A Lewis structure is a diagram that shows how atoms are connected in a molecule and how their valence electrons are arranged. It uses lines to represent covalent bonds and dots to represent electrons that are not shared in bonds. Learning to draw Lewis structures provides a foundation for understanding molecular shape, chemical bonding, reactivity, and many other concepts in chemistry.
The most reliable way to draw a Lewis structure is to count the total number of valence electrons, arrange the atoms, connect them with single bonds, distribute the remaining electrons, and then form multiple bonds if needed to satisfy the appropriate electron configurations. Finally, check the formal charges and overall charge to determine whether the structure is reasonable.
This method works for most common molecules and polyatomic ions. Some species, particularly those involving transition metals, an odd number of electrons, or expanded valence shells, require additional considerations.
What a Lewis structure shows
Atoms contain positively charged protons, neutral neutrons, and negatively charged electrons. The electrons most directly involved in chemical bonding are called valence electrons. For main-group elements, these are generally the electrons in the outermost occupied energy level.
Lewis structures focus on these valence electrons because their arrangement largely determines how atoms bond. A shared pair of electrons forms a covalent bond, which is commonly represented by a line between two atoms. An unshared pair, called a lone pair, remains associated primarily with one atom and is represented by two dots.
For example, a water molecule, H₂O, has two O–H bonds and two lone pairs on its oxygen atom. Each hydrogen shares one electron with oxygen, while oxygen contributes one electron to each bond. Oxygen also retains four nonbonding electrons, arranged as two lone pairs.
A Lewis structure is not a literal picture of a molecule. Its lines and dots are a simplified representation of electron distribution. It does not show the exact positions of electrons, the three-dimensional shape of the molecule, or the continuous nature of electron density. Nevertheless, it is an effective tool for predicting bonding patterns and identifying important chemical features.
Step 1: Count the total number of valence electrons
The first step is to determine how many valence electrons are available for the entire molecule or ion. Every subsequent decision depends on this count.
For main-group elements, the group number in the periodic table can usually be used to determine the number of valence electrons:
- Group 1: 1 valence electron
- Group 2: 2 valence electrons
- Group 13: 3 valence electrons
- Group 14: 4 valence electrons
- Group 15: 5 valence electrons
- Group 16: 6 valence electrons
- Group 17: 7 valence electrons
- Group 18: 8 valence electrons, except helium, which has 2
Multiply each element’s valence electron count by the number of atoms of that element, then add the results.
Consider carbon dioxide, CO₂. Carbon contributes four valence electrons, and each oxygen contributes six.
Total valence electrons = 4 + 2(6) = 16
Therefore, a Lewis structure for CO₂ must account for 16 valence electrons.
If the species is an ion, adjust the count for its charge. Add one electron for every negative charge and subtract one electron for every positive charge.
For example, the carbonate ion, CO₃²⁻, contains one carbon atom and three oxygen atoms. Its total is:
Total valence electrons = 4 + 3(6) + 2 = 24
The extra two electrons come from the ion’s overall 2− charge.
A useful check is to remember that the electron count must match both the chemical formula and the charge. Counting incorrectly at this stage can lead to an apparently plausible diagram that fails every later check.
Step 2: Choose the central atom and arrange the atoms
Next, decide which atom belongs in the center of the structure. The central atom is generally the one that can form the greatest number of bonds, which is often the least electronegative element other than hydrogen.
Electronegativity describes an atom’s tendency to attract shared electrons in a chemical bond. Atoms with lower electronegativity commonly occupy central positions, while more electronegative atoms often appear around the outside.
For example, carbon is the central atom in CO₂, giving the initial arrangement O–C–O. In ammonia, NH₃, nitrogen is central, with three hydrogen atoms attached to it.
Hydrogen is almost always a terminal atom because it can accommodate only two electrons in its valence shell. It forms one single covalent bond and does not serve as the central atom in ordinary Lewis structures.
Halogens such as fluorine, chlorine, bromine, and iodine are also commonly terminal atoms because they usually form one bond. Oxygen often appears at the outside of a structure, although it can occupy a central position in compounds such as water and ozone.
For molecules containing several different elements, the chemical formula, known connectivity, and typical bonding patterns help determine the arrangement. The least electronegative atom is a useful starting point, not an infallible rule. Some molecules have rings, chains, or several possible arrangements, and a Lewis structure cannot always be determined from the formula alone.
Step 3: Connect the atoms with single bonds
Once the atoms are arranged, connect each bonded pair with a single line. Each line represents two shared electrons.
For example, the initial structure for water is H–O–H. Two single bonds account for four electrons because each bond contains two electrons.
