Valence electrons are the electrons in an atom’s outermost occupied energy level that play the main role in chemical bonding. To find them for a Lewis structure, identify each element, determine how many valence electrons its atoms contribute, and add those numbers together. If the structure represents an ion, adjust the total for its charge. This total tells you how many electrons to account for when drawing bonds and lone pairs.
For most main-group elements, the periodic table provides the quickest way to find valence electrons. Once you know the pattern, you can calculate the total for a molecule or polyatomic ion and use it to build a Lewis structure that shows how its atoms share electrons.
What valence electrons are and why they matter
Atoms contain positively charged protons and neutral neutrons in their nuclei, surrounded by negatively charged electrons. These electrons occupy regions of space called orbitals, which are organized into energy levels. The electrons in the outermost occupied energy level are generally the most important for chemical bonding.
Atoms form chemical bonds by sharing, gaining, or losing electrons. In a covalent bond, two atoms share a pair of electrons. Electrons that participate in bonding or remain localized around a single atom as nonbonding pairs are represented in Lewis structures.
A Lewis structure is a diagram that shows the connections between atoms and the distribution of their valence electrons. Shared electron pairs are usually represented by lines between atoms, while lone pairs—pairs of valence electrons not used in a bond—are shown as dots.
Finding the correct number of valence electrons is the first essential step. If the electron total is wrong, a Lewis structure may contain an incorrect number of bonds, missing lone pairs, or inappropriate formal charges.
Use the periodic table to find valence electrons
For elements in groups 1, 2, and 13 through 18 of the periodic table, the group number provides a convenient way to determine the number of valence electrons. These are known as the main-group elements.
| Periodic table group | Valence electrons | Examples |
|---|---|---|
| Group 1 | 1 | Hydrogen, sodium |
| Group 2 | 2 | Beryllium, magnesium |
| Group 13 | 3 | Boron, aluminum |
| Group 14 | 4 | Carbon, silicon |
| Group 15 | 5 | Nitrogen, phosphorus |
| Group 16 | 6 | Oxygen, sulfur |
| Group 17 | 7 | Fluorine, chlorine |
| Group 18 | 8, except helium, which has 2 | Neon, argon, helium |
For groups 13 through 18, subtract 10 from the group number to obtain the number of valence electrons. For example, oxygen is in group 16, so it has six valence electrons. Chlorine is in group 17, so it has seven.
For groups 1 and 2, the group number directly gives the usual valence-electron count. Hydrogen has one valence electron, while magnesium has two.
Helium is an important exception to the group pattern. Although it appears in group 18, its only occupied energy level can hold just two electrons, and helium has two valence electrons.
This shortcut works well for most introductory Lewis structures involving main-group elements. Transition metals, which occupy the central block of the periodic table, can have more complicated electron configurations, so their bonding may require additional rules.
Count the valence electrons in a molecule
A molecule contains atoms that may contribute different numbers of valence electrons. To determine the total, multiply the number of atoms of each element by that element’s valence-electron count, then add the contributions.
The general rule is:
Total valence electrons = sum of the valence electrons contributed by every atom.
Consider water, H₂O. Each hydrogen atom contributes one valence electron, and oxygen contributes six.
- Two hydrogen atoms: 2 × 1 = 2 electrons
- One oxygen atom: 1 × 6 = 6 electrons
- Total: 2 + 6 = 8 valence electrons
Therefore, a Lewis structure of water must account for eight valence electrons.
Carbon dioxide, CO₂, provides another example. Carbon contributes four valence electrons, and each oxygen contributes six.
- One carbon atom: 1 × 4 = 4 electrons
- Two oxygen atoms: 2 × 6 = 12 electrons
- Total: 4 + 12 = 16 valence electrons
The Lewis structure of carbon dioxide must account for all 16 electrons, whether they appear in bonds or as lone pairs.
The subscripts in a chemical formula are important because they indicate how many atoms of each element are present. In CO₂, for example, the subscript 2 applies to oxygen only. Carbon contributes four electrons, not eight.
Adjust the electron count for an ion
A polyatomic ion is a group of bonded atoms that carries an overall electric charge. When calculating its valence electrons, start by adding the contributions from all the atoms, just as you would for a neutral molecule. Then adjust the total according to the ion’s charge.
For a negatively charged ion, add electrons. For a positively charged ion, subtract electrons.
The reason is straightforward: a negative charge indicates an excess of electrons relative to the neutral atoms, while a positive charge indicates a deficit.
Consider the hydroxide ion, OH⁻. Oxygen contributes six valence electrons, and hydrogen contributes one. The negative charge means one additional electron must be included.
- Oxygen: 6 electrons
- Hydrogen: 1 electron
- One additional electron for the −1 charge: 1 electron
- Total: 6 + 1 + 1 = 8 valence electrons
Now consider the ammonium ion, NH₄⁺. Nitrogen contributes five valence electrons, and four hydrogen atoms contribute four more. Because the ion has a positive charge of +1, subtract one electron.
- Nitrogen: 5 electrons
- Four hydrogen atoms: 4 × 1 = 4 electrons
- One electron subtracted for the +1 charge: −1 electron
- Total: 5 + 4 − 1 = 8 valence electrons
Both ions have eight valence electrons to distribute in their Lewis structures, but their bonding arrangements differ. Always account for the charge before drawing bonds and lone pairs.
