Figure 2.0.4 In this Lewis structure of Cl2CO, the electrons of the C=O double bond exert a greater repulsive force than do the electrons in the C–Cl single bonds, so the O–C–Cl bond angle is 124.3° rather than 120° as found in trigonal planar molecules such as BF3.
C O Cl Cl 124.3° 124.3° 111.4°
five negative charge centres
14.1.1
Predict the shape and bond angles for species with five and six negative charge centres using the VSEPR theory. ©IBO 2007
DEMO 2.1
Modelling five and six negative charge centres
Figure 2.1.1 (a) A trigonal bipyramid has six faces and five vertices. (b) When five atoms are bonded to a central atom, the shape of the molecule is called a trigonal bipyramid.
90°
a b
There are three atoms in a plane forming an equilateral triangle around the middle of the molecule (the equatorial positions), one atom at the top and one atom at the bottom of the molecule (axial positions). The plane of the equatorial atoms is at right angles to that of the axial atoms.
Example 1: PCl5
Phosphorus pentachloride (phosphorus(V) chloride) is a molecule in which a phosphorus atom has formed covalent bonds with five chlorine atoms. This results in the phosphorus atom having 10 electrons in its valence shell (an expanded valence shell). The shape of the molecule is described as trigonal bipyramidal and the molecule is non-polar. The bond angle between the equatorial atoms and the central phosphorus atom is 120°. This is consistent with the equilateral triangle formed by the three chlorine atoms in that position. The bond angle between an equatorial chlorine atom and an axial chlorine atom will be 90°.
Figure 2.1.2 Phosphorus pentachloride, PCl5, has a trigonal bipyramidal shape. Cl Cl Cl Cl Cl 120° P
There are other molecules with five negative charge centres around them; however, phosphorus, being in group 5, is the only member of period 3 that bonds with five atoms and has no pairs of non-bonding electrons. The other molecules have combinations of bonding and non-bonding pairs of electrons, so with the greater electrostatic repulsion of the non-bonding charge centres and the smaller numbers of atoms bonded to the central atom, a range of shapes is assumed.
Example 2: ClF3
Chlorine trifluoride is a very reactive, poisonous and corrosive gas. Chlorine has seven electrons in its valence shell and the formula tells us that three bonds are formed with fluorine. This leaves four non-bonding electrons (two pairs) in the valence shell. Consequently there are five negative charge centres around this chlorine atom—-three bonding and two non-bonding. The non-bonding pairs of electrons take up equatorial positions around the chlorine to minimize the electrostatic repulsions in the molecule. The angle between the equatorial fluorine and the axial fluorine atoms is slightly less than 90° (87.5°) due to the repulsive forces exerted by the non-bonding electrons. The shape of this molecule is commonly referred to as ‘T-shaped’.
Figure 2.1.3 Chlorine trifluoride is a ‘T-shaped’ molecule. F Cl F F 87.5° 185°
CHAPTER 2
bonding
In table 2.1.1 a range of shapes for molecules with five negative charge centres is shown. Notice the effect of the non-bonding electrons on the shapes (in particular the angle between axial atoms) of these molecules.
Table 2.1.1 Summary of moleculeS wiTh five negaTive charge cenTreS around The cenTral aTom
number of bonding pairs on the central atom
number of non- bonding pairs on
the central atom
example Shape of molecule representation of the shape
5 0 PCl5 Trigonal bipyramid Cl P Cl Cl Cl Cl 90° 120° 4 1 SF4 Seesaw or ‘sawhorse’ F S F F F 90° 116° 186.9° 3 2 ICl3 T-shaped Cl I Cl Cl 86.2° 187.2° 2 3 XeF2 Linear F F Xe
When there are six negative charge centres around a central atom they will move as far apart as possible, according to VSEPR theory, in order to minimize the repulsions between them. If 6 atoms are bonded to the central atom, all six negative charge centres are bonding electron pairs and the shape that the molecule takes on is called an octahedron. This name comes from the solid shape formed when the vertices are joined. An octahedron has eight faces and six vertices. This shape can be thought of as two square pyramids fused together at their bases, and so can be described as square bipyramidal. In an octahedral molecule four atoms in a plane form a square around the middle of the molecule (the equatorial positions) and there is one atom at the top and one atom at the bottom of the molecule (axial positions). The plane of the equatorial atoms is at right angles to that of the axial atoms. Six negative charge centres
Figure 2.1.4 (a) An octahedron has eight faces and six vertices. (b) When six atoms are bonded to a central atom the molecule is octahedral.
