AP®ChemistryCollege BoardRevision NotesUnit 2: Compound Structure & PropertiesVSEPR & HybridizationVSEPR Theory

VSEPR Theory (College Board AP® Chemistry): Revision Note

Martín

Written by: Martín

Reviewed by: Stewart Hird

Updated on

VSEPR Theory

  • The valence shell electron pair repulsion (VSEPR) theory is based on the principle that electron domains, such as lone pairs, single bonds, double bonds, triple bonds or single unpaired electrons, repel one another through Coulombic forces

  • To minimize this repulsion, the electron domains are arranged as far apart in space as possible

  • VSEPR theory has three rules:

    1. Bonding pairs and lone pairs are arranged as far apart in space as possible

    2. A double or triple bond counts as one electron domain

    3. Lone pairs repel more strongly than bonding pairs

  • Counting the electron domains around a central atom lets you predict the structural and electron properties of a molecule

Two electron domains

  • If there are two electron domains around the central atom, the bond angle is 180°, which maximizes their separation

  • The molecular geometry is linear

    • E.g. BeCl2, CO2 and ethyne (HC≡CH)

Examples of linear molecules

Three linear molecules with 180° bond angles marked: beryllium chloride (Cl–Be–Cl), carbon dioxide (O=C=O) and ethyne (H–C≡C–H)
Beryllium chloride, carbon dioxide and ethyne all have two electron domains

Three electron domains

  • If there are three electron domains around the central atom, the bond angle is 120°, which maximizes their separation

  • The molecular geometry is trigonal planar

    • E.g. BF3, CH2CH2 and CH2O

Examples of trigonal planar molecules

Three trigonal planar molecules with 120° bond angles marked: boron trifluoride, ethene with a C=C double bond, and methanal (H₂C=O)
Boron trifluoride, ethene and methanal all have three electron domains

  • If there are two bonding domains and one lone pair, the molecular geometry is bent and the bond angle is less than 120°

    • E.g. SO2

Four electron domains

  • Molecules with four electron domains have three-dimensional geometries

  • If there are four electron domains around the central atom, there are three possible geometries:

    • If all four electron domains are bonding domains, the bond angle is 109.5°, which maximizes their separation

      • The molecular geometry is tetrahedral

      • E.g. CH4, NH4+

Examples of tetrahedral molecules

Methane and the ammonium ion drawn in three dimensions with wedge and dash bonds, each with a 109.5° bond angle marked
Methane and ammonium ions have four electron domains

  • If there are three bonding domains and one lone pair, the bond angle is approximately 107°, which is slightly less than 109.5° because the lone pair repels the bonding pairs more strongly

    • The molecular geometry is trigonal pyramidal

    • E.g. NH3

NH3 has a trigonal pyramidal geometry

Ammonia drawn in three dimensions as a trigonal pyramid, with one N–H bond on a wedge and an H–N–H bond angle of about 107°, labeled molecular shape
The molecular geometry of ammonia

  • If there are two bonding domains and two lone pairs, the bond angle is approximately 104.5°, which is less than 109.5° because the two lone pairs repel the bonding pairs more strongly

Bond angles in water

Water with two lone-pair lobes and two bonding pairs on oxygen, labeled greatest repulsion between lone pairs, intermediate between a lone pair and a bonding pair, least between bonding pairs; H–O–H angle 104.5°
The order of electron pair repulsion is lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair

  • The molecular geometry is bent

  • E.g. H2O

Water has a bent geometry

Water drawn as a bent molecule with an H–O–H bond angle of about 104.5°, labeled molecular shape
The molecular geometry of water

Five electron domains

  • Molecules with five electron domains have three-dimensional geometries

  • If there are five electron domains around the central atom, there are four possible geometries:

    • If there are five bonding domains, the angle between the equatorial bonds is 120°, the angle between an axial and an equatorial bond is 90°, and the angle between the two axial bonds is 180°. This arrangement minimizes repulsion

