AP Chemistry 2.7: VSEPR and hybridization

• This is another huge Unit 2 topic. Khan Academy currently covers molecules with 2 through 6 electron clouds plus molecular polarity .

• VSEPR

  • stands for:

  • Valence Shell Electron Pair Repulsion

• The idea is simple:

  • electron groups repel one another.

  • Therefore, electron groups arrange themselves as far apart as possible.

  • That determines molecular geometry.

• Each of these counts as one electron domain/group around a central atom:

• single bond

• double bond

• triple bond

  • lone pair

  • Important:

• A triple bond counts as one electron group , not three.

• These are different.

• Electron geometry

  • considers:

  • bonds + lone pairs

• Molecular geometry

  • considers:

  • positions of atoms only

• Lone pairs affect the shape but are not included when naming molecular geometry.

• Electron geometry:

  • linear

• Ideal bond angle:

  • 180°

  • Example:

  • CO₂

  • O=C=O

• Carbon has:

  • 2 electron groups

  • 0 lone pairs

• Molecular geometry:

  • linear

• Electron geometry:

  • trigonal planar

• Ideal angles:

  • 120°

• Khan Academy states that three electron clouds adopt trigonal-planar electron geometry to minimize repulsion.

• Three bonds

  • zero lone pairs

  • Example:

  • BF₃

• Molecular geometry:

  • trigonal planar

• Bond angles:

  • ≈ 120°

• Two bonds

  • one lone pair

• Molecular geometry:

  • bent

  • Because the electron geometry is trigonal planar but one position is occupied by a lone pair.

• Electron geometry:

  • tetrahedral

• Ideal angle:

  • 109.5°

• Four bonds

  • zero lone pairs

  • Example:

  • CH₄

• Molecular geometry:

  • tetrahedral

• Angle:

  • 109.5°

• Three bonds

  • one lone pair

  • Example:

  • NH₃

• Molecular geometry:

  • trigonal pyramidal

• Bond angle:

  • slightly less than 109.5°

  • approximately:

  • 107°

• Two bonds

  • two lone pairs

  • Example:

  • H₂O

• Molecular geometry:

  • bent

• Bond angle:

  • approximately:

  • 104.5°

• Repulsion strengths generally follow:

  • lone pair–lone pair > lone pair–bond pair > bond pair–bond pair

  • Why?

• Lone-pair electron density is concentrated closer to the central atom and occupies more space.

  • Therefore lone pairs push bonding pairs closer together.

• Electron geometry:

  • trigonal bipyramidal

• There are two types of positions:

  • Equatorial

  • three positions around the center

  • 120° apart

  • Axial

  • two positions above and below

  • 90° from the equatorial plane

• Example:

  • PCl₅

• Molecular geometry:

  • trigonal bipyramidal

• Angles:

  • 90°, 120°, and 180°

• Lone pairs prefer:

  • equatorial positions

  • because equatorial positions experience fewer 90° interactions.

  • This reduces repulsion.

• Five electron groups:

  • 4 bonds

  • 1 lone pair

• Molecular geometry:

  • seesaw

• 3 bonds

  • 2 lone pairs

• Molecular geometry:

  • T-shaped

• 2 bonds

  • 3 lone pairs

• Molecular geometry:

  • linear

• The three lone pairs occupy equatorial positions.

• Electron geometry:

  • octahedral

• Ideal angles:

• 90° and 180°.

• Khan Academy's current Unit 2 material identifies octahedral, square-pyramidal, and square-planar arrangements for six electron groups.

• Example:

  • SF₆

• Molecular geometry:

  • octahedral

• 5 bonds

  • 1 lone pair

• Molecular geometry:

  • square pyramidal

• 4 bonds

  • 2 lone pairs

• The two lone pairs occupy opposite positions.

• Molecular geometry:

  • square planar

  • Example:

  • XeF₄

  • VSEPR Master Table

• This table is worth memorizing.

• Khan Academy includes molecular polarity inside its VSEPR section.

• Do not confuse:

  • bond polarity

  • with:

  • molecular polarity

• A molecule can contain polar bonds but still be:

  • nonpolar overall

• A polar bond has a dipole .

• Electron density is pulled toward the more electronegative atom.

