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Shapes of Simple Molecules & Ions

Lajoy Tucker

Teacher

Lajoy Tucker

Contents

Shapes of Simple Molecules

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Notation

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Working Out Electron Pairs

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Molecules with Double Bonds

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Shapes of Simple Molecules

The shape of a simple molecule is determined using valence shell electron pair repulsion (VSEPR) theory. It is based on the key principle that electron pairs repel each other and so position themselves as far from each other as possible. 

Electron Pairs Electron Geometry Angle Example  
2 Linear   Diagram showing the linear structure of carbon dioxide with a bond angle of 180 degrees between the two oxygen atoms.
3 Trigonal Planar Diagram showing the trigonal planar structure of boron trifluoride with a bond angle of 120 degrees between fluorine atoms.
4 Tetrahedral Diagram showing the tetrahedral shape of methane with hydrogen atoms bonded to carbon at bond angles of 109.5 degrees.
5

Trigonal Bipyramidal

 Diagram showing the trigonal bipyramidal structure of phosphorus pentachloride with bond angles of 90 degrees and 120 degrees between chlorine atoms.
6 Octahedral Diagram showing the octahedral structure of sulfur hexafluoride with 90 degree bond angles between fluorine atoms.

 

Notation

Each line represents a single covalent bond (a shared pair of electrons). To represent the shapes in three dimensions:

  • Solid lines represent bonds in the plane of the paper

  • Dashed lines represent bonds going into the page (away from the viewer)

  • Wedges represent bonds coming out of the page (towards the viewer)

Lone pairs can be represented as two dots and/or a lobe as below.

Lone pairs repel more strongly than bonding pairs.

This affects bond angles:

More lone pairs → smaller bond angles as they reduce bond angles by ~2.5 degrees for each lone pair.

Order of Repulsion:

Lone pair–lone pair > Lone pair–bond pair > Bond pair–bond pair

Example 1

3 bonding pairs and 1 lone pair around the central atom. e.g .

Electron geometry = tetrahedral (4 electron pairs total)

The lone pair reduces the typical tetrahedral bond angles ()by resulting in bond angles of .

Diagram showing the trigonal pyramidal structure of ammonia with a lone pair of electrons and a bond angle of 107 degrees between hydrogen atoms.

No answer provided.

Example 2

2 bonding pairs and 2 lone pairs around the central atom e.g.,

Electron geometry = tetrahedral (4 electron pairs total)

Diagram showing the bent structure of a water molecule with two hydrogen atoms bonded to an oxygen atom containing two lone pairs of electrons.

The lone pairs reduces the typical tetrahedral bond angles () by 2 x resulting in bond angles of .

Due to the increased repulsion of lone pairs, where there is more than one option for their placement, they are positioned such that overall repulsion is reduced.

No answer provided.

Example 3

4 bonding pairs, and 1 lone pair surrounding the central atom e.g

Electron geometry = trigonal bipyramidal (5 electron pairs)

The lone pair must be positioned to minimise repulsion and therefore takes an equatorial position where it is from two bonds and from two bonds.

Diagram showing the seesaw structure of sulfur tetrafluoride with four fluorine atoms bonded to sulfur and one lone pair of electrons on the sulfur atom.

This results in a see-saw shape.

No answer provided.

Example 4

4 bonding pairs, and 2 lone pairs surrounding the central atom e.g.,

Electron geometry = octahedral (6 electron pairs)

The 2 lone pairs must be positioned to minimise repulsion and therefore take the axial positions to be as far from each other as possible.

Diagram showing xenon tetrafluoride changing from an octahedral electron arrangement with two lone pairs to a square planar molecular shape with four fluorine atoms around xenon.

This results in a square planar shape

No answer provided.

Note – the bond angles here do not change as the lone pairs are positioned opposite each other.

No answer provided.

Working Out Electron Pairs

The number of electron pairs around a central atom can be determined using various methods.

This electron counting method works where only single covalent bonds are being formed e.g. bonds with hydrogen atoms or halogens.

