Compounds With Octahedral Shape

For compounds whose central atoms possess six bonding electron pairs, their bonded atoms are oriented towards the corners of an octahedron.

The best examples are the PF₆^- ( phosphorus hexafluoride or hexafluorophosphate) ion and the SF₆ ( sulfur hexafluoride ) molecule whose Lewis structures are shown below.

The ground state electronic configuration of the two central atoms:

P	1s2 2s2 2p6 3s2 3px1 3py1 3pz1
S	1s2 2s2 2p6 3s2 3px2 3py1 3pz1

In the excited state of the two central atoms above, their valence electrons are assumed to be distributed this way:

1s2 2s2 2p6 3s1 3px1 3py1 3pz1 3dz21 3dx2 - y21

Since six equal orbitals are required, it is assumed that sp³d² hybridization is used to form six hybrid orbitals:

1s2 2s2 2p6 (sp3d2)1 (sp3d2)1 (sp3d2)1 (sp3d2)1 (sp3d2)1 (sp3d2)1

Phosphorus acquires an extra electron for its sixth orbital.



Geometry of PF₆^-

• shape of ion: octahedral
• bond angle of axial atoms: 180°
• bond angle of equatorial atoms: 90°
• bond angle between an equatorial atom and an axial atom: 90°

Axial atoms are in white circles; equatorial atoms are in black circles.

An octahedron is superimposed on the PF₆^- ion.

Geometry of SF₆

• shape of molecule: octahedral
• bond angle of axial atoms: 180°
• bond angle of equatorial atoms: 90°
• bond angle between an equatorial atom and an axial atom: 90°

Molecules With Trigonal Bipyramidal Shape

Molecules like PCl₅ (phosphorus pentachloride) and PBr₂Cl₃ (phosphorus dibromo-trichloride) possess five covalent bonds whose electron pairs are arranged in a trigonal bipyramidal geometry.

The central atom of these molecules, phosphorus, has only three unpaired electrons out of its five valence electrons as can be seen from its electronic configuration below:

1s2 2s2 2p6 3s2 3px1 3py1 3pz1

ground state configuration

But the Lewis structures of the given molecules, as shown below, indicate that the phosphorus atom uses five unshared electrons to form five covalent bonds.

It is, therefore, assumed that phosphorus uses the empty d orbital to "promote" 1 electron from 3s orbital to d orbital,

1s2 2s2 2p6 3s1 3px1 3py1 3pz1 3dz21

excited configuration

and, through hybridization, form five sp³d hybrid orbitals which are then used to bond with chlorine and bromine atoms.

1s2 2s2 2p6 (sp3d)1 (sp3d)1 (sp3d)1 (sp3d)1 (sp3d)1

excited configuration

Geometry of PCl₅

• shape of molecule: trigonal bipyramidal
• bond angle of axial atoms: 180°
• bond angle of equatorial atoms: 120°
• bond angle between an equatorial atom and an axial atom: 90°

Axial atoms are in white circles; equatorial atoms are in black circles.

A trigonal bipyramid is superimposed on the PCl₅ molecule.

Geometry of PBr₂Cl₃

Due to the different axial and equatorial positions the two bromine atoms can assume, PBr₂Cl₃ (phosphorus dibromo-trichloride) has three geometric isomers as shown below.

Except for the first isomer in which the bromine atoms are in axial positions, the other two geometric isomers of PBr₂Cl₃ are polar.

Molecular Geometry: Molecules With Trigonal Planar Shapes

Molecules with trigonal planar shape are characterized by their central atoms having 3 bonding electron pairs.

These electron pairs refer to those ones involved in sigma bonds only.

Electron pairs from pi bonds of double and triple bonds do not count.

In this particular coordination geometry, each electron pair is at an angle of 120° from any of the other two electron pairs.

The illustrations of the geometry and the Lewis structures given below are for the following molecules:

BF3

BCl3

SO3

H2C=O

Geometry of BF₃

• shape of molecule: trigonal planar
• F-B-F bond angle: 120°

A trigonal plane is superimposed on the BF₃ molecule.

