Covalent Bonding 3

1. Covalent Bonding 3 Molecular Shape

2. Molecular Shape • Lewis diagrams are of use when we are learning what a covalent bond is, but the most useful representation of molecules in the structural formula. • In this representation each pair of electrons, bonding and non-bonding pairs alike, is shown as a simple line. • Non-bonding pairs can also be shown as two dots.

3. Molecular Shape • The actual shape of the molecule is also shown. • This shape has an important part to play in determining the chemical and physical properties of a molecule.

4. Molecular Shape • Each pair of electrons will be repelled from the others as far as possible in three-dimensional space because electrons all carry a negative charge. • This is known as the valence electron pair repulsion or VSEPR theory. • This theory states that the electron pairs around an atom repel each other.

5. Molecular Shape • The electrostatic repulsion of pairs of electrons determines the geometry of the atoms in the molecule. • The non-bonding pairs and bonding pairs of electrons are arranged around the central atom so as to minimize this electrostatic repulsion between the non-bonding and bonding pairs of electrons.

6. Shape of Molecule • The relative magnitude of the electron pair repulsion is: Non-bonding pair-non-bonding pair > bonding pair-non-bonding pair > bonding pair-bonding pair • A non-bonding pair will spread out more than a bonding pair and, therefore, the repulsion will be greatest between non-bonding pairs.

7. Molecular Shape • While non-bonding pairs of electrons can be important in determining the overall shape of a molecule, they are no actually considered part of the shape. • Shape describes the position of the atoms only; however, non-bonding electrons repel other pairs of electrons and so influence the final shape of a molecule.

8. Molecular Shape • Negative charge centre or region refers to pairs of electrons on the central atom. • This includes both the non-bonding pairs and bonding pairs of electrons in single, double or triple bonds. • Each double or triple bond is counted as one negative charge centre.

9. Molecular Shape • For central atoms that have an octet of electrons in the valence level, you can classify most shapes into five categories.

10. Linear • A combination of two double bonds or a single bond and a triple bond results in two negative charge centres. • This results in a linear shape. • Examples • Carbon dioxide (CO2) • Hydrogen cyanide (HCN) • The bond angle, the angle formed by two bonds, in linear molecules is 180o.

11. Carbon dioxide (CO2)

12. Trigonal Planar • Three negative charge centres around a central atom can be achieved by a combination of two single bonds and one double bond resulting in a trigonal planar shape. • Example • Methanal (CH2O) • Boron trifluoride (BF3) • All three bond angles are 120o.

13. Boron trifluoride (BF3) • There are some exceptions to the octet rule. • Boron is one of these exceptions. It is a metalloid in group 3 and it has 3 valence electrons. • It forms three covalent bonds and ends up with 6 electrons.

14. Tetrahedral • The most common situation encountered in molecules is the existence of four pairs of electrons, either bonding or non-bonding, surrounding each atom. • The most widely spaced arrangement of four pairs of electrons in three-dimensional space is known as the tetrahedral arrangement, in which each atom can be imagined at the vertex of a regular triangular-based pyramid. • The bond angle in this arrangement is 109.5o.

15. Methane

16. Molecular Shape • Shapes involving only bonding pairs are very symmetrical because repulsive forces between all electron groupings are the same. • When non-bonding pairs are introduced into a shape, the forces between electron groupings are no longer the same resulting in a distortion of the symmetry.

17. Pyramidal • One non-bonding pair and three single bonds result in a pyramidal shape. • The non-bonding pair exerts a stronger force on the bonded pairs than they do on each other, reducing the bond angle to less than 109.5ofound in the tetrahedral shape. • It is not possible to predict exactly how much the angle will be changed because the result depends on the properties of the elements and the bonds between the atoms.

18. Ammonia (NH3)

19. Bent • When the central atom has two non-bonding pairs and two single bonds, the resulting shape is bent. • The two non-bonding pairs exert greater force on each other than on the bonding pairs and reduce the angle between the two bonds even more than in the pyramidal shape.

20. Water (H2O)

22. Predicting Molecular Shape • Draw a preliminary Lewis structure of the molecule based on the formula given. • Determine the total number of negative charge centres around the central atom. Remember that a double or triple bond is counted as one charge centre. • Determine the types of charge centres (bonding or non-bonding pairs). • Determine which one of the five shapes will accommodate this combination of charge centres.

23. The Ammonium Ion (NH4+) • The ammonium ion is formed when a hydrogen ion bonds to the non-bonding pair of an ammonia molecule. • Although the bond between the H+ and the nitrogen is a covalent bond, it is formed in a slightly different way from the other covalent bonds in this molecule. • The H+ ion has no electrons to share in the bond, so both bonding electrons come from the nitrogen atom.

24. The Ammonium Ion (NH4+) • The H+ ion has no electrons to share in the bond, so both bonding electrons come from the nitrogen atom. • This is shown by the arrow in the structural formula.

25. Dative Covalent Bonds • A covalent bond which is formed in this way is called a dative covalent bond. • You might also see this called co-ordinate covalent bond. • A dative covalent bond behaves in the same way as any other single covalent bond.

26. The Ammonium Ion (NH4+) • In the ammonium ion, the nitrogen is surrounded by four bonding pairs of electrons, rather than three bonding and one non-bonding, so the ammonium ion has the same shape as a methane molecule—it is tetrahedral, with bond angles of 109.5o.

27. Sulfur dioxide (SO2) • The structure of sulfur dioxide is more unusual than those we have seen so far. • The central atom, sulfur, has 6 valence electrons. • Each oxygen atom has 6 electrons and needs to make two bonds to fill its valence shell. • This Lewis structure suggests that SO2 contains a single bond and a double bond.

28. Sulfur dioxide (SO2) • However, experimental measurements of bond lengths indicate that the bonds between the S and each O are identical. • The two bonds have properties that are somewhere between a single and a double bond. • In effect, the SO2 molecule contains two “one-and-a-half” bonds.

29. Sulfur dioxide (SO2) • To communicate the bonding in SO2 more accurately, chemists draw two different resonance Lewis structures side by side separated by a double-headed arrow. • Resonance structures are models that give the same relative position of atoms as in Lewis structures, but show different placing of their bonding and non-bonding pairs.

30. Resonance Structures • Many molecules and ions require resonance structures to represent their bonding. • It is important to keep in mind that resonance structures do not exist in reality. • An actual SO2 molecule is a combination—a hybrid—of its two resonance structures.

31. Practice Problems • Work on practice problems 6 & 7 on page 193 of the textbook.