Lewis dot structures are diagrams that show how electrons are arranged around atoms in a molecule or polyatomic ion. They help explain how atoms form chemical bonds, why certain molecules have particular shapes at the electron level, and how to identify lone pairs of electrons that influence chemical behavior.
To draw a Lewis dot structure, count the total number of valence electrons, choose a central atom, connect the atoms with single bonds, distribute the remaining electrons to complete outer electron shells, and form multiple bonds if needed. Finally, check the total electron count and evaluate the formal charges to determine whether the structure is reasonable.
With a consistent method, you can draw Lewis structures for simple molecules such as water and methane, as well as more complex species such as carbon dioxide, nitrate ions, and sulfur dioxide.
Understand what Lewis dot structures represent
A Lewis dot structure represents the valence electrons of atoms, which are the electrons in an atom’s outermost occupied shell. These electrons play the leading role in chemical bonding because they determine how atoms interact with one another.
In a Lewis structure, a line between two atoms represents a shared pair of electrons, called a covalent bond. One line represents a single bond, two lines represent a double bond, and three lines represent a triple bond. Dots placed around an atom represent lone pairs, which are valence electrons not shared in a bond.
For example, water, H₂O, has an oxygen atom bonded to two hydrogen atoms. Oxygen forms two single bonds and retains two lone pairs. Its Lewis structure therefore shows two O–H bonds and four additional electrons arranged as two lone pairs around oxygen.
Lewis structures are models rather than literal pictures of molecules. Electrons do not sit at fixed points around atoms as the dots might suggest. Instead, the diagrams provide a useful way to track valence electrons and predict aspects of chemical bonding.
Step 1: Count the total number of valence electrons
The first step is to determine how many valence electrons are available to place in the structure. Every subsequent step depends on getting this number right.
For many main-group elements, the group number on the periodic table indicates the number of valence electrons. For example, carbon has four, nitrogen has five, oxygen has six, and the halogens, such as fluorine and chlorine, have seven. Hydrogen has one valence electron.
For a neutral molecule, add the valence electrons contributed by every atom. Consider water, H₂O:
- Oxygen contributes 6 electrons.
- Each hydrogen contributes 1 electron.
- Total: 6 + 1 + 1 = 8 valence electrons.
For carbon dioxide, CO₂, carbon contributes four electrons and each oxygen contributes six:
4 + 6 + 6 = 16 valence electrons.
Charged species require an adjustment. For a negatively charged ion, add one electron for each negative charge. For a positively charged ion, subtract one electron for each positive charge.
For example, the nitrate ion, NO₃⁻, contains one nitrogen atom and three oxygen atoms. Its total valence electron count is:
5 + (3 × 6) + 1 = 24 electrons.
The extra electron comes from the ion’s −1 charge.
Write down the total before drawing any bonds. Keeping track of this number makes it easier to catch mistakes later.
Step 2: Choose the central atom
Next, decide which atom should occupy the center of the structure. The central atom is generally the one connected to the greatest number of other atoms.
Carbon commonly serves as the central atom in compounds containing carbon and other nonmetals. In water, oxygen is central because the hydrogen atoms can each form only one ordinary covalent bond. In carbon dioxide, carbon sits between the two oxygen atoms.
Hydrogen is almost always a terminal atom, meaning it appears at the outside of a structure. It can form only one ordinary covalent bond and follows the duet rule, requiring two electrons around it rather than eight.
Halogens such as fluorine and chlorine are also commonly terminal atoms because they usually form one single bond in simple compounds. Oxygen and nitrogen can occupy central positions, depending on the molecule.
A practical guideline is to choose the least electronegative atom as the central atom, excluding hydrogen. Electronegativity describes how strongly an atom attracts shared electrons in a chemical bond. Atoms with lower electronegativity often share electrons with several neighboring atoms, making them suitable central atoms.
This guideline is not absolute. Molecular formulas, known connectivity, and chemical bonding patterns may determine the arrangement more reliably than electronegativity alone.