      • The molecular geometry is trigonal bipyramidal

      • E.g. PCl5

PCl5 has a trigonal bipyramidal geometry

Phosphorus pentachloride drawn as a trigonal bipyramid: 120° between equatorial bonds, 90° between axial and equatorial bonds, 180° between the two axial bonds
The molecular geometry of phosphorus pentachloride

  • If there are four bonding domains and one lone pair, the lone pair sits in an equatorial position

    • The angle between the equatorial bonds is less than 120° and the angle between an axial and an equatorial bond is less than 90°, because the lone pair repels the bonding pairs more strongly

    • The molecular geometry is seesaw

    • E.g. SF4

SF4 has a seesaw geometry

Sulfur tetrafluoride drawn as a seesaw, with one lone pair in an equatorial position and angles below 120° and below 90° marked
The molecular geometry of sulfur tetrafluoride

  • If there are three bonding domains and two lone pairs, the molecule is flat

  • The bond angle is slightly less than 90° because the lone pairs repel the bonding pairs more strongly

    • The molecular geometry is T-shaped

    • E.g. ClF3

ClF3 has a T-shaped geometry

Chlorine trifluoride drawn as T-shaped, with two lone pairs in equatorial positions and an F–Cl–F angle just under 90°
The molecular geometry of chlorine trifluoride

  • If there are two bonding domains and three lone pairs, the molecule is flat

  • The bond angle is 180° because the three lone pairs sit in the equatorial positions, where their repulsions cancel

    • The molecular geometry is linear

    • E.g. I3-

I3- has a linear geometry

Triiodide ion drawn as linear inside brackets with a negative charge, with three lone pairs around the central iodine atom
The molecular geometry of the triiodide ion

Six electron domains

  • Molecules with six electron domains have three-dimensional geometries

  • If there are six electron domains around the central atom, there are three possible geometries:

    • If there are six bonding domains, the bond angle is 90°, which maximizes their separation

      • The molecular geometry is octahedral

      • E.g. SF6

SF6 has an octahedral geometry

Sulfur hexafluoride drawn in three dimensions as an octahedron with six S–F bonds
The molecular geometry of sulfur hexafluoride

  • If there are five bonding domains and one lone pair, the bond angle is slightly less than 90° because the lone pair repels the bonding pairs more strongly

    • The molecular geometry is square pyramidal

    • E.g. BrF5

BrF5 has a square pyramidal geometry

Bromine pentafluoride drawn in three dimensions as a square pyramid, with one lone pair below the bromine atom
The molecular geometry of bromine pentafluoride

  • If there are four bonding domains and two lone pairs, the molecule is flat

  • The bond angle is 90°, and the lone pairs are placed at 180° to each other, which minimizes the repulsion between them

    • The molecular geometry is square planar

    • E.g. XeF4

XeF4 has a square planar geometry

Xenon tetrafluoride drawn as square planar, with one lone pair above and one below the xenon atom
The molecular geometry of xenon tetrafluoride

  • The table summarizes the VSEPR shapes:

Summary of the VSEPR theory

Electron domains

Electron-domain arrangement

Bonding domains

Lone pairs

Molecular geometry

Bond angles

2

linear

2

0

linear

180°

3

trigonal planar

3

0

trigonal planar

120°

3

trigonal planar

2

1

bent

<120°

4

tetrahedral

4

0

tetrahedral

109.5°

4

tetrahedral

3

1

trigonal pyramidal

<109.5° (107° in NH3)

4

tetrahedral

2

2

bent

<109.5° (104.5° in H2O)

5

trigonal bipyramidal

5

0

trigonal bipyramidal

120° (equatorial)

90° (axial–equatorial)

180° (axial–axial)

5

trigonal bipyramidal

4

1

seesaw

<120° (equatorial)

<90° (axial–equatorial)

