Lewis structures show how atoms share electrons and how lone pairs are distributed in a molecule or ion. For polyatomic ions, which are groups of two or more covalently bonded atoms carrying an overall electric charge, drawing an accurate Lewis structure requires accounting for the total number of valence electrons, arranging the atoms, forming bonds, and checking the placement of electrons and formal charges. The key difference from drawing a neutral molecule is that the ion’s charge changes the electron count and must be indicated in the final structure.
A reliable method works for common ions such as ammonium (NH₄⁺), hydroxide (OH⁻), carbonate (CO₃²⁻), and nitrate (NO₃⁻). Although some ions have several valid Lewis structures, the same basic principles apply to all of them.
Understand what a Lewis structure represents
A Lewis structure is a diagram of the valence electrons involved in bonding and the electrons that remain as lone pairs. Valence electrons are the electrons in an atom’s outermost occupied shell. They determine how atoms typically form chemical bonds.
A shared pair of electrons between two atoms is represented by a line, called a bond. A single line represents two shared electrons, while two lines represent a double bond and three lines represent a triple bond. Electrons that belong to one atom rather than being shared in a bond are represented by dots, usually arranged in pairs.
For example, oxygen typically has six valence electrons. In a simple Lewis structure, it often forms two bonds and retains two lone pairs to complete an octet, or eight electrons in its valence shell. Nitrogen commonly forms three bonds and retains one lone pair, while hydrogen generally forms one bond and is stable with two electrons.
These patterns are useful guidelines, not universal rules. Elements in the third period and beyond can appear in Lewis structures that do not follow the simple octet pattern, and some species have incomplete octets or an odd number of electrons.
Polyatomic ions follow these bonding principles but have an overall charge because their total number of electrons differs from the number of protons in their constituent atoms.
Step 1: Count the total number of valence electrons
The first and most important step is to determine how many valence electrons the entire ion contains. Every atom contributes its valence electrons, and the ion’s charge determines whether electrons must be added or removed.
For common main-group elements, the periodic table provides a convenient way to determine valence electron counts. Group 1 elements generally contribute one valence electron, Group 2 elements contribute two, Group 13 elements contribute three, Group 14 elements contribute four, Group 15 elements contribute five, Group 16 elements contribute six, and Group 17 elements contribute seven.
To calculate the total:
- Add the valence electrons contributed by every atom.
- Add electrons for a negative charge.
- Subtract electrons for a positive charge.
A negative charge means the species has more electrons than the corresponding collection of neutral atoms. A positive charge means it has fewer electrons.
Consider the nitrate ion, NO₃⁻. Nitrogen contributes five valence electrons, and each of the three oxygen atoms contributes six. The negative charge adds one electron.
The total is:
5 + 3(6) + 1 = 24 valence electrons.
The carbonate ion, CO₃²⁻, has four valence electrons from carbon and 18 from its three oxygen atoms. Its 2− charge adds two more electrons.
The total is:
4 + 3(6) + 2 = 24 valence electrons.
By contrast, the ammonium ion, NH₄⁺, contains five valence electrons from nitrogen and four from its hydrogen atoms. Its positive charge requires subtracting one electron.
The total is:
5 + 4(1) − 1 = 8 valence electrons.
These totals include every electron that must be represented in the Lewis structure, whether it participates in bonding or remains as a lone pair. If the count is incorrect, the finished structure cannot correctly represent the ion.
Step 2: Choose the central atom and arrange the atoms
Next, determine which atom belongs in the center of the structure. In most polyatomic ions, the central atom is the least electronegative atom other than hydrogen. Electronegativity describes how strongly an atom attracts shared electrons in a chemical bond.
Hydrogen is almost always a terminal atom because it forms only one bond. Halogens, such as fluorine and chlorine, also commonly occupy terminal positions because they typically form one bond in simple Lewis structures.
For example, nitrate is arranged with nitrogen in the center and three oxygen atoms surrounding it. Carbonate has a central carbon atom bonded to three oxygen atoms. In ammonium, nitrogen occupies the center and is bonded to four hydrogen atoms.
Once the atoms are arranged, connect the central atom to each surrounding atom with a single bond. Each single bond represents two electrons.
For nitrate, three nitrogen–oxygen single bonds use six electrons. Since the total electron count is 24, there are 18 electrons left to place.
For carbonate, three carbon–oxygen single bonds also use six electrons, leaving 18 electrons.
For ammonium, four nitrogen–hydrogen single bonds use all eight available electrons. No electrons remain to be placed as lone pairs.
At this stage, the arrangement is a starting framework rather than necessarily the final structure. Additional bonds may be needed to satisfy electron-shell requirements and produce a more favorable distribution of formal charges.