  • Think of each bond dipole as a vector with:

  • direction

  • magnitude

• To determine whether the whole molecule is polar:

• Determine whether the bonds are polar.

• Determine the molecular shape.

• Decide whether bond dipoles cancel.

• CO₂:

  • O=C=O

  • Each C=O bond is polar.

• But the molecule is:

  • linear

• The two identical dipoles point in opposite directions.

  • They cancel.

  • Therefore:

  • CO₂ is nonpolar

  • even though:

  • C=O bonds are polar.

• O—H bonds are polar.

• Water is:

  • bent

• The dipoles are not directly opposite.

  • They do not cancel.

  • Therefore:

• H₂O is polar.

• B—F bonds are polar.

• BF₃ is trigonal planar and symmetrical.

  • Its three equal bond dipoles cancel.

  • Therefore:

• BF₃ is nonpolar.

• N—H bonds are polar.

• NH₃ is trigonal pyramidal because nitrogen has a lone pair.

• The molecule is not symmetrical enough for the bond dipoles to cancel.

  • Therefore:

• NH₃ is polar.

• If:

  • outer atoms are identical

  • geometry is highly symmetric

  • dipoles often cancel.

  • Examples:

  • CO₂ → nonpolar

  • BF₃ → nonpolar

  • CH₄ → nonpolar

  • SF₆ → nonpolar

• But never rely only on memorizing this shortcut.

• Always examine:

  • bond polarity + geometry

• Khan Academy's current VSEPR practice directly asks students to use Lewis structures and geometry to compare molecular polarity.

• Khan Academy's current hybridization section covers sp³, steric number, sp², sp, and identifying hybridization in molecules .

• Simple atomic s and p orbitals do not always explain observed molecular geometries very well.

• The hybridization model says atomic orbitals can combine mathematically to form new orbitals oriented for bonding.

  • These are called:

  • hybrid orbitals

• For AP Chemistry Unit 2:

  • sp

  • sp²

  • sp³

  • These correspond closely to electron-group geometry.

• Steric number tells you the number of electron groups around an atom.

• A convenient AP rule:

  • steric number = number of sigma bonds + number of lone pairs

  • Important:

• Double and triple bonds each contain only:

  • one sigma bond

  • so they each count as one electron group.

• If an atom has four electron groups:

  • sp³ hybridized

• One s orbital combines with:

  • three p orbitals

  • giving:

  • four sp³ hybrid orbitals

• Khan Academy describes sp³ as combining one s and three p orbitals, producing four equivalent hybrid orbitals.

  • 8.5 sp³ Geometry

• Electron geometry:

  • tetrahedral

  • Examples:

  • CH₄

  • NH₃ central N

  • H₂O central O

• All three central atoms can be treated as:

  • sp³

  • because all have:

• 4 electron groups.

• Carbon in CH₄ has:

  • 4 C—H sigma bonds

  • 0 lone pairs

• Steric number:

  • 4

  • Therefore:

  • sp³

• Geometry:

  • tetrahedral

• Oxygen has:

  • 2 O—H sigma bonds

  • 2 lone pairs

• Steric number:

  • 2 + 2 = 4

  • Therefore:

  • sp³

• Even though the molecular shape is bent.

  • This is important:

  • hybridization follows electron groups, not just the visible molecular shape.

• Three electron groups:

  • sp² hybridization

• One s orbital combines with:

  • two p orbitals

  • forming:

  • three sp² orbitals

• One unhybridized p orbital remains.

  • 8.9 sp² Geometry

• Electron geometry:

  • trigonal planar

• Ideal angle:

  • 120°

• A carbon involved in a typical carbon-carbon double bond is often:

  • sp² hybridized

• In:

  • H₂C=CH₂

  • each carbon has:

  • two C—H sigma bonds

  • one C—C sigma bond

  • Total:

  • 3 sigma-bond groups

  • Steric number = 3

  • Therefore:

  • each carbon is sp²

• Two electron groups:

  • sp hybridization

• One s orbital combines with:

  • one p orbital.

  • This produces:

  • two sp orbitals

• Two unhybridized p orbitals remain.

