The counting rule
For any atom in a molecule, count the total number of σ bonds plus lone pairs on that atom. Don't count π bonds — they don't use hybrid orbitals. Two groups → sp. Three → sp². Four → sp³. (Five would be sp³d, six sp³d², used for hypervalent atoms like P and S in extended-octet structures.)
sp (linear)
Two groups, 180° apart. Examples: the carbons of an alkyne C≡C (each has one σ to C, one σ to H or R, plus two π bonds we don't count). Carbon in CO₂. Bond angles are linear, and the unhybridized p-orbitals carry the π bonds.
sp² (trigonal planar)
Three groups, 120° apart, all in one plane. Examples: the carbons in an alkene C=C, the carbon of a carbonyl C=O, every carbon in benzene, the boron in BF₃. The remaining p-orbital is perpendicular to the plane and (for C=C, C=O, aromatic rings) carries the π bond.
sp³ (tetrahedral)
Four groups, 109.5° apart, pointing toward the corners of a tetrahedron. Examples: every saturated carbon (alkanes, alcohols, ethers, amines on the sp³ N), water's oxygen (2 bonds + 2 lone pairs = 4), ammonia's nitrogen (3 bonds + 1 lone pair = 4). Lone pairs squeeze the bond angles a bit — water is 104.5° instead of the ideal 109.5° because lone pairs push harder than bonded pairs.
Draw this on the whiteboard
Open the OChem Board whiteboard — benzene rings, wedge/dash bonds, and a clickable periodic table built in. No account needed.