Lewis Structures
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You should be comfortable counting valence electrons for main-group elements and understand that covalent bonds involve the sharing of electron pairs between atoms.
Why this matters
Lewis structures provide the foundational map of electron distribution in molecules.
Before we can predict molecular geometry, reactivity, or polarity, we must establish where the electrons are located. Lewis structures offer a simple, visual representation of how valence electrons are arranged as bonding pairs and lone pairs. This lecture focuses on the systematic method for constructing these structures and the criteria for selecting the most chemically plausible arrangement when multiple options exist.
The Logic of Electron Counting
Think of a Lewis structure as a budget for electrons. First, we calculate the total "wealth" of the molecule by summing the valence electrons of all constituent atoms. If the molecule is an ion, we adjust this total by adding electrons for negative charges or subtracting them for positive charges. Once we have this total pool, we must distribute it according to the rules of chemical stability. The primary goal is to satisfy the octet rule, which states that main-group atoms generally seek to be surrounded by eight electrons (or two for hydrogen) to achieve a noble gas configuration. We achieve this by forming bonds (shared pairs) and placing remaining electrons as lone pairs on individual atoms. The process is not random; it follows a logical hierarchy where central atoms are bonded first, and outer atoms are satisfied before the central atom, ensuring the most stable electronic configuration is reached.
Systematic Construction Steps
Constructing a Lewis structure follows a strict algorithmic sequence to avoid errors. First, identify the central atom, which is typically the least electronegative element (excluding hydrogen, which is always terminal). Second, calculate the total number of valence electrons using the formula \( V_{\text{total}}=\textstyle\sum V_i\mp q \), where \( q \) is the net charge. Third, draw single bonds between the central atom and all surrounding atoms. Each single bond consumes two electrons from the total pool. Fourth, distribute the remaining electrons as lone pairs on the terminal atoms to complete their octets. Finally, place any leftover electrons on the central atom. If the central atom lacks an octet but has available valence orbitals, multiple bonds (double or triple) may be formed by converting lone pairs from terminal atoms into bonding pairs.
Evaluating Structural Validity
Not all valid electron distributions are equally likely. When multiple Lewis structures can be drawn for a given formula, we must use formal charge to determine the best representation. Formal charge is a bookkeeping device that assigns electrons to atoms based on the assumption that bonding electrons are shared equally. It is calculated using the expression \( FC=V-L-\tfrac{B}{2} \), where \( V \) is the number of valence electrons in the free atom, \( L \) is the number of non-bonding (lone pair) electrons, and \( B \) is the number of bonding electrons. The ideal structure minimises the magnitude of formal charges on all atoms. If formal charges are unavoidable, the negative charge should reside on the more electronegative atom, while the positive charge should reside on the less electronegative atom.
Exceptions to the Octet Rule
While the octet rule is a powerful guideline, several important exceptions exist. Boron and aluminium frequently form compounds with only six electrons around the central atom (electron-deficient species) because they have only three valence electrons to share. Conversely, atoms of Period 3 and below (such as phosphorus, sulfur and chlorine) are often drawn with more than eight electrons, as an "expanded octet". These atoms are large enough to bond to many neighbours. Modern calculations show that d orbitals play little part, and the bonding is better described as partly ionic, multi-centre bonding, but the drawings remain a useful bookkeeping device. For example, in sulfur hexafluoride (SF6), sulfur is surrounded by twelve electrons. Additionally, odd-electron species, such as nitric oxide (NO), cannot satisfy the octet rule for all atoms and are often represented with a radical dot.
The result
Lewis structures and the octet rule
A molecule's total valence electrons can be distributed as bonding pairs (shared between two atoms) and lone pairs (localised on one atom) so that, wherever possible, every main-group atom is surrounded by eight electrons (two, for hydrogen). Where more than one such distribution is possible, the formal charge \(FC=V-L-\tfrac{B}{2}\) (valence electrons \(V\), lone-pair electrons \(L\), bonding electrons \(B\)) on each atom selects the best structure: minimise the magnitude of formal charges, and place any necessary negative formal charge on the more electronegative atom.
A common misconception is that formal charge represents the actual physical charge on an atom. It does not. Formal charge is a hypothetical value derived from an arbitrary division of bonding electrons. The actual partial charges in a molecule are determined by electronegativity differences and are best described by dipole moments, not formal charge. Do not confuse the two; a structure with zero formal charges is not necessarily non-polar. C
Worked example
Problem. Determine the best Lewis structure for the nitrate ion (NO3−). Calculate the formal charge on the nitrogen atom and one oxygen atom in the structure with one double bond.
- Calculate total valence electrons: Nitrogen (5) + 3 × Oxygen (6) + 1 (for negative charge) = 24 electrons.
- Place Nitrogen as the central atom and bond it to three Oxygen atoms using single bonds. This uses 6 electrons (3 bonds × 2 e−).
- Distribute remaining 18 electrons as lone pairs on the Oxygen atoms to complete their octets (3 × 6 e− = 18 e−). Nitrogen currently has only 6 electrons.
- Form a double bond by converting one lone pair from an Oxygen atom into a bonding pair. Nitrogen now has an octet (8 electrons).
- Calculate Formal Charge (FC) for Nitrogen: \( V=5 \), \( L=0 \), \( B=8 \). \( FC = 5 - 0 - \frac{8}{2} = +1 \).
- Calculate Formal Charge for the double-bonded Oxygen: \( V=6 \), \( L=4 \), \( B=4 \). \( FC = 6 - 4 - \frac{4}{2} = 0 \).
- Calculate Formal Charge for the single-bonded Oxygen: \( V=6 \), \( L=6 \), \( B=2 \). \( FC = 6 - 6 - \frac{2}{2} = -1 \).
Units check. Formal charge is a dimensionless integer representing electron count difference.
Three quick tests
- How many valence electrons are in the sulfite ion (SO32−)?
Answer
26 electrons (S: 6 + 3×O: 18 + 2 charge: 2). - Why is the structure H–O–H preferred over H–H–O for water?
Answer
Hydrogen can only form one bond (duet rule), so it must be terminal; oxygen is more electronegative and central. - What is the formal charge on a carbon atom in methane (CH4)?
Answer
0 (V=4, L=0, B=8; 4 - 0 - 4 = 0).
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Every lecture and every result in this unit, with prerequisites and threads.
Lecture notes drafted with a local model and checked line by line by the editor; every number in the worked example was recomputed.