Step 3: Connect the atoms with single bonds
Once you have chosen the central atom, draw the basic arrangement of the atoms and connect each terminal atom to the central atom with a single bond.
Each single bond represents two electrons. Therefore, subtract two electrons from the total for every bond you draw.
For water, connect oxygen to both hydrogen atoms:
H–O–H
Two single bonds use four electrons. Since water has eight valence electrons, four remain to be placed.
For carbon dioxide, begin with the arrangement:
O–C–O
The two single bonds use four of the 16 available electrons, leaving 12.
At this stage, the drawing establishes which atoms are connected, but it may not yet represent the most stable bonding arrangement. Multiple bonds may be needed later to satisfy the usual electron requirements of the atoms.
Step 4: Complete the outer atoms’ electron shells
After drawing the single bonds, distribute the remaining electrons as lone pairs around the terminal atoms first.
Most main-group atoms in common covalent compounds tend to achieve an octet, or eight electrons in their valence shell. Hydrogen is the main exception in introductory Lewis structures because its first electron shell holds a maximum of two electrons.
Remember that bonding electrons count toward the electron shell of both atoms sharing the bond. A single bond contributes two electrons to each bonded atom’s octet count.
Return to water. The two O–H bonds use four electrons, leaving four. Place those four electrons around oxygen as two lone pairs. Each hydrogen now has two electrons from its bond, and oxygen has eight electrons around it: four in bonds and four in lone pairs.
For carbon dioxide’s initial structure, O–C–O, place the remaining 12 electrons around the two terminal oxygen atoms. Give each oxygen three lone pairs, using six electrons per oxygen. All 16 electrons have now been accounted for.
However, this initial carbon dioxide structure is incomplete in another sense: carbon has only four electrons around it from its two single bonds. Carbon ordinarily favors an octet, so the structure must be revised.
Step 5: Complete the central atom’s octet
Once the terminal atoms have their usual electron shells, examine the central atom. If it has fewer than eight electrons, consider converting one or more lone pairs on neighboring atoms into additional bonds.
Each additional bond uses a pair of electrons that was previously a lone pair. The total number of valence electrons in the structure does not change.
For carbon dioxide, each oxygen has three lone pairs in the initial arrangement. Convert one lone pair from each oxygen into a bonding pair between oxygen and carbon. This produces two double bonds:
O=C=O
Each double bond contains four electrons. Carbon now has eight electrons around it, and each oxygen also has an octet. The structure accounts for all 16 valence electrons.
This process is how you recognize when a single-bond arrangement needs double or triple bonds. Carbon, nitrogen, oxygen, and other atoms that commonly follow the octet rule may require multiple bonds to reach a satisfactory structure.
Do not add extra electrons simply to complete an octet. Instead, rearrange existing electrons into additional bonds. The total electron count must remain consistent with the original formula and charge.
Step 6: Check the total electron count
Before deciding that a structure is complete, count every electron represented in the drawing.
Each bond line represents two electrons, regardless of whether the bond is single, double, or triple. Each lone-pair set represents two electrons as well. Count each shared bonding pair only once when calculating the total number of electrons in the entire structure.
For water, the two single bonds account for four electrons, and the two lone pairs account for another four. The total is eight.
For carbon dioxide, the two double bonds account for eight electrons, and the two lone pairs on each oxygen account for eight more. The total is 16.
The electron count must equal the number calculated in the first step. If it does not, look for a missing lone pair, an extra pair of dots, or an incorrect number of bond lines.
For ions, place the entire Lewis structure inside square brackets and write the overall charge outside the brackets. This convention makes clear that the diagram represents a charged species rather than a neutral molecule.
Step 7: Calculate formal charges
A structure can have the correct number of electrons and still place them in an unfavorable arrangement. Formal charge helps compare possible Lewis structures.
Formal charge is an accounting method that assigns bonding electrons equally between the bonded atoms, even when the atoms have different electronegativities. It does not directly measure an atom’s actual electrical charge in a molecule.