Step 1: Count valence electrons on central atom (based on group number)

Step 2: Add the number of bonds being formed

Step 3: Add electrons if there is a negative charge or subtract electrons if there’s a positive charge

Step 4: Divide by 2 to determine total number of electron pairs

Step 5: Deduce number of bonding pairs and lone pairs

Example 1 -

1. Number of electrons around central carbon atom = 4

2. Add 4 for the number of bonds being formed = 8

3. No charge so skip this step

4. Divide by 2 🡪 8/2 = 4 electron pairs

5. All 4 of the electron pairs are bonding pairs (as there are 4 F’s)

Therefore, 4 bonding pairs and 0 lone pairs

Shape = Tetrahedral

Bond angle =

Diagram showing the tetrahedral molecular structure of carbon tetrafluoride with four fluorine atoms bonded to a central carbon atom.

No answer provided.

Note – drawing a dot and cross diagram can be used instead and gives the same answer.

Dot and cross diagram showing the covalent bonding and shared electron pairs in a carbon tetrafluoride molecule.

No answer provided.

Example 2 -

1. Number of electrons around central chlorine atom = 7

2. Add 2 for the number of bonds being formed = 9

3. Add 1 for the 1- charge = 10

4. Divide by 2 🡪 10/2 = 5 electron pairs

Therefore, 2 bonding pairs and 3 lone pairs

Based on trigonal bipyramidal

3 lone pairs in equatorial positions (around the middle) to minimise repulsion

Shape = Linear

Bond angle =

Diagram showing the three-dimensional and displayed formula representations of the tetrafluorochlorate ion with fluorine atoms bonded to a central chlorine atom.

No answer provided.

Molecules with Double Bonds

Oxygen forms double bonds sharing two pairs of electrons with bonded atoms. As such, the above electron counting method is not suitable.

Constructing dot and cross diagrams is appropriate here to ensure the double bonds are accounted for.

Dot and cross diagram showing the covalent bonding and shared electron pairs in a sulfur dioxide molecule.

Double bonds here can be treated the same as single bonds when using VSEPR. They can be called bonding ‘regions’ as opposed to bonding pairs.

has 2 bonding regions and 1 lone pair

Shape based on trigonal planar (3 electron pairs)

Shape = Bent

Bond angle = 120-2.5 =

The O=S=O bond angle is actually due to increased repulsion between the double bonds but the general rule of subtracting can be used in an exam context.

Displayed formula showing sulfur dioxide with two double bonds between sulfur and oxygen atoms and a lone pair on the sulfur atom.

Question:

Draw the shape of

1. Number of electrons around central sulfur atom = 6

2. Add 5 for the number of bonds being formed = 11

3. Add 1 for the 1- charge = 12

4. Divide by 2 🡪 12/2 = 6 electron pairs

5. Only 5 of these are bonding pairs (as there are 5 F’s)

Therefore, 5 bonding pairs and 1 lone pair. Based on octahedral shape (6 electron pairs)

Note – as all bond angles in an octahedron are the same, it does not matter where the lone pair is positioned. All positions are the same.

Shape = Square pyramidal

 

Summary

Total number of electron pairs

Number of bonding pairs

Number of lone pairs

Shape

2

2

0

 Linear Molecules with Double Bonds

Linear

3

3

0

 Trigonal planar

Trigonal planar

2

1

 Bent (V-shape)

Bent (V-shape)

4

4

0

 Tetrahedral

Tetrahedral

3

1

 Trigonal pyramidal

Trigonal pyramidal

2

2

 Bent (V-shape)

Bent (V-shape)

5

5

0

 Trigonal bipyramidal

Trigonal bipyramidal

4

1

 Trigonal pyramidal or see-saw orTrigonal pyramidal or see-saw

Trigonal pyramidal or see-saw

3

2

 Trigonal planar or T-shape or Trigonal planar or T-shape

Trigonal planar or T-shape

6

6

0

 Octahedral

Octahedral

5

1

 Square pyramidal

Square pyramidal

4

2

 Square planar

Square planar
No answer provided.
Electronegativity and Polarity