Lewis structures of BF₃ and BCl₃

Due to its 3 valence electrons, boron does not follow the octet rule and thus forms electron-deficient compounds such as BF₃ ( boron trifluoride ) and BCl₃ ( boron trichloride ).

The characteristic reaction of the above compounds is to accept/share an electron pair from other atoms or compounds.

The reaction is called Lewis acid-base reaction.

A Lewis acid accepts a pair of electrons; a Lewis base donates a pair of electrons.



Geometry of BCl₃

Geometry of SO₃

Lewis structures of SO₃ and H₂C=O

Shown above is the Lewis structures of SO₃ ( sulfur trioxide ) and H₂C=O ( formaldehyde ) molecules.



Geometry of H₂C=O

Molecular Geometry: Compounds With Tetrahedral Shapes

A compound with four electron pairs at its central atom has a tetrahedral shape.

If one of the four electron pairs is unshared or non-bonding, then the compound becomes pyramidal in shape.

Although triangular pyramid is also a tetrahedron, there is a distinction between compounds with tetrahedral and pyramidal shapes: the location of their central atoms.

A compound with triangular pyramidal or simply pyramidal shape has its central atom located at the apex of the pyramid whereas a compound with tetrahedral shape has its central atom located at the center of tetrahedron.

Here are the lists of compounds included in this post:

compounds with tetrahedral shapes

CH4

F3SN

CCl4

ClO4-

NH4+

BF4-

SO4=

compounds with pyramidal shapes

NH3

SO3=

ClO3-

NF3

XeO3

F2SeO

Along with the illustration of a compound's shape, its Lewis structure is also included together with a brief discussion regarding the effect of lone electron pairs or double/triple bonds on the bond angles and/or the use of hybrid orbitals by certain atoms for covalent bonding.

The first three compounds discussed in details below serve to represent the rest of the compounds.

Methane, CH₄

• shape of molecule: tetrahedral
• H-C-H bond angle: 109.5°

The Lewis structure and one of the four equal H-C-H bond angles of CH₄.

A tetrahedron superimposed on a CH₄ molecule. The four hydrogen atoms are oriented towards the corners of the tetrahedron.

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Ammonia, NH₃

The ammonia molecule has a non-bonding electron pair at the top of the tetrahedron as depicted below.

The Lewis structure of NH₃ shows it has a lone pair of electrons which occupies more space than any of the bonding electron pairs resulting to H-N-H bond angles being less than the ideal tetrahedral angle of 109.5°.

Since molecular shapes involve atoms only, the shape of ammonia will be minus the lone pair of electrons. It is a triangular pyramid with the nitrogen atom at the apex.

• shape of molecule: triangular pyramidal
• H-N-H bond angle: 106.6°

A triangular pyramid superimposed on the ammonia molecule.

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Trifluorothionitrile, F₃S≡N

• shape of molecule: tetrahedral
• F-S-F bond angle: less than 109.5°
• N-S-F bond angle: greater than 109.5°

Referring to the Lewis structure above, sulfur atom exceeds the octet configuration or the eight valence electrons required for bonding by utilizing empty d orbitals and creating hybrid orbitals (for a discussion of constructing Lewis structure of compounds, see Drawing Lewis Electron Dot Structure or Formula).

F₃SN has a tetrahedral shape which means it has four electron pairs directed towards the four corners of a tetrahedron but its Lewis structure shows 6 bonding electron pairs (see my previous post for a discussion of compounds' different coordination geometries and the corresponding number of electron pairs of their central atoms).

In determining the shapes of compounds, only electron pairs involved in sigma bonds are considered.

As for the bond angles, greater repulsion exerted by electron pairs from double and triple bonds cause the bond angles containing them to be greater than those bond angles containing single bonds only.

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Compounds With Tetrahedral Shape

The Geometry of CCl₄

The Lewis structures of CCl₄ and ClO₄^-.