  • 8.12 sp Geometry

• Geometry:

  • linear

• Bond angle:

  • 180°

• Atoms involved in typical triple bonds are often:

  • sp hybridized

• HC≡CH

• Each carbon has:

  • one C—H sigma bond

  • one C—C sigma bond

• Total electron groups:

  • 2

  • Therefore:

  • sp hybridized

  • Hybridization Table

• Memorize this.

• A sigma bond (σ) forms through direct, head-on orbital overlap.

• All single bonds are:

  • sigma bonds

  • For example:

  • C—C

  • contains:

  • 1 σ bond

• A pi bond (π) forms through sideways overlap of parallel p orbitals.

• Pi bonds occur in:

  • double and triple bonds.

• A single bond contains:

  • 1 σ

  • So:

  • C—C

  • = 1 sigma

• A double bond contains:

  • 1 σ + 1 π

  • So:

  • C=C

  • does not contain two sigma bonds.

• A triple bond contains:

  • 1 σ + 2 π

  • So:

  • C≡C

  • = one sigma + two pi.

  • This is extremely important.

• You can determine hybridization quickly.

• Step 1

• Draw or inspect Lewis structure.

• Step 2

• Pick the atom.

• Step 3

• Count:

  • sigma bonds + lone pairs.

• Step 4

• 2 → sp

  • 3 → sp²

  • 4 → sp³

• Khan Academy's current practice explicitly asks students to determine hybridization and identify the orbitals that overlap to make sigma bonds.

• These connections are more important than memorizing random facts.

Connection 1: Electronegativity → Bond Type

• Small difference:

  • → mostly covalent

• Large difference:

  • → increasing ionic character

Connection 2: Coulomb's Law → Ionic Strength

• Larger charges:

  • → stronger attraction

• Smaller ions:

  • → shorter distance

  • → stronger attraction

Connection 3: Bond Order → Strength and Length

• Higher bond order:

  • → more shared electron density

  • → stronger attraction

  • → stronger bond

  • → shorter bond

  • Therefore:

  • triple > double > single in strength

  • and:

  • triple < double < single in length

  • for comparable atoms.

Connection 4: Lewis Diagram → VSEPR

• Lewis diagram tells you:

  • number of bonds

  • number of lone pairs

  • Those give:

  • electron groups

  • which determine:

• molecular geometry

Connection 5: Molecular Geometry → Polarity

• Lewis structure

  • ↓

  • VSEPR geometry

  • ↓

  • bond dipole directions

  • ↓

  • whether dipoles cancel

  • ↓

  • polar or nonpolar molecule

Connection 6: VSEPR → Hybridization

• 2 electron groups:

  • linear

  • → sp

• 3 electron groups:

  • trigonal planar

  • → sp²

• 4 electron groups:

  • tetrahedral

  • → sp³

Connection 7: Hybridization → Sigma/Pi Bonding

• Single:

  • σ

• Double:

  • σ + π

• Triple:

  • σ + 2π

Unit 2 Master Vocabulary

• Ionic bond — electrostatic attraction between oppositely charged ions.

• Covalent bond — bond involving shared electrons.

• Metallic bond — attraction involving metal cores and delocalized electrons.

• Electronegativity — ability of an atom to attract shared electrons.

• Bond polarity — unequal distribution of bonding electrons.

• Bond order — number of bonding pairs between atoms.

• Bond length — equilibrium distance between bonded nuclei.

• Bond energy — energy required to break a bond.

• Crystal lattice — repeating arrangement of particles in a solid.

• Lattice energy — energy associated with forming/separating an ionic lattice.

• Alloy — mixture containing multiple elements with at least one metal.

• Substitutional alloy — atoms replace other metal atoms in lattice positions.

• Interstitial alloy — smaller atoms occupy gaps in a metal lattice.

• Lewis diagram — representation of valence electrons and bonds.

• Lone pair — nonbonding pair of electrons.

• Octet rule — tendency of many atoms to reach eight valence electrons.

• Resonance — multiple valid Lewis representations differing only in electron placement.

• Resonance hybrid — actual delocalized electronic structure represented by resonance forms.

• Formal charge — bookkeeping charge assigned to an atom in a Lewis structure.

• VSEPR — model predicting molecular shape from electron-group repulsions.

• Electron geometry — arrangement of all electron groups.