For ammonia, three N–H single bonds account for six electrons. For CO₂, the initial O–C–O arrangement contains two single bonds and uses four electrons.
Subtract the electrons used in these bonds from the total valence electron count. The remaining electrons will be placed as lone pairs or used to create additional bonds.
For CO₂, 16 total electrons minus four bonding electrons leaves 12 electrons to distribute.
At this stage, the single-bond arrangement is only a starting point. Some molecules can be represented adequately with single bonds, but others require double or triple bonds to give their atoms reasonable electron configurations.
Step 4: Complete the outer atoms’ valence shells
Place the remaining electrons around the outer atoms first, usually as lone pairs. For most main-group elements, the goal is to give each atom an octet, meaning eight electrons in its valence shell.
Hydrogen is the key exception. It follows the duet rule, requiring only two electrons.
Each single bond contributes two electrons to the count around both atoms connected by that bond. A pair of dots contributes two electrons to the count around the atom on which it is placed.
Return to the initial structure of CO₂, O–C–O. Each oxygen has one single bond, accounting for two electrons around it. To reach an octet, each oxygen needs six additional electrons, or three lone pairs.
Placing three lone pairs on each oxygen uses all 12 remaining electrons. The resulting structure gives both oxygen atoms octets, but carbon has only four electrons around it from its two single bonds.
This illustrates why completing the outer atoms is not always enough. The central atom must also be checked, and multiple bonds may be necessary.
Step 5: Place any remaining electrons on the central atom
After the outer atoms have appropriate electron configurations, place any electrons still unassigned on the central atom.
This step is especially important when the central atom has not reached an octet. If all electrons have already been placed but the central atom has fewer than eight electrons, consider converting one or more lone pairs on neighboring atoms into shared bonding pairs.
Some molecules have enough electrons to give the central atom an octet immediately. For example, in methane, CH₄, carbon forms four single bonds to hydrogen. These four bonds account for eight electrons around carbon, while each hydrogen has two.
Other structures require additional bonds, and a few valid molecules have central atoms that do not follow the octet rule. The correct approach is to apply the usual rules first, then consider exceptions when the standard method does not produce a reasonable structure.
Step 6: Form double or triple bonds when necessary
If the central atom has fewer than eight electrons and all available electrons have been assigned, check whether an outer atom can share a lone pair with it to form an additional bond.
A double bond contains two shared electron pairs, or four electrons. A triple bond contains three shared electron pairs, or six electrons.
Carbon dioxide provides a classic example. In the initial single-bond structure, carbon has only four electrons around it. Convert one lone pair from each oxygen into an additional shared pair with carbon. This produces the structure O=C=O.
Each oxygen now has two lone pairs and a double bond. Each oxygen has eight electrons around it, and carbon has eight electrons from its two double bonds. All 16 valence electrons are accounted for.
Nitrogen gas, N₂, offers another example. Each nitrogen atom has five valence electrons, giving 10 in total. A triple bond between the atoms uses six electrons, and each nitrogen retains one lone pair, using the remaining four electrons.
The resulting structure is :N≡N:, with one lone pair on each nitrogen. Both atoms have octets.
Multiple bonds are not simply decorative additions. They allow atoms to share electrons in ways that produce more favorable electron distributions. In many common molecules, they are essential for satisfying the octet rule and reducing formal charges.
Step 7: Calculate formal charges
Once a complete structure has been drawn, calculate the formal charge on each atom. Formal charge is a bookkeeping tool that helps evaluate how electrons are assigned in a Lewis structure.
The formula is:
Formal charge = valence electrons − nonbonding electrons − ½(bonding electrons)
Here, nonbonding electrons are the dots assigned to the atom, while bonding electrons are those in the lines connecting it to other atoms. Because a bond contains two electrons, the atom is assigned half of the bonding electrons for formal-charge purposes.
A convenient alternative is:
Formal charge = valence electrons − lone-pair electrons − number of bonds
The second form works when each line represents one bond, including the individual bonds in double and triple bonds.
For carbon dioxide in its O=C=O structure, carbon has four valence electrons, no lone-pair electrons, and four bonds when the double bonds are counted individually. Its formal charge is 4 − 0 − 4 = 0.
Each oxygen has six valence electrons, four nonbonding electrons, and two bonds. Its formal charge is 6 − 4 − 2 = 0.
All atoms therefore have a formal charge of zero.
By contrast, an alternative CO₂ structure with two single bonds and three lone pairs on each oxygen gives carbon a formal charge of +2 and each oxygen a formal charge of −1. The charges add to zero, but the arrangement has an incomplete octet on carbon and a less favorable formal-charge distribution.
