A structural formula shows how atoms in a molecule are connected to each other and, depending on the style, how electrons are arranged. There are several types, from fully detailed Lewis structures to the shorthand skeletal formulas used in organic chemistry. Each follows specific rules, and once you learn those rules, drawing them becomes a repeatable process rather than guesswork.
Three Types of Structural Formulas
Chemists use three main formats, each offering a different level of detail. Choosing the right one depends on what you need to communicate.
- Lewis (Kekulé) structures show every atom, every bond (as a line), and all lone pair electrons. These are the most detailed and the best starting point for learning.
- Condensed formulas write out atoms as text without drawing bonds. For example, propane becomes CH₃CH₂CH₃. The order of the symbols tells you which atoms connect to which. These save space but can get hard to read for complex molecules.
- Skeletal (bond-line) formulas strip away carbon and hydrogen atoms entirely. Carbon is assumed at every bend and endpoint of a zigzag line, and hydrogens are filled in mentally to give each carbon four bonds. Atoms other than carbon and hydrogen (called heteroatoms) are always written out. This is the standard in organic chemistry because it makes the overall shape and functional groups easy to see at a glance.
A molecule like retinol drawn as a full Lewis structure would be dense and hard to read. The same molecule as a skeletal formula clearly shows its ring, its chain of double bonds, and its OH group. For small inorganic molecules, though, Lewis structures are usually the better choice.
Drawing a Lewis Structure Step by Step
Lewis structures are the foundation. If you can draw one correctly, the other formats are simplifications of the same information. Follow these five steps in order.
Step 1: Count Total Valence Electrons
Add up the valence electrons for every atom in the molecule. Valence electrons correspond to an element’s group number on the periodic table: carbon (group 4) has 4, nitrogen (group 5) has 5, oxygen (group 6) has 6, and so on. For a polyatomic ion, add one electron for each negative charge or subtract one for each positive charge. This total is the electron budget you have to work with, and every electron must be placed somewhere in the final structure.
Step 2: Arrange Atoms and Draw Single Bonds
Place the least electronegative atom in the center and arrange the other atoms around it. Hydrogen and fluorine can only form one bond, so they always go on the outside. Connect each outer atom to the central atom with a single line, which represents a shared pair of two electrons. Each bond you draw uses two electrons from your total.
Step 3: Distribute Remaining Electrons to Outer Atoms
After drawing the single bonds, place the leftover electrons as lone pairs (dots) on the outer atoms until each has eight electrons in its valence shell. Hydrogen is the exception: it only needs two. Start with the most electronegative atoms first.
Step 4: Place Leftover Electrons on the Central Atom
If you still have unassigned electrons after completing the outer atoms’ octets, put them on the central atom as lone pairs.
Step 5: Form Multiple Bonds if Needed
Check whether the central atom has a full octet. If it doesn’t, move a lone pair from an outer atom to form a double or triple bond with the central atom. For example, in CO₂, each oxygen shares two pairs with carbon, creating two double bonds so that every atom reaches eight electrons.
The Octet Rule and Its Limits
Most atoms in a structural formula should end up with eight electrons in their valence shell (four pairs). This is the octet rule, and it holds reliably for carbon, nitrogen, oxygen, and fluorine. Hydrogen is satisfied with just two electrons.
Some elements in the third row of the periodic table and below can hold more than eight electrons. Phosphorus, for instance, forms five bonds in compounds like PCl₅, and sulfur forms six bonds in SF₆. These are called hypervalent molecules. Boron goes the other direction: in BF₃, it has only six electrons around it and is stable that way. When you draw structures involving these elements, don’t force the octet rule if the chemistry doesn’t support it.
Checking Your Work With Formal Charge
Formal charge tells you whether the electrons in your structure are distributed realistically. The calculation for any atom is: valence electrons of the isolated atom, minus the number of bonds to that atom, minus the number of nonbonding (lone pair) electrons on that atom.