For each atom, calculate formal charge using:
Formal charge = valence electrons − nonbonding electrons − ½(bonding electrons)
Nonbonding electrons are the electrons in lone pairs. Bonding electrons are all the electrons shared in bonds involving that atom.
Consider an oxygen atom in carbon dioxide. Oxygen has six valence electrons, four nonbonding electrons in two lone pairs, and four bonding electrons in a double bond:
6 − 4 − ½(4) = 0
Carbon also has a formal charge of zero. It has four valence electrons, no nonbonding electrons, and eight bonding electrons across its two double bonds:
4 − 0 − ½(8) = 0
The resulting carbon dioxide structure has zero formal charge on every atom.
When several plausible structures are possible, chemists generally favor structures with smaller formal charges and less charge separation, all else being equal. If negative formal charges occur, they are often more favorable on more electronegative atoms.
These are guidelines, not guarantees. Some important molecules and ions have formal charges even in their most useful Lewis structures. The sum of all formal charges must equal the overall charge of the molecule or ion.
Step 8: Put ions in brackets and consider resonance
Polyatomic ions contain multiple covalently bonded atoms and carry an overall electrical charge. Their Lewis structures use the same electron-counting process as neutral molecules, with the appropriate adjustment for charge.
For example, nitrate, NO₃⁻, has 24 valence electrons. Place nitrogen in the center and connect it to three oxygen atoms with single bonds. Complete the oxygen octets and then examine nitrogen’s electron count. Nitrogen has only six electrons from its three single bonds, so one oxygen lone pair can be converted into a double bond.
The resulting structure contains one N=O double bond and two N–O single bonds. With the usual formal-charge assignment, nitrogen carries a +1 formal charge, the two singly bonded oxygen atoms each carry a −1 formal charge, and the double-bonded oxygen carries zero. The formal charges add to −1, matching the ion’s overall charge.
The double bond can be placed between nitrogen and any one of the three oxygen atoms. These alternatives are called resonance structures. They represent different valid electron arrangements for the same arrangement of atomic nuclei.
For nitrate, all three nitrogen–oxygen bonds are equivalent in the actual ion. No single resonance structure fully describes the electron distribution. The real structure is better represented as a resonance hybrid, with electron density distributed across the three N–O bonds.
When drawing a set of resonance structures, keep the positions of the atoms the same and move only electrons. Use a double-headed resonance arrow, ↔, between the structures. This arrow is not an equilibrium arrow: the molecule is not simply switching back and forth between separate forms.
Step 9: Recognize exceptions to the octet rule
The octet rule is a useful starting point, but not every stable species follows it. Recognizing the major exceptions prevents you from forcing a structure into an arrangement that does not fit its chemistry.
Incomplete octets. Some atoms are stable with fewer than eight electrons around them. Boron commonly forms three bonds in compounds such as boron trifluoride, BF₃, leaving boron with six electrons around it in the conventional Lewis structure. Beryllium compounds can also have incomplete octets.
Odd-electron species. Molecules with an odd total number of valence electrons cannot give every atom a paired-electron arrangement. Nitric oxide, NO, has 11 valence electrons, so at least one electron must remain unpaired. Such species are called radicals and often have distinctive chemical reactivity.
Expanded valence shells in conventional structures. Elements in the third period and beyond, including sulfur and phosphorus, are often drawn with more than eight electrons around the central atom in traditional Lewis structures. Sulfur hexafluoride, SF₆, is commonly represented with six S–F single bonds, giving sulfur 12 electrons in its Lewis diagram.
These expanded-octet drawings are useful electron-counting models, but they should not be interpreted as proof that the central atom uses ordinary atomic orbitals to hold a fixed number of localized electrons beyond the octet. Modern bonding descriptions account for electron distribution and bonding in more nuanced ways.
When an ordinary octet arrangement seems impossible, check the electron count, formal charges, and chemical context before concluding that an exception is necessary.