5

trigonal bipyramidal

3

2

T-shaped

<90°

5

trigonal bipyramidal

2

3

linear

180°

6

octahedral

6

0

octahedral

90°

6

octahedral

5

1

square pyramidal

<90°

6

octahedral

4

2

square planar

90°

Examiner Tips and Tricks

  • Lewis diagrams and VSEPR are used together to predict molecular geometry, bond angles and whether a molecule has a dipole moment

  • Always explain shapes in terms of the Coulombic repulsion between electron domains

  • The number of electron domains also tells you the hybridization (2 = sp, 3 = sp2, 4 = sp3), but hybridization only names the arrangement; it does not explain it

  • Answer the question asked: if it asks for a geometry, give the shape name, not the hybridization

Molecular Geometry

Predicting molecular geometry

  • The molecular geometry of any molecule can be determined by following these steps:

    1. Draw an accurate Lewis diagram for the molecule or ion

    2. Count the number of electron domains

    3. Count the number of bonding domains and the number of lone pairs

    4. Apply the VSEPR theory to predict the molecular geometry

Examiner Tips and Tricks

  • Count electron domains around the central atom first.

    • Each lone pair, single bond, double bond, triple bond or single unpaired electron is one domain

  • The electron-domain arrangement uses all domains; the molecular geometry is named only from the positions of the atoms.

    • NH3 has a tetrahedral arrangement of domains but a trigonal pyramidal geometry

  • A bond angle is a number and a geometry is a name

    • Give exactly what the question asks for, and make sure your angle matches your geometry (a linear molecule has 180°, not 120°)

Worked Example

Predict the molecular geometry in the following molecules or ions:

  1. H2S

  2. NH2Cl

  3. ICl4-

Answers:

  • Answer 1: The total number of valence electrons in H2S is 1 + 1 + 6 = 8, so there are four pairs of electrons around S

    • H2S has 4 electron domains (2 bonding domains and 2 lone pairs), so the molecular geometry is bent

    Four steps for H₂S: Lewis diagram with two lone pairs on sulfur; 4 electron domains; 2 lone pairs and 2 bonding domains; molecular geometry bent

  • Answer 2: The total number of valence electrons in NH2Cl is 5 + 1 + 1 + 7 = 14

    • NH2Cl has 4 electron domains around N (3 bonding domains and 1 lone pair), so the molecular geometry is trigonal pyramidal

    Four steps for NH₂Cl: Lewis diagram with one lone pair on nitrogen and three on chlorine; 4 electron domains; 1 lone pair and 3 bonding domains; trigonal pyramidal

  • Answer 3: The total number of valence electrons in ICl4- is (7 × 5) + 1 = 36

    • ICl4- has 6 electron domains around I (4 bonding domains and 2 lone pairs), so the molecular geometry is square planar

    Four steps for ICl₄⁻: Lewis diagram with two lone pairs on iodine and three on each chlorine; 6 electron domains; 2 lone pairs and 4 bonding domains; square planar

Bond Angles

Predicting bond angles

  • The bond angles in any molecule can be determined by following these steps:

    1. Draw an accurate Lewis diagram for the molecule or ion

    2. Count the number of electron domains and place them as far apart as possible

    3. Count the number of bonding domains and the number of lone pairs

    4. Apply the VSEPR theory to predict the bond angle

Worked Example

Predict the bond angles in the following molecules:

  1. CH2O

  2. SF4

  3. XeF4

Answers:

  • Answer 1: CH2O has 4 + (2 × 1) + 6 = 12 valence electrons

    • Carbon is the central atom, with a double bond to oxygen and a single bond to each hydrogen

    • There are 3 electron domains around carbon, all bonding, and no lone pairs

    • The molecular geometry is trigonal planar, so the H–C–H and H–C=O bond angles are about 120°