A carbon with four bonds and no lone pairs has a formal charge of zero (4 − 4 − 0 = 0), which is what you’d expect. If that same carbon had only three bonds and a lone pair, the formal charge would be −1 (4 − 3 − 2 = −1). The best Lewis structure is generally the one where formal charges are as close to zero as possible, negative charges sit on the more electronegative atoms, and no atom carries an unreasonably large charge. If your structure gives a +2 on oxygen, something is wrong.
Drawing Skeletal (Bond-Line) Formulas
Skeletal formulas follow three rules that make them fast to draw and easy to read. First, carbon atoms are not written out. Instead, every intersection of two lines and every endpoint of a line represents a carbon. Second, hydrogens bonded to carbon are omitted. Since carbon always forms exactly four bonds, you mentally fill in however many hydrogens are needed to reach four. A carbon at a bend with two lines coming off it has two implicit hydrogens; a carbon at an endpoint has three. Third, every atom that is not carbon or hydrogen must be written explicitly. Oxygen, nitrogen, sulfur, halogens, and any other heteroatom always appears as its letter symbol.
The carbon backbone is drawn as a zigzag line, with each zig or zag representing a carbon-carbon bond. This zigzag shape isn’t arbitrary: it roughly reflects the 109.5° bond angles of a tetrahedral carbon. Double bonds are shown as two parallel lines, triple bonds as three. Functional groups like OH, NH₂, and COOH are written at the appropriate position along the chain.
Groups like CH₃, OH, and NH₂ are usually written with carbon, oxygen, or nitrogen first. But if the group is on the left side of a chain, the order is sometimes reversed to H₃C, HO, or H₂N so that the bonding atom faces the right direction. This keeps the connectivity clear.
Common Shorthand for Functional Groups
In more advanced structural formulas, you’ll see abbreviations that stand in for common groups. “Me” means a methyl group (CH₃), “Et” means ethyl (CH₃CH₂), “Pr” means propyl (CH₃CH₂CH₂), and “Ph” means a phenyl ring (C₆H₅). The generic “R” stands for any alkyl or aryl group and is used when the specific identity of that part of the molecule doesn’t matter for the discussion. These abbreviations let you focus on the chemically interesting part of a molecule without cluttering the drawing.
Showing 3D Shape on a Flat Page
Structural formulas are two-dimensional, but molecules exist in three dimensions. Chemists handle this with wedge-and-dash notation. A solid wedge represents a bond coming out of the page toward you. A dashed (or hatched) wedge represents a bond going behind the page, away from you. A plain line represents a bond in the plane of the page. This notation is essential when a molecule’s 3D arrangement matters, particularly at carbon atoms bonded to four different groups, where the spatial arrangement determines the molecule’s biological activity or chemical behavior.
When drawing a tetrahedral carbon this way, place two bonds as plain lines roughly in the plane, one as a wedge, and one as a dash. The angle between bonds at a tetrahedral center is 109.5°. For a trigonal planar center (three groups, like a carbon in a double bond), the bonds are separated by 120°. Linear arrangements have 180° between bonds.
Drawing Resonance Structures
Some molecules can’t be accurately represented by a single structural formula because their electrons are spread out (delocalized) across multiple bonds. In these cases, you draw two or more resonance structures connected by a double-headed arrow. This arrow does not mean the molecule flips between forms. It means the real structure is a blend of all the forms shown.
To generate a resonance structure from an existing one, use curved arrows to show electron movement. A curved arrow with two barbs represents two electrons moving. The tail of the arrow starts where the electrons are (a lone pair or a pi bond), and the head points to where they’re going. Only three types of electron movement are valid in resonance: a lone pair forming a new pi bond to an adjacent atom, a pi bond shifting to become a new pi bond on the next atom over, or a pi bond breaking to become a lone pair on an adjacent atom. No atoms move, only electrons.
After drawing each resonance form, verify that no atom exceeds its octet (for second-row elements), that all formal charges are correct, and that the total electron count hasn’t changed. The resonance form where every atom has a full octet and formal charges are minimized is typically the major contributor to the real structure.