• Molecular geometry — arrangement of bonded atoms.

• Bond dipole — separation of partial charge across a polar bond.

• Molecular polarity — overall uneven charge distribution in a molecule.

• Hybridization — model describing mixing of atomic orbitals.

• Steric number — number of sigma bonds plus lone pairs around an atom.

• Sigma bond — bond formed by head-on orbital overlap.

• Pi bond — bond formed by sideways p-orbital overlap.

Unit 2 Things You Need to Memorize

• Bond order

  • Single = 1

  • Double = 2

  • Triple = 3

  • Bond strength

  • triple > double > single

  • Bond length

  • single > double > triple

  • Sigma and pi

  • Single = 1σ

  • Double = 1σ + 1π

  • Triple = 1σ + 2π

  • Hybridization

  • 2 groups = sp

  • 3 groups = sp²

  • 4 groups = sp³

  • Basic geometries

  • 2 groups → linear

  • 3 → trigonal planar

  • 4 → tetrahedral

  • 5 → trigonal bipyramidal

  • 6 → octahedral

  • Formal charge

  • FC = valence electrons − nonbonding electrons − number of bonds

Most Common Unit 2 Exam Mistakes

• Assuming every metal/nonmetal bond is perfectly ionic.

• Forgetting that bond character exists on a continuum.

• Saying bond breaking releases energy. Breaking bonds requires energy.

• Thinking an ionic crystal consists of separate molecules.

• Forgetting that larger ionic charges produce stronger attractions.

• Forgetting that smaller ion distance produces stronger attraction.

• Mixing up substitutional and interstitial alloys.

• Counting the wrong number of valence electrons in Lewis structures.

• Forgetting to add electrons for negative ions .

• Forgetting to subtract electrons for positive ions .

• Putting hydrogen in the center of a Lewis structure.

• Giving hydrogen eight electrons instead of two.

• Giving C, N, O, or F an expanded octet.

• Moving atoms when creating resonance structures. Only electrons move.

• Thinking the real molecule rapidly jumps between resonance structures.

• Forgetting that equivalent resonance can produce fractional average bond orders.

• Forgetting to check that formal charges add to the overall molecular/ionic charge.

• Counting a double bond as two VSEPR electron groups. It counts as one.

• Counting a triple bond as three VSEPR groups. It counts as one.

• Confusing electron geometry with molecular geometry.

• Ignoring lone pairs when finding electron geometry.

• Assuming every molecule with polar bonds is polar.

• Forgetting that symmetric bond dipoles can cancel.

• Thinking bent H₂O is nonpolar because it contains two identical bonds.

• Thinking NH₃ is tetrahedral in molecular shape. Its electron geometry is tetrahedral, but molecular geometry is trigonal pyramidal .

• Thinking H₂O's molecular geometry is tetrahedral. It is bent .

• Forgetting that lone pairs repel more strongly than bonding pairs.

• Forgetting that lone pairs prefer equatorial positions in trigonal-bipyramidal arrangements.

• Counting double/triple bonds multiple times when calculating steric number.

• Thinking a double bond contains two sigma bonds. It contains one sigma and one pi .

• Thinking a triple bond contains three sigma bonds. It contains one sigma and two pi .

• Determining hybridization from the number of atoms instead of the number of electron groups.

The Unit 2 Problem-Solving Chain

• For a molecule question, this order will solve a huge percentage of Unit 2 problems:

  • 1. Count valence electrons

  • ↓

  • 2. Draw Lewis structure

  • ↓

• 3. Check octets and formal charges

  • ↓

  • 4. Check for resonance

  • ↓

  • 5. Count electron groups

  • ↓

  • 6. Find electron geometry

  • ↓

  • 7. Find molecular geometry

  • ↓

  • 8. Determine bond angles

  • ↓

  • 9. Determine molecular polarity

  • ↓

  • 10. Determine hybridization

  • ↓

• 11. Count σ and π bonds

• If a student can confidently work through that chain, they have most of the hardest skills in AP Chemistry Unit 2 covered.

  • These notes follow all eight topics in Khan Academy’s current Unit 2 course, including its current emphasis on bond character, alloys, Lewis structures, resonance/formal charge, molecular polarity, and hybridization.