Example 3: SF6
Sulfur hexafluoride (sulfur(VI) fluoride) is a molecule in which a sulfur atom has formed covalent bonds with six chlorine atoms. This results in the sulfur atom having 12 electrons in its valence shell (another example of an expanded valence shell). The shape of the molecule is described as octahedral or square bipyramidal and the molecule is non-polar. The bond angle between the equatorial atoms and the central sulfur atom (F–S–F) is 90°. This is consistent with the square formed by the four fluorine atoms in that position. The bond angle between an equatorial fluorine atom and an axial fluorine atom will be 90°.
Figure 2.1.5 Sulfur hexafluoride, SF6, has an octahedral shape.
F F F F F F S Example 4: PF6−
The phosphorus hexafluoride ion, PF6−, is a Lewis base that is formed when PF5 reacts with XeF2. While its chemistry is unusual, its shape is the same as that of sulfur hexafluoride—octahedral, or square bipyramidal. This molecule is an anion (negative ion) because in addition to the five fluorine atoms that would normally covalently bond with a phosphorus atom, an extra fluoride ion, F−, has bonded in a dative covalent (coordinate) bond with the phosphorus atom. In PF6−, the phosphorus atom has 12 electrons in its valence shell. There are other molecules with six negative charge centres around them; however, sulfur, being in group 6, is the only member of period 3 that bonds with six atoms to form a neutral molecule and has no non-bonding pairs of electrons. The other molecules have combinations of bonding and non-bonding pairs of electrons, so with the greater electrostatic repulsion of the non- bonding charge centres and the smaller numbers of atoms bonded to the central atom, a number of different shapes is assumed.
Example 5: XeF4
Under conditions of high voltage and with the very electronegative element fluorine, xenon will react to form xenon tetrafluoride, XeF4. Xenon has 8 electrons in its valence shell. According to the formula, four bonds are formed, so there are four non-bonding electrons (two non-bonding pairs). These non- bonding pairs take up the axial positions in the octahedral arrangement of negative charge centres to minimize the electrostatic repulsions in the molecule. As a result, the XeF4 molecule takes on a square planar shape in which the F–Xe–F bond angle is 90°. The XeF4 molecule would be non-polar. Notice that there are 12 electrons in the valence shell of the xenon atom. The Figure 2.1.6 The phosphorus
hexafluoride ion, PF6−, has an octahedral shape. P F F F F F F
Figure 2.1.7 Xenon tetrafluoride, XeF4, is a square planar molecule.
Xe F F F
CHAPTER 2
bonding
Figure 2.1.8 When xenon tetrafluoride is isolated, it forms colourless, roughly cubic crystals.
Table 2.1.2 summarizes the shapes that can be taken on by molecules with six negative charge centres. Notice that in the case of square pyramidal molecules, the one non-bonding pair of electrons repels the square plane upwards to give an angle that is less than 90° with the atom at the ‘top’ of the pyramid.
Table 2.1.2 Summary of moleculeS wiTh Six negaTive charge cenTreS around The cenTral aTom
number of bonding pairs on the central atom number of non-bonding pairs on the central atom example Shape of
molecule representation of the shape
6 0 SF6−, PF6− Octahedral or square bipyramidal S F F F F F 90° 90° F 5 1 BrF5 Square pyramidal Br F F F F F 85° 90° 4 2 XeF4 Square planar Xe F F F F 90° WORKSHEET 2.1
VSEPR for five and six negative charge centres
VSEPR: Basic molecular configurations table
1 The valence-shell electron pair repulsion (VSEPR) theory is used to