A complete example: Drawing the Lewis structure of ammonia
Ammonia, NH₃, provides a straightforward example of the entire process.
First, count the valence electrons. Nitrogen contributes five, and the three hydrogen atoms contribute one each. The total is eight.
Second, choose the central atom. Nitrogen is central, with three hydrogen atoms surrounding it.
Third, draw single bonds. Connect nitrogen to each hydrogen with three N–H single bonds. These bonds use six electrons, leaving two.
Fourth, place the remaining electrons. Put the remaining two electrons together as one lone pair on nitrogen.
Fifth, check the electron shells. Nitrogen has six electrons from its three bonds and two from its lone pair, for a total of eight. Each hydrogen has two electrons from its bond.
Finally, verify the count and formal charges. The three bonds account for six electrons, and the lone pair accounts for two, giving eight overall. Nitrogen has a formal charge of zero, and each hydrogen also has a formal charge of zero.
The completed structure has three single bonds and one lone pair on nitrogen. The lone pair is especially important because it influences ammonia’s molecular shape and chemical behavior.
Use Lewis structures to understand molecular shape and bonding
Lewis structures do more than track electrons. They provide the starting point for understanding why molecules have particular shapes and how those shapes affect their properties.
Electron pairs repel one another. In the valence-shell electron-pair repulsion model, often called VSEPR theory, bonding pairs and lone pairs arrange themselves to reduce repulsion around a central atom. Lewis structures help identify the number of bonding groups and lone pairs needed for this prediction.
Water has two bonding pairs and two lone pairs around oxygen. These four electron groups arrange themselves approximately in a tetrahedral pattern, but the positions of the atoms produce a bent molecular shape. Ammonia has three bonding pairs and one lone pair around nitrogen, producing a trigonal pyramidal molecular shape. Carbon dioxide has two bonding groups around carbon and no lone pairs on the central atom, so its molecular shape is linear.
Bonding and electron distribution also help explain polarity. A polar bond has an unequal distribution of shared electrons. Whether an entire molecule is polar depends on both the polarity of its bonds and its geometry. In carbon dioxide, the two polar C=O bonds point in opposite directions and cancel one another’s overall dipole moments. Water’s bent shape prevents its O–H bond dipoles from canceling, making water a polar molecule.
Lewis structures cannot fully describe every feature of chemical bonding. They do not directly show three-dimensional geometry, exact bond lengths, or the detailed quantum-mechanical distribution of electrons. Even so, they provide a powerful foundation for understanding covalent bonds, lone pairs, formal charges, resonance, and molecular geometry.
Common mistakes to avoid when drawing Lewis structures
The most reliable way to improve is to check a few common sources of error every time you finish a diagram.
- Miscounting valence electrons. Count the contribution from every atom and adjust for the overall charge before drawing bonds.
- Giving hydrogen an octet. Hydrogen follows the duet rule and generally forms only one single bond in ordinary Lewis structures.
- Forgetting that bonds contain two electrons. A single bond uses two electrons, a double bond four, and a triple bond six.
- Counting shared electrons twice in the total. When checking the structure’s overall electron count, each bonding pair is counted once, even though it contributes to the octet of both bonded atoms.
- Completing terminal octets but ignoring the central atom. If the central atom has an incomplete octet, consider whether multiple bonds are needed.
- Changing the electron total while adding double bonds. Convert an existing lone pair into a bonding pair instead of adding two new electrons.
- Ignoring formal charges. Compare plausible arrangements by checking their formal charges and whether the charges sum to the overall molecular or ionic charge.
- Treating one resonance structure as the entire molecule. When several valid resonance structures exist, the actual electronic structure is represented by their combined contribution, not by a single drawing.
The essential habit is to treat Lewis structures as a controlled accounting process. Count the electrons, establish the atom connections, distribute the electrons, check the shells, and evaluate formal charges. Repeating these steps in order makes even unfamiliar molecules and polyatomic ions much easier to analyze.