Step 3: Complete the outer atoms’ valence shells
After drawing the initial single bonds, distribute the remaining electrons to the surrounding atoms first. For many common main-group ions, this means giving each terminal oxygen, fluorine, or chlorine atom an octet.
Remember that each bond already supplies two electrons toward the valence shell of both bonded atoms. A terminal oxygen with one single bond generally needs three lone pairs to reach eight electrons around it: two electrons in the bond and six in lone pairs.
Continue adding lone pairs until each suitable terminal atom has a complete valence shell or until no electrons remain.
For nitrate, the three nitrogen–oxygen single bonds use six electrons. Adding three lone pairs to each oxygen requires another 18 electrons, bringing the total to 24. All three oxygen atoms now have octets, but the central nitrogen has only six electrons around it because it shares one pair with each oxygen.
For carbonate, the same initial process uses all 24 electrons and gives each oxygen an octet. However, the central carbon also has only six electrons around it.
This is a common point of confusion: completing the outer atoms does not guarantee that the central atom has a complete valence shell. The next step addresses this problem.
Some ions have different requirements. In hydroxide, OH⁻, the oxygen has one bond to hydrogen and three lone pairs. These account for eight electrons around oxygen, while hydrogen has its required two electrons. The ion has a complete structure without any additional bonds.
Step 4: Form multiple bonds when necessary
If the central atom does not have an appropriate valence shell after the terminal atoms are completed, consider converting a lone pair on a neighboring atom into an additional bond.
To do this, move one lone pair from a terminal atom into the space between that atom and the central atom. The lone pair becomes a shared electron pair, creating a double bond. This changes how the electrons are assigned but does not change the total number of electrons.
For nitrate, the initial structure gives nitrogen only six electrons in its valence shell. Converting one lone pair from an oxygen atom into an additional nitrogen–oxygen bond produces one double bond and two single bonds. Nitrogen now has eight electrons around it.
Carbonate follows a similar pattern. One oxygen forms a double bond with carbon, while the other two remain connected by single bonds. Carbon now has an octet.
A double bond contains four shared electrons, whereas a single bond contains two. When drawing Lewis structures, count the electrons in all bonds and lone pairs to confirm that the total remains equal to the number calculated in the first step.
Multiple bonds are not required in every polyatomic ion. Ammonium, for instance, has four single bonds around nitrogen and no lone pairs on the central atom. Hydroxide has one single bond and three lone pairs on oxygen. In both cases, the available electrons and the atoms’ valence-shell requirements are satisfied without a double bond.
Step 5: Calculate formal charges
A structure may satisfy the electron-counting rules and still contain a less favorable distribution of electrons. Formal charge helps evaluate the competing possibilities.
Formal charge is a bookkeeping tool that assigns each atom a charge based on the electrons it contributes and the electrons assigned to it in a Lewis structure. It assumes that bonding electrons are divided equally between the bonded atoms, even when the actual bond is polar.
The formula is:
Formal charge = valence electrons − nonbonding electrons − ½(bonding electrons)
Nonbonding electrons are the electrons in lone pairs. Bonding electrons are the electrons shared in bonds.
A convenient shortcut is to count the number of bonds rather than individual bonding electrons:
Formal charge = valence electrons − nonbonding electrons − number of bonds
In this shortcut, a double bond counts as two bonds and a triple bond counts as three.
For example, consider an oxygen atom with one single bond and three lone pairs. Oxygen has six valence electrons, six nonbonding electrons, and one bond.
Its formal charge is:
6 − 6 − 1 = −1.
An oxygen atom with one double bond and two lone pairs has six valence electrons, four nonbonding electrons, and two bonds.
Its formal charge is:
6 − 4 − 2 = 0.
The same method applies to nitrogen and carbon. Nitrogen in ammonium has five valence electrons, no nonbonding electrons, and four bonds.
Its formal charge is:
5 − 0 − 4 = +1.
The sum of the formal charges on all atoms must equal the overall charge of the ion. This is an essential check.
In a typical nitrate Lewis structure, the double-bonded oxygen has a formal charge of zero, the two single-bonded oxygen atoms each have a formal charge of −1, and nitrogen has a formal charge of +1. The total is −1, matching the charge of NO₃⁻.
In a typical carbonate Lewis structure, the double-bonded oxygen has a formal charge of zero, the two single-bonded oxygen atoms each have a formal charge of −1, and carbon has a formal charge of zero. The total is −2, matching the charge of CO₃²⁻.
Formal charges are not necessarily the actual charges measured on individual atoms. They are a useful model for comparing Lewis structures, not a complete description of the electron distribution in a real ion.