When comparing possible Lewis structures, the most favorable options generally satisfy the following guidelines:
- Formal charges should be as small as reasonably possible.
- Negative formal charges are generally more favorable on more electronegative atoms.
- The sum of all formal charges must equal the overall charge of the molecule or ion.
- For many common main-group molecules, structures that satisfy the octet rule are preferred when they also have reasonable formal charges.
These are guidelines rather than absolute laws. A structure with formal charges is not necessarily incorrect, and some important molecules require them.
Step 8: Check the total electron count and overall charge
Before accepting a structure, verify that every valence electron has been accounted for. Count two electrons for each bond and two for each lone pair. The total must equal the number calculated at the beginning.
Then add the formal charges on all atoms. Their sum must equal the overall charge of the species.
For a neutral molecule, the formal charges must sum to zero. For a polyatomic ion, they must sum to the ion’s stated charge.
If the sum does not match, revisit the electron count, the number of bonds, the placement of lone pairs, and any multiple bonds. A missing lone pair, an incorrect bond order, or a forgotten charge adjustment can produce the discrepancy.
Also verify that hydrogen has no more than two electrons around it and that the other atoms follow reasonable valence patterns. Most second-period elements, including carbon, nitrogen, oxygen, and fluorine, cannot exceed an octet in ordinary Lewis structures. If one of these atoms appears to have ten or twelve electrons around it, the structure needs to be reconsidered.
A complete example: Drawing the Lewis structure of water
Water, H₂O, demonstrates the method without requiring multiple bonds.
First, count the valence electrons. Oxygen contributes six, and the two hydrogen atoms contribute one each, giving eight electrons in total.
Next, place oxygen in the center and connect it to both hydrogen atoms. The two O–H single bonds use four electrons, leaving four.
Place the remaining four electrons on oxygen as two lone pairs. Oxygen now has eight electrons around it when its two bonds and two lone pairs are counted. Each hydrogen has two electrons from its bond.
The final Lewis structure has two single bonds and two lone pairs on oxygen. The formal charge on oxygen is 6 − 4 − 2 = 0. Each hydrogen has a formal charge of 1 − 0 − 1 = 0.
The structure uses all eight valence electrons, satisfies the duet rule for hydrogen and the octet rule for oxygen, and has a total charge of zero.
Although the Lewis structure shows the bonding arrangement, it does not fully describe water’s three-dimensional shape. The two lone pairs repel the bonding electron pairs, producing a bent molecular geometry. This distinction illustrates how Lewis structures provide a starting point for understanding molecular shape without directly depicting it.
A complete example: Drawing the carbonate ion
The carbonate ion, CO₃²⁻, illustrates how to handle a polyatomic ion with multiple valid Lewis structures.
Begin by counting the electrons. Carbon contributes four, the three oxygen atoms contribute 18, and the 2− charge adds two more, for a total of 24 valence electrons.
Place carbon in the center and connect it to all three oxygen atoms with single bonds. These three bonds use six electrons, leaving 18.
Give each oxygen three lone pairs. This uses the remaining 18 electrons and gives every oxygen an octet. However, carbon has only six electrons around it from its three single bonds.
Convert one lone pair from one oxygen into an additional bond with carbon. The resulting structure contains one C=O double bond and two C–O single bonds. Carbon now has eight electrons around it, and all three oxygen atoms have octets.
Calculate the formal charges. Carbon has a formal charge of zero. The double-bonded oxygen has a formal charge of zero, while each single-bonded oxygen has a formal charge of −1. The charges add to −2, matching the ion’s overall charge.
Any one of the three oxygen atoms can be assigned the double bond in an equivalent Lewis structure. These alternatives are called resonance structures. They represent different valid electron assignments for the same arrangement of atoms.
In the actual carbonate ion, the three carbon–oxygen bonds are equivalent rather than one being permanently double and the other two permanently single. The bonding electrons are delocalized, meaning their distribution extends across more than one bond. The resonance structures are a way to represent this shared electron distribution within the limits of Lewis notation.
When drawing a polyatomic ion, place brackets around the entire structure and write the overall charge outside the brackets. This makes clear that the charge belongs to the whole ion.
Understand the octet rule and its exceptions
The octet rule is one of the most useful guides for drawing Lewis structures, but it is not universal. Many common atoms form bonds that allow them to achieve eight valence-shell electrons, while other species have stable or chemically meaningful structures that do not follow this pattern.
Hydrogen follows the duet rule because its first electron shell can hold only two electrons. Helium also has a filled shell with two electrons.
