How do I know which atom goes in the center of a Lewis structure?

Choose the least electronegative atom, which is usually the one that appears earliest in a period or lowest in a group, and often the atom that appears only once in the formula. Hydrogen is never central because it can form only one bond, and halogens are almost never central for the same reason. In CH_2O, for example, carbon is central; in H_2SO_4, sulfur is central and the hydrogens attach to oxygens.

What if I run out of electrons before every atom has an octet?

That is the signal to form a multiple bond. You cannot add electrons, but you can relocate them: take a lone pair from an outer atom that already has an octet and turn it into a second shared pair with the electron-deficient central atom. Repeat until the central atom reaches eight. Then recount the total to make sure it still matches your original valence electron sum.

What is the difference between a covalent bond and an ionic bond?

An ionic bond forms when a metal transfers electrons to a nonmetal, creating oppositely charged ions held together by electrostatic attraction. A covalent bond forms between two nonmetals that both pull strongly on electrons, so the electrons are shared rather than transferred. Covalent compounds exist as discrete molecules with defined formulas, while ionic compounds form extended lattices described by empirical formulas.

Do I need to use formal charge on every structure?

No. For most simple molecules such as H_2O, NH_3, CH_4, and CO_2, following the five-step procedure gives the accepted structure directly. Formal charge is a tie-breaker you reach for when more than one arrangement satisfies the octet rule, such as deciding where a double bond belongs in SCN^- or which atom is central. The preferred structure has formal charges closest to zero, with any negative value on the most electronegative atom.

CHEM-4.2

Covalent Bonds & Lewis Structures

Learn how nonmetals share valence electrons, count electrons correctly, and draw Lewis structures with single, double, and triple bonds — step by step.

What you'll do in this lesson

A voice-first session with the Crimsora tutor on Covalent Bonds & Lewis Structures, then targeted practice and FRQs — with the tutor adapting to where you get stuck.

Start Covalent Bonds & Lewis Structures free

What this lesson covers

When a metal and a nonmetal meet, electrons get transferred and ions form. But when two nonmetals meet, neither atom can pull electrons away completely — both hold onto their own valence electrons tightly. The compromise is a covalent bond: a shared pair of electrons that counts toward the octet of both atoms at once.

Lewis structures are the bookkeeping system chemists use to show exactly where those shared electrons went. A correct Lewis structure accounts for every single valence electron in the molecule, either as a bonding pair between two atoms or as a lone pair sitting on one atom. Nothing appears from nowhere, and nothing disappears. In this lesson you will learn the counting procedure, how to place atoms in a skeleton, and what to do when the electrons run short and the molecule needs a double or triple bond. These structures become the foundation for predicting molecular shape and polarity in the lessons that follow.

Why Nonmetals Share Instead of Transfer

A nonmetal atom has a high electronegativity — a strong pull on electrons — and usually needs only one to three more electrons to reach a noble-gas configuration. When two such atoms approach, neither one can strip electrons from the other, because both are pulling hard. Instead, the valence orbitals overlap and a pair of electrons becomes shared territory, attracted to both nuclei at once.

That shared pair is the covalent bond. It is stable because the two positive nuclei are both attracted to the same concentrated region of negative charge between them, which lowers the total energy of the system. Pull the atoms too close and nucleus-nucleus repulsion takes over; too far apart and the attraction weakens. The distance where energy is lowest is the bond length, and the energy needed to pull the atoms back apart is the bond energy.

The payoff for each atom is a filled valence shell. In

F_2

, each fluorine brings seven valence electrons. One electron from each atom forms the shared pair, and now each fluorine can count eight electrons around itself — six of its own as lone pairs plus the two shared. Neither atom actually owns eight, but each one experiences eight. This is the octet rule.

A common misconception is that covalent compounds have no charged parts at all. Sharing is often unequal, and that unequal sharing creates partial charges. That is bond polarity, and it comes later in this unit. For now, treat every bond in a Lewis structure simply as a shared pair, drawn as a line.

Building a Lewis Structure Step by Step

The procedure is mechanical once you trust it. First, count total valence electrons by adding the group-based valence count for every atom. Group 1 contributes 1, Group 2 contributes 2, Group 13 contributes 3, and Groups 14 through 17 contribute 4, 5, 6, and 7 respectively. If the species is an ion, add one electron for each negative charge and subtract one for each positive charge.

Second, draw a skeleton. The least electronegative atom (never hydrogen) goes in the center, and the other atoms surround it. Hydrogen and halogens are almost always on the outside because they typically form only one bond. In oxyacids like

HNO_3

, hydrogen attaches to an oxygen, not to the central atom.

Third, place one bonding pair between each connected pair of atoms and subtract two electrons from your total for each bond drawn.

Fourth, distribute the remaining electrons as lone pairs on the outer atoms first, giving each an octet (hydrogen needs only two, a duplet). Then put any electrons still left over on the central atom.

