet its wide use.

Lewis structures also have a famous failure. Draw oxygen gas, O2, by the usual rules and you get a double bond with every electron neatly paired. Yet liquid oxygen is attracted to a magnet: pour it between the poles of a strong magnet and it clings there, while liquid nitrogen runs straight through. Attraction of this kind requires unpaired electrons, and an oxygen molecule has two. Explaining this needs molecular orbital theory, which treats electrons as spread over the whole molecule. A Lewis structure is best treated as a useful bookkeeping model, not a literal map of where electrons are.

Some teachers use a quick formula to predict how many bonds a molecule needs before you draw anything. First add up the electrons every atom would need for a full shell: eight for most atoms and two for hydrogen. Then subtract the valence electrons actually available. Half the difference is the number of bonds. Try hydrogen cyanide, HCN. Hydrogen needs 2, carbon 8 and nitrogen 8, a total of 18. The atoms supply 1, 4 and 5, which is 10. The difference is 8, so the molecule needs four bonds: one between hydrogen and carbon, and a triple bond between carbon and nitrogen.

The two electrons left over after those eight bonding electrons form one lone pair on nitrogen, and every formal charge comes out at zero. Formaldehyde, CH2O, works the same way: 20 electrons needed and 12 available, so four bonds, meaning two carbon to hydrogen bonds and a carbon to oxygen double bond, with the last four electrons as two lone pairs on oxygen. The shortcut has limits. For boron trifluoride it predicts four bonds, implying a double bond to one fluorine, yet chemists usually draw three single bonds and accept six electrons on boron. Treat the formula as a starting guess, then check octets and formal charges.

Drawing Lewis structures is mostly careful bookkeeping. Count valence electrons first, use two electrons per bond, add lone pairs, check octets where the rule applies, and make sure the total electron count never changes. If several structures work, formal charge can help you judge which representation is more reasonable.