Chemical Bonding & Molecular Geometry Cheat Sheet

A printable reference covering Lewis structures, formal charge, electronegativity, VSEPR shapes, bond angles, polarity, and intermolecular forces for grades 10-11.

Download PNG

Chemical bonding explains how atoms connect to form compounds and why different substances have different properties. This cheat sheet helps students organize Lewis structures, bond types, formal charge, and molecular shapes in one clear reference. It is especially useful when predicting geometry, polarity, and intermolecular forces from a chemical formula. The core ideas are valence electrons, electron sharing or transfer, and electron pair repulsion. Lewis structures show bonding pairs and lone pairs, while formal charge helps choose the best structure. VSEPR theory uses the number of electron groups around a central atom to predict shape, bond angles, and polarity.

Key Facts

Bond polarity increases as the electronegativity difference increases, using $\Delta EN = |EN_A - EN_B|$ .
For Lewis structures, the total valence electron count equals the sum of valence electrons from all atoms, plus electrons for anions and minus electrons for cations.
Formal charge is calculated with $FC = V - N - \frac{B}{2}$ , where $V$ is valence electrons, $N$ is nonbonding electrons, and $B$ is bonding electrons.
The best Lewis structure usually has the smallest formal charges, negative formal charge on the more electronegative atom, and complete octets when possible.
The steric number is $SN = \text{number of bonded atoms} + \text{number of lone pairs}$ around the central atom.
VSEPR predicts electron geometries of linear for $SN = 2$ , trigonal planar for $SN = 3$ , tetrahedral for $SN = 4$ , trigonal bipyramidal for $SN = 5$ , and octahedral for $SN = 6$ .
Common ideal bond angles are $180^\circ$ for linear, $120^\circ$ for trigonal planar, $109.5^\circ$ for tetrahedral, $90^\circ$ and $120^\circ$ for trigonal bipyramidal, and $90^\circ$ for octahedral.
A molecule is nonpolar when its bond dipoles cancel, which means the vector sum is $\sum \vec{\mu} = 0$ .

Vocabulary

Valence electron: A valence electron is an outer-shell electron that can participate in chemical bonding.
Electronegativity: Electronegativity is an atom's ability to attract shared electrons in a chemical bond.
Lewis structure: A Lewis structure is a diagram that shows valence electrons as dots and covalent bonds as lines.
Formal charge: Formal charge is the apparent charge on an atom in a Lewis structure, found using $FC = V - N - \frac{B}{2}$ .
VSEPR theory: VSEPR theory predicts molecular shape by assuming electron groups repel each other and spread out as far apart as possible.
Dipole moment: A dipole moment is a measure of charge separation in a bond or molecule, represented by the vector $\vec{\mu}$ .

Common Mistakes to Avoid

Forgetting to adjust electron count for ions is wrong because an anion gains electrons and a cation loses electrons before the Lewis structure is drawn.
Counting a double bond as two electron groups in VSEPR is wrong because any single, double, or triple bond counts as one electron group around the central atom.
Ignoring lone pairs when predicting shape is wrong because lone pairs repel bonding pairs more strongly and can change both molecular geometry and bond angles.
Assuming every polar bond makes a polar molecule is wrong because symmetrical molecules can have bond dipoles that cancel, giving $\sum \vec{\mu} = 0$ .
Choosing the Lewis structure with more formal charge is wrong when a structure with smaller formal charges and complete octets is available.

Practice Questions

1 Calculate the formal charge on oxygen in a Lewis structure where oxygen has $6$ valence electrons, $4$ nonbonding electrons, and $4$ bonding electrons.
2 For $\mathrm{NH_3}$ , find the steric number of nitrogen and predict the molecular geometry and approximate bond angle.
3 Use $\Delta EN = |EN_A - EN_B|$ to compare bond polarity for $\mathrm{H-Cl}$ with $EN_H = 2.20$ and $EN_{Cl} = 3.16$ .
4 Explain why $\mathrm{CO_2}$ is nonpolar even though each $\mathrm{C=O}$ bond is polar.