Fifth, check every atom. If the central atom is short of an octet, convert a lone pair from an adjacent atom into another bonding pair.

Step	Action	Running check
1	Sum valence electrons, adjust for charge	Total is fixed and cannot change
2	Choose central atom, draw skeleton	H and halogens go outside
3	One pair per bond	Subtract 2 per bond
4	Lone pairs to outer atoms, then center	Outer atoms reach 8 (H reaches 2)
5	Fix short octets with multiple bonds	Total still matches Step 1

Where students most often go wrong is skipping Step 5's recount. If your finished drawing contains more or fewer dots and lines than Step 1 allowed, the structure is wrong no matter how good it looks.

When One Bond Is Not Enough: Double and Triple Bonds

Sometimes you finish distributing electrons and the central atom has only four or six around it. You cannot invent more electrons — the total was fixed in Step 1. The fix is to move electrons you already have: take a lone pair from an outer atom and slide it into the bonding region, creating a second shared pair between the same two atoms.

In

CO_2

there are

4 + 6 + 6 = 16

valence electrons. After two single bonds and full octets on both oxygens, carbon has only four electrons. Converting one lone pair from each oxygen into a bonding pair produces

O=C=O

, where carbon now has eight and each oxygen still has eight (two bonding pairs plus two lone pairs). The electron count is unchanged at 16.

In

N_2

there are 10 valence electrons. A single bond leaves each nitrogen with six. Two more shared pairs are required, giving a triple bond with one lone pair on each nitrogen.

Bond type	Shared pairs	Relative length	Relative strength
Single	1	Longest	Weakest
Double	2	Shorter	Stronger
Triple	3	Shortest	Strongest

The pattern to remember: more shared pairs means more electron density pulling the nuclei together, so the bond is shorter and harder to break. That is why

N_2

in the atmosphere is famously unreactive.

A frequent error is adding a double bond to an atom that already had an octet, which produces ten electrons on that atom. Only make a multiple bond when an atom is genuinely short. Another error is forgetting that hydrogen never participates in double bonds; it holds a maximum of two electrons.

Exceptions, Ions, and Checking Your Work

The octet rule is a strong pattern, not a law. Hydrogen is complete with two electrons. Beryllium in

BeCl_2

is often drawn with four, and boron in

BF_3

with six — these are electron-deficient central atoms, and their hunger for electrons explains why

BF_3

readily grabs a lone pair from another molecule. Atoms in period 3 and beyond, such as sulfur and phosphorus, can hold more than eight in an expanded octet, as in

SF_6

PCl_5

. Molecules with an odd total number of valence electrons, like

NO

with 11, must leave one electron unpaired.

For polyatomic ions, adjust the total first and then enclose the finished structure in brackets with the charge written outside. For

OH^-

, the count is

6 + 1 + 1 = 8

: one bonding pair plus three lone pairs on oxygen, drawn in brackets with a negative charge.

To verify a structure, count the electrons shown — each line is two, each dot pair is two — and confirm the total matches Step 1. Then circle each atom mentally and count its surroundings: eight for most, two for hydrogen, with justified exceptions.

A useful cross-check is formal charge, calculated as valence electrons minus lone-pair electrons minus half the bonding electrons. The best structure keeps formal charges close to zero and puts any negative formal charge on the most electronegative atom. This is how chemists choose between competing arrangements, such as deciding that carbon rather than oxygen sits in the center of

CO_2

Key terms

Covalent bond.: An attraction formed when two atoms share one or more pairs of valence electrons, with both nuclei attracted to the same shared region.
Lewis structure.: A diagram showing every valence electron in a molecule or polyatomic ion as either a bonding pair (line) or a lone pair (two dots).
Lone pair.: A pair of valence electrons localized on a single atom and not shared with any other atom.
Octet rule.: The tendency of main-group atoms to be surrounded by eight valence electrons, matching a noble-gas configuration; hydrogen is satisfied with two.
Bond order.: The number of shared electron pairs between two atoms: 1 for a single bond, 2 for a double bond, 3 for a triple bond.
Bond length.: The average distance between two bonded nuclei, corresponding to the lowest-energy separation; it decreases as bond order increases.
Formal charge.: A bookkeeping value equal to valence electrons minus lone-pair electrons minus half of the bonding electrons, used to compare possible structures.
Expanded octet.: An arrangement in which a central atom from period 3 or below is surrounded by more than eight valence electrons, as in $SF_6$ .

Worked example

Draw the complete Lewis structure for hydrogen cyanide,

HCN

, and state the bond order of each bond.

Step 1, count valence electrons. Hydrogen (Group 1) brings 1, carbon (Group 14) brings 4, nitrogen (Group 15) brings 5. The total is

1 + 4 + 5 = 10

valence electrons. The molecule is neutral, so no adjustment is needed.

Step 2, draw the skeleton. Hydrogen can form only one bond, so it must be terminal. Carbon is less electronegative than nitrogen and can form four bonds, so carbon is central: the arrangement is

H-C-N

.