Step 6: Choose the most reasonable Lewis structure
Some ions can be represented by more than one Lewis structure with the same arrangement of atoms. To evaluate these possibilities, chemists consider the formal charges and the locations of those charges.
In general, a more favorable Lewis structure tends to minimize the magnitude of formal charges and avoid unnecessary separation of positive and negative charges. When a negative formal charge is present, it is generally more favorable on a more electronegative atom than on a less electronegative one.
These guidelines must be applied together with electron-counting and valence-shell requirements. A structure with small formal charges is not automatically preferable if it violates an important bonding or electron-shell constraint.
Nitrate illustrates why this evaluation matters. If nitrogen forms three single bonds to oxygen and all three oxygen atoms have complete octets, nitrogen carries a formal charge of +2 and each oxygen carries −1. The overall charge is still −1, but this arrangement has greater formal-charge separation than a structure containing one nitrogen–oxygen double bond.
The structure with one double bond and two single bonds has a formal charge of +1 on nitrogen, −1 on each single-bonded oxygen, and zero on the double-bonded oxygen. It is generally a more favorable Lewis representation.
For carbonate, a structure with three single bonds gives carbon a formal charge of +1 and each oxygen a formal charge of −1, for an overall charge of −2. Introducing one double bond reduces the formal charge on carbon to zero and gives the double-bonded oxygen a formal charge of zero. The resulting structure is generally preferred over the all-single-bond arrangement.
These comparisons help identify reasonable structures, but formal charges should not be mistaken for exact measurements of electron density.
Step 7: Put the ion in brackets and show its charge
Once the electron distribution and formal charges have been checked, enclose the entire Lewis structure in square brackets and write the overall charge outside the brackets at the upper right.
For example, the ammonium ion is represented as [NH₄]⁺, while the hydroxide ion is represented as [OH]⁻. The carbonate ion is written as [CO₃]²⁻, and nitrate is written as [NO₃]⁻.
When drawing the full Lewis structure, include the bonds and lone pairs inside the brackets. The overall charge outside the brackets indicates the net charge of the entire ion, not the charge on any one atom.
This distinction is important because an ion can contain atoms with different formal charges while carrying a single overall charge. In ammonium, nitrogen has a formal charge of +1, while the hydrogen atoms have formal charges of zero. In nitrate, positive and negative formal charges coexist, but their sum is −1.
The brackets and external charge make clear that the diagram represents a charged group of bonded atoms rather than a neutral molecule.
Understand resonance structures
A single Lewis structure does not always fully describe the bonding in a polyatomic ion. In some ions, more than one equivalent arrangement of multiple bonds and formal charges can be drawn without changing the positions of the atoms.
These alternatives are called resonance structures. They differ in the placement of electrons, not in the arrangement of the atomic nuclei.
Nitrate provides a clear example. A Lewis structure can show a double bond between nitrogen and any one of the three oxygen atoms, with single bonds to the other two. This produces three equivalent resonance structures.
Carbonate also has three equivalent resonance structures, each with a carbon–oxygen double bond to a different oxygen atom.
The real ion is not rapidly switching between separate structures. Instead, its electron distribution is better described as a resonance hybrid, which reflects the contributions of the possible Lewis structures. In nitrate, the three nitrogen–oxygen bonds are equivalent in the actual ion. The same is true of the three carbon–oxygen bonds in carbonate.
Consequently, a single Lewis structure is useful for electron counting and formal-charge analysis, but it can give an incomplete impression of the bonding if it is interpreted as a literal picture of distinct single and double bonds.
When an ion has equivalent resonance structures, drawing all of them helps communicate its bonding more accurately. Use a double-headed resonance arrow between the structures, rather than an equilibrium arrow, because resonance structures are alternative representations of the same species, not separate substances in equilibrium.
Work through common polyatomic ions
The principles become easier to apply when they are used consistently across several familiar ions.
Ammonium (NH₄⁺) contains one nitrogen atom and four hydrogen atoms. The total valence electron count is five + four − one = eight. Place nitrogen in the center and connect it to the four hydrogen atoms with single bonds. All eight electrons are used in bonding, nitrogen has an octet, and each hydrogen has two electrons. Nitrogen has a formal charge of +1, and each hydrogen has a formal charge of zero. Enclose the structure in brackets with a + charge outside.
Hydroxide (OH⁻) contains one oxygen atom and one hydrogen atom. Oxygen contributes six valence electrons, hydrogen contributes one, and the negative charge adds one, giving eight electrons. Draw a single oxygen–hydrogen bond, then place three lone pairs on oxygen. Oxygen has an octet, hydrogen has two electrons, and oxygen carries a formal charge of −1. The overall ion charge is also −1.