The chlorine atom in ClO₄^- or perchlorate ion acquires one electron to complete its eight valence electrons and utilizes empty d orbitals to bond with four oxygen atoms by sharing all of its eight valence electrons.

Referring to the illustration above, atoms in partial grey and red shades are atoms with formal charges. To learn more about the calculation of formal charges, read Drawing Lewis Electron Dot Structure or Formula.

The Geometry of ClO₄^-

The Geometry of NH₄⁺

The Lewis structures of NH₄⁺ and BF₄^-.

The NH₄⁺ or ammonium ion is formed by NH₃ acting as a Bronsted base and accepting a proton.

On the other hand, BF₄^- or boron tetrafluoride ion is formed by BF₃ acting as a Lewis acid and bonding with a fluoride ion acting as a Lewis base.

For a discussion of Bronsted acid-base reactions, see Solving Weak Acid/Base Dissociation Problems.



The Geometry of BF₄^-

The Geometry of SO₄⁼

The Lewis structures of SO₄⁼

Given the geometry of SO₄⁼ or sulfate ion as shown above, there are two possible Lewis structures for it. The second Lewis structure is considered to be the correct one because it is more stable and consistent with the experimental data.

The sulfur atom in the second Lewis structure exceeds the octet configuration or the 8 valence electrons required for bonding by acquiring two extra electrons and utilizing empty d orbitals.

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Compounds With Pyramidal Shape

The Geometry of SO₃⁼

The Lewis structures of SO₃⁼ and ClO₃^-

The sulfur and chlorine atoms in the above Lewis structures acquire extra electrons to fill up its valence electron shells and are assumed to use hybrid orbitals involving empty d orbitals to form covalent bonds with oxygen atoms.

The Geometry of ClO₃^-

The Geometry of NF₃

The Lewis structures of NF₃ and XeO₃

The Geometry of XeO₃

The Geometry of F₂SeO

The Lewis structure of F₂SeO

Geometries of Molecules and Ions

The shapes of compounds, either molecules or polyatomic ions, are very important in helping us understand better their reactions.

Fortunately, the geometries of most molecules and ions can be predicted quite reliably even by considering only their electron-electron pair interactions.

The idea is that the repulsive forces that exist between bonding and non-bonding pairs of electrons of a molecule or an ion cause those pairs of electrons to adapt certain spatial arrangement that allows minimum repulsion.

The spatial arrangement of the electron pairs of a molecule depends on its number of atoms and the number of valence electrons of its central atom.

So, in order to predict the geometry of a molecule or an ion, one needs to know its number of bonding and non-bonding pairs of electrons by determining the following:

• total number of atoms in the molecule or ion
• number of valence electrons of the molecule's central atom
• the Lewis structure of the molecule or ion

In the following illustrations, all of the possible coordination geometries for different compounds are depicted using ball-and-stick models.

In each illustration, information such as number of electron pairs of the compound and the number of its atoms are given.

Linear Geometry

• number of electron pairs: 2
• coordination geometry: linear
• number of atoms: 3 ( 1 central atom, 2 bonding atoms)

Trigonal Planar Geometry

• number of electron pairs: 3
• coordination geometry: trigonal planar
• number of atoms: 4 ( 1 central atom, 3 bonding atoms)

Tetrahedral Geometry

• number of electron pairs: 4
• coordination geometry: tetrahedral
• number of atoms: 4-5 ( 1 central atom, 3-4 bonding atoms)

Trigonal Bipyramidal Geometry

• number of electron pairs: 5
• coordination geometry: trigonal bipyramidal
• number of atoms: 6 ( 1 central atom, 5 bonding atoms)

Octahedral Geometry

• number of electron pairs: 6
• coordination geometry: octahedral
• number of atoms: 7 ( 1 central atom, 6 bonding atoms)

Pentagonal Bipyramidal Geometry

• number of electron pairs: 7
• coordination geometry: pentagonal bipyramidal
• number of atoms: 8 ( 1 central atom, 7 bonding atoms)