Step 3, place one bonding pair per connection. Two bonds use

2 \times 2 = 4

electrons, leaving

10 - 4 = 6

electrons to distribute.

Step 4, fill the outer atoms. Hydrogen already has two electrons from its bond and needs nothing more. Nitrogen gets the remaining 6 electrons as three lone pairs, which gives it an octet. Zero electrons remain.

Step 5, check the central atom. Carbon has only the two bonding pairs, or 4 electrons — it is 4 short of an octet. Convert two of nitrogen's lone pairs into bonding pairs with carbon. Now carbon is joined to nitrogen by three shared pairs.

Final structure:

H-C \equiv N

with one lone pair remaining on nitrogen. Verify the count: the single bond (2) plus the triple bond (6) plus nitrogen's lone pair (2) equals 10 electrons, matching Step 1. Carbon has 8 (two from the H bond, six from the triple bond), nitrogen has 8 (six bonding, two lone), hydrogen has 2.

Bond orders: the

H-C

bond has bond order 1, and the

C \equiv N

bond has bond order 3.

Practice questions

How many total valence electrons must be accounted for in the Lewis structure of the carbonate ion,

CO_3^{2-}

Answer: 24

Carbon is in Group 14 and contributes 4. Each of the three oxygens is in Group 16 and contributes 6, for 18. That gives 22 from the neutral atoms. The 2− charge means two extra electrons were gained, so add 2 to reach 24. Forgetting to adjust for the charge is the most common slip here and produces 22. The finished structure uses one carbon-oxygen double bond and two single bonds, with the two singly bonded oxygens carrying a negative formal charge.

Nitrogen gas,

N_2

, contains a triple bond, while fluorine gas,

F_2

, contains a single bond. Explain why, using valence electron counts, and predict which bond is shorter.

Answer: Each nitrogen has 5 valence electrons and needs 3 more for an octet, so the two atoms must share three pairs; each fluorine has 7 valence electrons and needs only 1 more, so one shared pair is enough. The $N \equiv N$ bond is shorter.

Start from the counts.

N_2

has 10 valence electrons total; a single bond would leave each nitrogen with only 6, so lone pairs must be converted into bonding pairs until each atom reaches 8, producing a triple bond plus one lone pair per atom.

F_2

has 14 valence electrons; one shared pair plus three lone pairs on each atom already gives each fluorine 8, so no multiple bond is needed. Higher bond order means more shared electron density between the nuclei, pulling them closer together, so the triple bond in

N_2

is both shorter and much stronger — which is why atmospheric nitrogen is so unreactive.

A student draws

H_2O

with two

O-H

single bonds and four lone pairs on oxygen. Identify the error and give the correct structure.

Answer: The student used 12 electrons instead of 8. Water has only two lone pairs on oxygen, not four.

Count first: hydrogen contributes 1 each and oxygen contributes 6, so the total is

1 + 1 + 6 = 8

valence electrons. Two

O-H

bonds use 4, leaving 4 electrons — exactly two lone pairs for oxygen. With two bonding pairs and two lone pairs, oxygen has 8 electrons and each hydrogen has 2, so the structure is complete. Four lone pairs would put 12 electrons around oxygen, which both breaks the octet rule and invents four electrons the molecule does not have. Always recount at the end against the Step 1 total.

FAQ

How do I know which atom goes in the center of a Lewis structure?: Choose the least electronegative atom, which is usually the one that appears earliest in a period or lowest in a group, and often the atom that appears only once in the formula. Hydrogen is never central because it can form only one bond, and halogens are almost never central for the same reason. In $CH_2O$ , for example, carbon is central; in $H_2SO_4$ , sulfur is central and the hydrogens attach to oxygens.
What if I run out of electrons before every atom has an octet?: That is the signal to form a multiple bond. You cannot add electrons, but you can relocate them: take a lone pair from an outer atom that already has an octet and turn it into a second shared pair with the electron-deficient central atom. Repeat until the central atom reaches eight. Then recount the total to make sure it still matches your original valence electron sum.
What is the difference between a covalent bond and an ionic bond?: An ionic bond forms when a metal transfers electrons to a nonmetal, creating oppositely charged ions held together by electrostatic attraction. A covalent bond forms between two nonmetals that both pull strongly on electrons, so the electrons are shared rather than transferred. Covalent compounds exist as discrete molecules with defined formulas, while ionic compounds form extended lattices described by empirical formulas.
Do I need to use formal charge on every structure?: No. For most simple molecules such as $H_2O$ , $NH_3$ , $CH_4$ , and $CO_2$ , following the five-step procedure gives the accepted structure directly. Formal charge is a tie-breaker you reach for when more than one arrangement satisfies the octet rule, such as deciding where a double bond belongs in $SCN^-$ or which atom is central. The preferred structure has formal charges closest to zero, with any negative value on the most electronegative atom.

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