Some molecules have an incomplete octet. Boron trifluoride, BF₃, is a familiar example. Boron forms three single bonds to fluorine, giving it six electrons around it. Although boron has fewer than eight electrons, this is a useful representation of the molecule. In some chemical reactions, BF₃ can accept an electron pair from another species.
Other molecules contain an odd total number of valence electrons, making it impossible for every electron to be paired. Nitric oxide, NO, has 11 valence electrons: five from nitrogen and six from oxygen. Its Lewis structure must include an unpaired electron. Such species are called radicals and may be particularly reactive, although their behavior depends on their specific electronic structure.
Some elements in the third period and beyond are traditionally represented with more than eight electrons around the central atom in Lewis structures. Sulfur hexafluoride, SF₆, is commonly drawn with six S–F single bonds, giving sulfur 12 electrons in the Lewis representation. Such structures are useful bookkeeping models, but the bonding in these molecules is more complex than a simple picture of localized electron pairs. The expanded-octet notation should not be interpreted as proof that an atom has literally expanded a set of ordinary valence orbitals to hold extra electron pairs in the way the simplest Lewis model suggests.
Transition-metal compounds often require additional approaches because their bonding and electron configurations do not fit the basic main-group rules neatly. For these species, Lewis structures alone may not adequately describe the chemistry.
Use resonance when one Lewis structure is not enough
Sometimes a molecule or ion can be represented by two or more Lewis structures that differ only in the placement of electrons, not the positions of the atoms. These are resonance structures.
Resonance is especially important when a double bond, lone pair, or formal charge can be placed in more than one equivalent position. The carbonate ion is one example. The nitrate ion, NO₃⁻, also has three equivalent resonance structures in which the double bond can be assigned to any of the three oxygen atoms.
The actual electronic structure is not a molecule rapidly switching between separate drawings. Instead, its electron distribution is better described by a combined picture that reflects contributions from the possible resonance structures. This combined description is called a resonance hybrid.
When drawing resonance structures, keep the atoms in the same positions and move only electrons. The number of valence electrons and the overall charge must remain unchanged. Valid resonance structures may differ in bond orders, lone-pair positions, and formal charges, but they must represent the same atomic framework.
Resonance is not necessary for every molecule. It becomes useful when no single Lewis structure adequately represents the distribution of electrons or when several comparable structures contribute to the bonding.
Common mistakes to avoid
One frequent mistake is counting the wrong number of valence electrons. Always calculate the total before drawing bonds, and remember to add electrons for negative charges or subtract them for positive charges.
Another is giving hydrogen more than one bond. Hydrogen ordinarily forms one covalent bond and has a duet rather than an octet. Fluorine and the other halogens also typically form one bond in common neutral molecules, although their bonding behavior depends on the compound.
A third mistake is assuming that all atoms must always have eight electrons. The octet rule works well for many familiar compounds, but incomplete octets, radicals, and expanded-octet representations are important exceptions. Forcing every atom into an octet can lead to an incorrect electron count or unreasonable formal charges.
It is also easy to forget that a line represents two electrons. A double bond uses four electrons, and a triple bond uses six. Every line and lone pair must be included when checking the total.
Finally, do not confuse formal charge with the actual charge distribution in a molecule. Formal charges are calculated using an idealized assignment of bonding electrons, while real electrons are distributed according to the molecule’s electronic structure. Formal charge helps compare Lewis structures, but it is not a direct measurement of the electrical charge concentrated on an atom.
What a Lewis structure can and cannot tell you
A Lewis structure is most useful for identifying which atoms are bonded, how many bonds connect them, where lone pairs occur, and whether formal charges or resonance are important. It can also help predict where a molecule may accept or donate electron pairs and provide a starting point for understanding chemical reactions.
However, a Lewis structure does not fully determine molecular geometry, bond angles, bond lengths, or three-dimensional electron density. To estimate molecular shape, chemists often use valence-shell electron-pair repulsion theory (VSEPR), which predicts that electron groups around a central atom tend to arrange themselves to reduce repulsion. More detailed descriptions of bonding may require molecular orbital theory or computational methods.
The central lesson is that drawing a Lewis structure is a systematic accounting exercise guided by chemical principles. Count the valence electrons, arrange and connect the atoms, distribute the electrons, form multiple bonds when appropriate, and check formal charges and the overall charge. When the ordinary rules fail, consider resonance, radicals, incomplete octets, or other exceptions rather than forcing the molecule into an unsuitable pattern. This approach makes Lewis structures a practical and reliable foundation for understanding how many molecules and ions bond.
