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 .
- 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 , where is valence electrons, is nonbonding electrons, and 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 around the central atom.
- VSEPR predicts electron geometries of linear for , trigonal planar for , tetrahedral for , trigonal bipyramidal for , and octahedral for .
- Common ideal bond angles are for linear, for trigonal planar, for tetrahedral, and for trigonal bipyramidal, and for octahedral.
- A molecule is nonpolar when its bond dipoles cancel, which means the vector sum is .
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 .
- 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 .
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 .
- 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 valence electrons, nonbonding electrons, and bonding electrons.
- 2 For , find the steric number of nitrogen and predict the molecular geometry and approximate bond angle.
- 3 Use to compare bond polarity for with and .
- 4 Explain why is nonpolar even though each bond is polar.
Understanding Chemical Bonding & Molecular Geometry
A reliable way to solve a structure problem is to work in a fixed order. Start by choosing the central atom. It is usually the least electronegative atom, but hydrogen is never central.
Connect the atoms with single bonds first, then place remaining electrons around outer atoms before placing any on the central atom. If the central atom lacks an octet, turn a lone pair from a nearby atom into a double or triple bond.
Check the electron total at the end. This catches many mistakes, especially missed lone pairs and incorrect charges on ions.
Formal charge is a bookkeeping tool, not the actual charge sitting on one atom. It helps compare possible drawings of the same substance. Some molecules need more than one valid Lewis structure.
These are called resonance forms. In nitrate ion, for example, the double bond can be drawn in different positions. The real ion is not switching between separate drawings.
Its electrons are spread across the atoms, so the nitrogen oxygen bonds have equal average lengths. This matters because resonance often makes a particle more stable and changes how strongly it reacts.
Electron geometry and molecular shape are related but not identical. Electron geometry counts every region of electron density around the central atom. Molecular shape describes the positions of atoms only.
A central atom with four electron groups has a tetrahedral electron arrangement. If one group is a lone pair, the visible shape becomes trigonal pyramidal. If two groups are lone pairs, it becomes bent.
Lone pairs take up more space than bonding pairs because their electron density stays closer to the central atom. They push bonds closer together, so real bond angles are often smaller than ideal values. Water is bent for this reason, while carbon dioxide is linear because its central carbon has no lone pairs.
Polarity depends on the whole three dimensional structure, not just whether individual bonds are polar. Carbon dioxide has polar carbon oxygen bonds, yet its straight symmetrical shape causes the pulls to cancel. Water has polar oxygen hydrogen bonds arranged in a bent shape, so it has a net pull toward oxygen.
This difference affects everyday properties. Polar water dissolves many ionic compounds and polar substances. Nonpolar oils do not mix well with water because their attractions are different.
Intermolecular forces explain boiling points, melting points, viscosity, and surface tension. Hydrogen bonding is especially important in water, alcohols, DNA, and proteins. When studying, draw shapes carefully, mark lone pairs, then use symmetry to judge polarity before predicting the forces between molecules.