  • Answer 2: SF4 has 6 + (4 × 7) = 34 valence electrons

    • There are 5 electron domains around sulfur: 4 bonding domains and 1 lone pair

    • The molecular geometry is seesaw

    • The lone pair repels the bonding pairs more strongly, so the equatorial F–S–F angle is less than 120° and the axial–equatorial F–S–F angles are less than 90°

  • Answer 3: XeF4 has 8 + (4 × 7) = 36 valence electrons

    • There are 6 electron domains around xenon: 4 bonding domains and 2 lone pairs

    • The two lone pairs sit opposite each other, so the molecular geometry is square planar

    • The F–Xe–F bond angles are 90°

Relative Bond Energies

Bond energy

  • The bond energy is the amount of energy required to break 1 mol of bonds in the gas phase

    • It reflects the strength of a chemical bond between two atoms. The higher the energy, the stronger the bond

  • It depends on the bond order: more shared pairs between the two nuclei give a stronger, shorter bond

  • There are some important rules to follow when predicting relative bond energy:

    • Single bonds are generally weaker than double bonds, and double bonds are weaker than triple bonds, so single bonds have the lowest bond energy

    • Multiple bonds are stronger because there is a greater Coulombic attraction between the shared electrons and the nuclei of the atoms

    • Triple bonds have the highest bond energy, so they are the strongest

Bond energies of carbon–carbon bonds

Bond

Bond length (pm)

Bond energy (kJ/mol)

C–C

154

347

C=C

134

614

C≡C

120

839

  • Bond order can also be used to compare bond energies

Bond order

  • The bond order is the number of shared electron pairs between two bonded atoms:

    • 1 for a single bond

    • 2 for a double bond

    • 3 for a triple bond

  • A higher bond order between two atoms means a shorter bond with a higher bond energy

  • Where resonance spreads a bond over several positions, the bond order can be a fraction (e.g. 1.5 for each O–O bond in O3)

Worked Example

Which of these molecules has the carbon–carbon bond with the highest bond energy: ethane, ethene or ethyne?

Structural formulas of ethane, ethene with a C=C double bond, and ethyne with a C≡C triple bond

Answer:

  • Ethane: C–C single bond, bond order 1

  • Ethene: C=C double bond, bond order 2

  • Ethyne: C≡C triple bond, bond order 3

  • The higher the bond order, the more electrons are shared between the two carbon nuclei

    • So, the Coulombic attraction is greater and the bond is stronger

  • Ethyne's C≡C bond has the highest bond energy

    • Ethyne = 839 kJ/mol, compared with 614 kJ/mol for C=C and 347 kJ/mol for C–C

Relative Bond Lengths

Bond length

  • The bond length is the distance between the nuclei of two bonded atoms

  • It depends on the bond order (single, double or triple) and on the atomic radii of the bonded atoms

  • There are some important rules to follow when predicting relative bond lengths:

    • If the atomic radius of the atoms involved increases, the bond length increases

      • This is because the shared electrons are attracted less strongly by the nuclei of the atoms

    • Multiple bonds decrease the bond length because there is a greater Coulombic attraction between the shared electrons and the nuclei of the atoms

    • Triple bonds are the shortest because of the greater Coulombic attraction between the shared electrons and the nuclei

Bond lengths of carbon–carbon bonds

Bond

Bond length (pm)

Bond energy (kJ/mol)

C–C

154

347

C=C

134

614

C≡C

120

839

  • Bond order can also be used to compare bond lengths

    • To see how bond order is found, see the previous section, Relative Bond Energies

Worked Example

Which of the following molecules will have the longest single covalent bond?