Use the total to draw a Lewis structure
Once you have calculated the total valence electrons, you can begin constructing the Lewis structure. The electron count determines how many electrons must be represented, while the chemical identities of the atoms and their usual bonding patterns help determine where those electrons belong.
For most simple molecules and polyatomic ions, follow this sequence:
1. Choose the central atom. The central atom is usually the least electronegative element other than hydrogen. Electronegativity describes how strongly an atom attracts shared electrons in a bond. Hydrogen is almost always terminal, meaning it bonds to only one other atom.
2. Connect the atoms with single bonds. Draw one line between each pair of bonded atoms. Each single bond represents two electrons, so subtract two electrons from the total for every single bond drawn.
3. Complete the outer atoms’ electron shells. Add lone pairs to the surrounding atoms until they satisfy the appropriate electron-counting rule. Hydrogen needs two electrons in its first energy level, while many common main-group atoms tend to form structures that give them eight electrons around them.
4. Place remaining electrons on the central atom. After the outer atoms have been assigned their electrons, place any remaining electrons on the central atom as lone pairs.
5. Check the central atom’s electron count. If a central atom that typically follows the octet rule has fewer than eight electrons, consider converting lone pairs on adjacent atoms into additional bonds. Each additional bond represents another shared pair of electrons.
6. Verify the structure. Count every electron represented by the bonds and lone pairs. The sum must match the original total valence-electron count. For an ion, place the structure in brackets and show its overall charge.
These steps provide a reliable starting method, but the octet rule has exceptions. Some molecules have electron-deficient central atoms, odd numbers of electrons, or central atoms that can accommodate more than eight electrons in conventional Lewis structures.
Work through an example with water
Water illustrates how the initial electron count guides the entire drawing process.
The formula H₂O contains two hydrogen atoms and one oxygen atom. The total is eight valence electrons.
First, place oxygen in the center and connect it to each hydrogen with a single bond. The two bonds account for four electrons, leaving four electrons.
Next, place the remaining four electrons on oxygen as two lone pairs. Each hydrogen now has two electrons associated with its bond, satisfying the duet rule for hydrogen. Oxygen has eight electrons around it: four from the two bonding pairs and four from its two lone pairs.
The completed Lewis structure has two single bonds and two lone pairs on oxygen. All eight valence electrons are accounted for, and the atoms satisfy their usual electron-counting rules.
This example shows why counting valence electrons comes before assigning lone pairs. The total establishes the electron budget; the bonding arrangement determines how that budget is distributed.
Understand the octet rule and its exceptions
The octet rule states that many main-group atoms tend to form bonds that give them eight electrons in their valence shell. This arrangement resembles the electron configuration of a noble gas and helps explain common bonding patterns.
Carbon, nitrogen, oxygen, and the halogens frequently follow the octet rule in introductory Lewis structures. Hydrogen is different because its first energy level can hold only two electrons.
However, the octet rule is a useful guideline rather than a universal law. Boron, for example, can form stable compounds in which it has only six electrons around it. Beryllium can also appear with fewer than eight electrons. Molecules containing an odd total number of valence electrons cannot give every atom a complete octet using ordinary paired-electron arrangements.
Some compounds of elements in period 3 and beyond are conventionally represented with more than eight electrons around a central atom. These structures can be useful models, although the actual bonding in such compounds is more complex than a simple picture of localized electron pairs.
When the usual rules do not produce a satisfactory structure, compare possible arrangements using formal charge. Formal charge is an accounting tool that assigns electrons in bonds equally between the bonded atoms and compares each atom’s assigned electron count with that of its neutral form. Structures with smaller formal charges are often preferred, especially when negative formal charges are placed on more electronegative atoms. The total formal charge must equal the overall charge of the molecule or ion.
Avoid common mistakes when counting valence electrons
Several simple errors can lead to incorrect Lewis structures.
Confusing valence electrons with all electrons. Oxygen has eight electrons in a neutral atom, but only six are valence electrons. The remaining two occupy its inner energy level and are not normally included in introductory Lewis structures.
Forgetting to multiply by the number of atoms. A formula such as NH₃ contains three hydrogen atoms, each contributing one electron. Counting hydrogen only once would produce an incorrect total.
Ignoring the charge on an ion. A negative charge means electrons must be added to the neutral-atom total, while a positive charge means electrons must be subtracted.
Counting a bond as one electron. A single covalent bond contains two electrons, even though it is represented by one line. A double bond represents four bonding electrons, and a triple bond represents six.
Counting shared electrons twice. When checking a complete Lewis structure, count each bond’s electrons once, not once for each atom. A line between two atoms represents one shared pair, or two electrons total.
Assuming every atom must have an octet. Hydrogen follows the duet rule, and several other elements can appear in valid structures that do not follow the octet rule. Use the electron total and the chemical context rather than forcing every atom to have eight electrons.
Check your answer before finishing
A reliable Lewis structure should satisfy three checks. First, the total number of electrons represented by all bonds and lone pairs must equal the number calculated from the chemical formula and charge. Second, the electron arrangements should be chemically reasonable, including appropriate octets or recognized exceptions. Third, the formal charges should sum to the overall charge of the species.
For many introductory problems, the most important skill is calculating the total correctly. Identify each element, determine its valence electrons from the periodic table, multiply by the number of atoms, and adjust for any ionic charge. With that total established, you can distribute the electrons into bonds and lone pairs and evaluate whether the resulting Lewis structure is reasonable.