Nitrate (NO₃⁻) contains one nitrogen atom and three oxygen atoms. Its total is 24 valence electrons. Place nitrogen in the center and connect it to the three oxygen atoms with single bonds. Complete the oxygen octets, then convert one oxygen lone pair into a nitrogen–oxygen double bond to give nitrogen an octet. The resulting structure has one double bond, two single bonds, a +1 formal charge on nitrogen, and −1 formal charges on the two single-bonded oxygen atoms. The overall charge is −1. Draw all three equivalent resonance structures to represent the equivalent nitrogen–oxygen bonds in the real ion.
Carbonate (CO₃²⁻) contains one carbon atom and three oxygen atoms. Its total is 24 valence electrons. Place carbon in the center, connect it to the three oxygen atoms with single bonds, and complete the oxygen octets. Then convert one oxygen lone pair into a carbon–oxygen double bond to give carbon an octet. The resulting structure has one double bond and two single bonds. The double-bonded oxygen has a formal charge of zero, the two single-bonded oxygen atoms each have a formal charge of −1, and carbon has a formal charge of zero. The sum is −2. Draw all three equivalent resonance structures.
Sulfate (SO₄²⁻) contains one sulfur atom and four oxygen atoms. Sulfur contributes six valence electrons, and the four oxygen atoms contribute 24. The 2− charge adds two more, giving a total of 32 electrons. Place sulfur in the center, connect it to the four oxygen atoms, and complete their octets. One common Lewis representation uses two sulfur–oxygen double bonds and two single bonds. In that representation, the two single-bonded oxygen atoms each have a formal charge of −1, while sulfur and the double-bonded oxygen atoms have formal charges of zero. The formal charges sum to −2. Because sulfur is in the third period, it is commonly represented in introductory Lewis structures with an expanded valence shell. Other representations are possible, and a Lewis structure alone does not fully describe the bonding in sulfate.
These examples demonstrate that the total electron count, the arrangement of atoms, and the formal-charge calculation work together. The exact bonding pattern depends on the elements involved and the electronic structure of the ion.
Recognize common mistakes
Several recurring errors can make a Lewis structure incorrect even when the drawing looks plausible.
The first is mishandling the overall charge during electron counting. Always add electrons for a negative charge and subtract electrons for a positive charge. Treating a 2− ion as though it were neutral removes two electrons from the calculation that should be included.
The second is forgetting that bonding electrons count toward the valence shells of both bonded atoms. A single bond contains two electrons shared between two atoms, so both atoms count those electrons when checking their valence shells. However, the two electrons are counted only once when determining the total number of electrons in the entire structure.
The third is adding lone pairs without keeping track of the remaining electron total. Write down the starting electron count and subtract two electrons for every single bond, double bond, or lone pair added to the structure. A double bond uses four electrons, not two.
Another mistake is assuming that every atom must always have exactly eight electrons. Hydrogen follows the duet rule, while some elements can form stable structures with incomplete octets. Elements in the third period and beyond may be represented with expanded valence shells in conventional Lewis structures. Species with odd numbers of electrons may also contain an atom that does not achieve an octet.
A further mistake is confusing formal charges with the overall ionic charge. Formal charges are assigned to individual atoms, and their sum must match the charge of the entire ion. The charge written outside the brackets represents the total.
Finally, avoid treating a single resonance structure as the only possible electron arrangement when equivalent alternatives exist. Nitrate and carbonate are not accurately described as having one permanently localized double bond and two permanently localized single bonds. Their bonding is delocalized across the oxygen atoms.
Check the completed structure
Before considering a Lewis structure finished, review it systematically.
First, confirm that the total number of electrons represented by all bonds and lone pairs matches the number calculated from the atoms and the overall charge. Count each shared electron only once in the total.
Second, check the valence shell of every atom. Hydrogen should have two electrons, while many common main-group atoms should have eight. Identify any legitimate exceptions rather than forcing an octet where it does not apply.
Third, calculate the formal charge on each atom and add the results. The sum must equal the charge of the polyatomic ion.
Fourth, consider whether multiple bonds are needed to satisfy the valence-shell requirements and whether an alternative arrangement would give a more favorable distribution of formal charges. If equivalent resonance structures exist, recognize that they are different representations of the same ion.
Finally, make sure the entire structure is enclosed in brackets and that the overall charge appears outside the brackets at the upper right.
Following these checks turns Lewis-structure drawing from a matter of guesswork into a repeatable chemical reasoning process. Count the electrons, arrange the atoms, distribute lone pairs, form necessary multiple bonds, evaluate formal charges, and represent the overall charge correctly. With practice, these steps provide a dependable way to draw and interpret Lewis structures for many common polyatomic ions.