(A) HCl

(B) HF

(C) HBr

(D) HI

Answer:

  • The correct answer is D

  • This is because the iodine atom has the largest radius of I, Cl, F and Br

  • This can be found from the periodic table: iodine is below F, Cl and Br in group 17

  • Therefore, the Coulombic attraction between the iodine nucleus and the shared electrons is the smallest of the four halogens

Dipole Moments

  • A dipole moment measures the separation of positive and negative charge within a molecule

  • A molecule with a net dipole moment is polar

  • Two factors decide whether a molecule is polar:

    • The presence of polar covalent bonds in the molecule

    • The molecular geometry

  • Some molecules have polar covalent bonds, but their bond dipoles are canceled out by the molecular geometry

    • This happens because the atoms are arranged symmetrically, so the individual bond dipoles cancel each other out

    • If the bond dipoles are equal in size and point in opposite directions, they cancel and the molecule has no net dipole moment

Polar and nonpolar molecules

  • For example, CH3Cl is a polar molecule. Its net dipole moment points toward the chlorine atom, which attracts the shared electrons with the strongest Coulombic force

The CH3Cl molecule

Chloromethane with bond dipole arrows toward carbon from each hydrogen and toward chlorine from carbon; overall dipole points to chlorine; labeled polar
There are four polar covalent bonds in CH3Cl which do not cancel each other out, so CH3Cl is a polar molecule; the overall dipole is toward the electronegative chlorine atom

  • CCl4 is a nonpolar molecule

    • The C–Cl bonds are polar because of the difference in electronegativity between C and Cl, but the four equal bond dipoles are arranged symmetrically

    • So, the net dipole moment is zero

The CCl4 molecule

Tetrachloromethane with four bond dipole arrows pointing from carbon to each chlorine, carbon δ+ and each chlorine δ−, which cancel; labeled nonpolar
Though CCl4 has four polar covalent bonds, the individual bond dipoles cancel each other out, so CCl4 is a nonpolar molecule

Unlock more, it's free!

Join the 100,000+ Students that ❤️ Save My Exams

the (exam) results speak for themselves:

I would just like to say a massive thank you for putting together such a brilliant, easy to use website.I really think using this site helped me secure my top gradesin science and maths. You really did save my exams! Thank you.

Beth
IGCSE Student

This website is soooo useful and I can’t ever thank you enough for organising questions by topic like this. Furthermore, the name of the website could not have been more appropriate as it literally did SAVE MY EXAMS!

Fathima
A Level Student

Incredible! SO worth my money, the revision notes have everything I need to know and are so easy to understand. I actually enjoy revising! It makes me feel a lot more confident for my GCSEs in a few months.

Kate
GCSE Student

Absolutely brilliant, both my girls used it for A levels and GCSE. It's saves on paper copies, also beneficial exam questions ranked from easy to hard. It's removed a lot of stress from the exams.

Sameera
Parent

Just to say that your resources are the best I have seen and I have been teaching chemistry at different levels for about 40 years

Mark
Chemistry Teacher

Excellent

Our rating represented in Trustpilot stars
Read more reviews

Build on this topic

Check your understandingTry an Exam Question on this topicTest your recallReinforce key facts and definitions with Flashcards
Previous:Limitations of Lewis StructuresNext:Hybridization
Author
Reviewer
Martín

Author: Martín

Expertise: Chemistry Content Creator

Martín, a dedicated chemistry teacher and tutor, excels in guiding students through IB, AP, and IGCSE Chemistry. As an IB Chemistry student, he came from hands-on preparation, focusing on practical exam techniques and rigorous practice. While at Universidad San Francisco de Quito, his academic journey sparked a passion for computational and physical chemistry. Martín specializes in chemistry, and he knows that SaveMyExams is the right place if he wants to have a positive impact all around the world.

Stewart Hird

Reviewer: Stewart Hird

Expertise: Chemistry Content Creator

Stewart has been an enthusiastic GCSE, IGCSE, A Level and IB teacher for more than 30 years in the UK as well as overseas, and has also been an examiner for IB and A Level. As a long-standing Head of Science, Stewart brings a wealth of experience to creating Topic Questions and revision materials for Save My Exams. Stewart specialises in Chemistry, but has also taught Physics and Environmental Systems and Societies.