DPST1031 Chap.4 Lewis Structures, Formal Charge and VSEPR
Lewis Structures, Formal Charge and VSEPR
Lewis diagrams turn a formula into a picture of shared and unshared valence electrons, and VSEPR theory turns that picture into a three-dimensional shape. This chapter sets out a five-step drawing routine, the three kinds of octet exception, and the formal charge calculation used to choose the most realistic structure.
It then counts electron domains to give electron geometry, removes lone pairs to name the molecular shape, and predicts bond angles for two to six domains, including seesaw, T-shaped and square planar species. Resonance and delocalisation, the Mastery extension, close the chapter.
What this chapter covers
- 01
Counting valence electrons for molecules and ions
- 02
Lewis diagrams of organic functional groups
- 03
Octet exceptions and expanded valence shells
- 04
Formal charge and the preferred structure
- 05
Resonance and delocalisation
- 06
Electron domains and electron geometry
- 07
Molecular shape and bond angles
- 08
Molecular polarity from shape
Worked example · free
Shape and bond angle of ammonia
- +1Count: 5 + 3 × 1 = 8 valence electrons.
- +1Three N−H bonds use 6 electrons, leaving one lone pair on nitrogen.
- +1Four domains give a tetrahedral electron geometry; with one lone pair the shape is trigonal pyramidal.
- +1The lone pair compresses the angle from 109.5° to about 107°.
Key terms
- Lewis diagram
- A drawing of a species that shows every valence electron as a bonding pair or a lone pair.
- Lone pair
- A pair of valence electrons that belongs to one atom and is not shared in a bond.
- Octet exception
- A stable species whose central atom has fewer or more than eight valence electrons.
- Resonance structures
- Two or more Lewis diagrams with the same atom positions that differ only in where electrons are placed.
- Electron domain geometry
- The arrangement of all electron domains, bonds and lone pairs, around a central atom.
- Molecular shape
- The arrangement of the atoms alone, ignoring lone pairs, around a central atom.
- Bond angle
- The angle between two bonds that share a central atom.
Lewis Structures, Formal Charge and VSEPR FAQ
How do I decide which Lewis structure is best?
Calculate the formal charge of every atom in each candidate. Prefer the structure with charges closest to zero, and if a negative formal charge is unavoidable, place it on the most electronegative atom. The charges must always add up to the overall charge of the species, which is a quick arithmetic check.
Why can sulfur and phosphorus have more than eight electrons?
Atoms in period 3 and beyond are large enough, and have available orbitals, to accommodate more than eight electrons around them. That is why PCl₅ and SF₆ are stable. Carbon, nitrogen, oxygen and fluorine are in period 2 and never exceed an octet in a correct Lewis diagram.
What is the difference between electron geometry and molecular shape?
Electron geometry describes where all the electron domains point, including lone pairs. Molecular shape describes only where the atoms are. Ammonia has a tetrahedral electron geometry but a trigonal pyramidal shape, because one of its four domains is a lone pair.
Why do lone pairs reduce bond angles?
A lone pair is held by only one nucleus, so its electron density spreads out closer to the central atom and repels neighbouring bonding pairs more strongly than another bond would. The bonds are pushed together, so angles fall from 109.5° in methane to about 107° in ammonia and 104.5° in water.
How do I count electron domains around a central atom?
Count each lone pair on the central atom as one domain and each bond to another atom as one domain, whether it is single, double or triple. Carbon dioxide therefore has two domains and is linear, while the sulfite ion has four domains, three bonds and one lone pair, and is trigonal pyramidal.
Exam move
Practise the five-step Lewis routine on twenty species until the electron count is automatic, including at least five ions and five organic molecules with oxygen or nitrogen. For each, assign formal charges and check that they sum to the overall charge. Then add VSEPR: write the domain count, the electron geometry, the molecular shape and the ideal angle in a four-column table.
Pay particular attention to five- and six-domain species, where lone pairs take equatorial or opposite positions. Finish with resonance: draw nitrate, carbonate and ozone, and state the bond order you expect from the number of equivalent structures.
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