1 A molecule of ammonia NH 3 is made up of one nitrogen and three hydrogen atoms: Coordinate bond The nitrogen atom forms three bonds and the hydrogen atoms one bond each. In this case, one pair of electrons is not involved in bond formation and this is called a lone pair of electrons.

2 It is possible to have a shared electron pair in which the pair of electrons comes just from one atom and not from both. Such bond is called coordinate covalent bond. Even though the ammonia molecule has a stable configuration, it can react with hydrogen H + by donating the lone pair of electrons, forming the ammonium ion NH 4 + :

3 In the chlorine molecule Cl – Cl the pair of electrons of the covalent bond is shared equally between both chlorine atom. Because there is not a charge separation between the chlorine atoms, Cl 2 molecule is nonpolar. Partial ionic character of covalent bonds

4 On the contrary, in HCl molecule, there is a shift of electrons toward the chlorine atom which is more electronegative than hydrogen. Such molecule, in which a charge separation exists is called a polar molecule or dipole molecule. The polar molecule of hydrochloric acid + + e-e- H Cl d ++ --

5 The separation between the positive and negative charges is given by the dipole moment μ. The dipole moment is the product between the magnitude of the charges (δ) and the distance separating them (d): μ = δ · d The symbol δ suggests small magnitude of charge, less than the charge of an electron (1.602 · C). The unit for the dipole moment is Debye (D): 1D = 3.34 · C m

6 The charge δ for HCl molecule represents about 16% of the electron charge ( C). We can say that the covalent H – Cl bond has about 16% ionic character. Dipole moment values for some molecules: Carbon dioxide CO 2 μ = 0 D Carbon monoxide COμ = D Water H 2 Oμ = 1.85 D Hydrochloric acid HClμ = 1.03 D dist. between H and Cl atoms is d = 136 pm ( m) We can calculate the charge δ for HCl molecule:

7 c) Metallic bond The metallic bond represents the electromagnetic attraction forces between delocalized electrons and the metal nuclei. The metallic bond is a strong chemical bond, as indicated by the high melting and boiling points of metals. A metal can be regarded as a lattice of positive metal “ions” in a “sea” of delocalised electrons.

8 Metal atoms contain few electrons in their outer shells. Metals cannot form ionic or covalent bonds. Sodium has the electronic structure 1s 2 2s 2 2p 6 3s 1. When sodium atoms come together, the electron from the 3s atomic orbital of one sodium atom shares space with the corresponding electron of a neighbouring atom to form a molecular orbital. All the 3s orbitals of all the atoms overlap to give a vast number of molecular orbitals which extend over the whole piece of metal. There is a huge numbers of molecular orbitals because any orbital can only hold two electrons.

9 The electrons can move freely within these molecular orbitals and so each electron becomes detached from its parent atom. The electrons are called delocalized electrons. The “free“ electrons of the metal are responsible for the characteristic metallic properties: ability to conduct electricity and heat, malleability (ability to be flattened into sheets), ductility (ability to be drawn into wires) and lustrous appearance. Crystal structure of sodium

10 Weak chemical bonds – Intermolecular bonds 1. Van der Waals forces Intermolecular forces are attractions between one molecule and neighboring molecules. All molecules are under the influence of intermolecular attractions, although in some cases those attractions are very weak. These intermolecular interactions are known as van der Waals forces. Even in a gas like hydrogen (H 2 ), if the molecules are slow down by cooling the gas, the attractions become large enough so the molecules will stick together and form a liquid and then a solid.

11 The attractions between the H 2 molecules are so weak that the molecules have to be cooled to 21 K (-252  C) before the attractions are enough to form liquid hydrogen. Helium’ s intermolecular attractions are even weaker – the molecules won’t stick together to form a liquid until the temperature drops to 4 K ( -269  C). Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, it doesn’t seems to be any electrical distortion to produce positive or negative parts.

12 H 2 symmetrical molecule ++ - HH - But that’ s only true in average. In the next figure the symmetrical molecule of H 2 is represented. On average there is no electrical distortion.

13 But the electrons are mobile and at any one instant they might find themselves towards one end of the molecule. This end of the molecule becomes slightly negative (charge -  ). The other end will be temporarily short of electrons and so becomes slightly positive (+  ). An instant later the electrons may move to the other end, reversing the polarity of the temporary dipole of molecule. Temporary dipole of H 2 + ++ - HH

14 This phenomena even happens in monoatomic molecules of rare gases, like helium, which consists of a simple atom. If both the helium electrons happen to be on one side of the atom at the same time, the nucleus is no longer properly covered by electrons for that instant. Temporary dipole of He ++ --

15 How temporary dipoles give intermolecular bonds? A molecule which has a temporary polarity approaches another molecule which happens to be non-polar at that moment. The electrons form the non-polar molecule will be attracted by the slightly positive end of the polar molecule. This is how an induced dipole is forming. ++ -- ++ -- ++ -- original temporary dipole non-polar molecule induced dipole Dipole-dipole attraction (van der Waals forces)

16 An instant later the electrons in the left hand molecule can move up to the other end. So, they will repel the electrons in the right hand molecule. The polarity of both molecules reverses, but there is still attraction between  - end and  + end. As long as the molecules stay close to each other, the polarities will continue to fluctuate in synchronization so that the attraction is always maintained. This phenomena can occur over huge numbers of molecules. The following diagram shows how a whole lattice of molecules could be held together in a solid.

17 Molecular distribution in a solid The interactions between temporary dipoles and induced dipoles are known as van der Waals dispersion forces. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ --

18 Hydrogen bond If we plot the boiling points of the hydrides of the elements of groups 15, 16 end 17 we find that the boiling point of the first elements in each group is abnormally high.

19 In case of ammonia NH 3, water H 2 O and hydrofluoric acid HF, there must be some additional intermolecular forces of attraction, requiring significantly more heat energy to break them. These relatively powerful intermolecular forces are called hydrogen bonds. Hydrogen bonds are stronger than van der Waal dispersion forces, but weaker than covalent or ionic bonds. Hydrogen bonds can form if:  hydrogen is attached directly to one electronegative element (F, O, N)  each of the elements to which the hydrogen atom is attached have one “active“ lone pair of electrons.

20 Let’s consider two water molecules coming close together: The slightly  + charge of hydrogen is strongly attracted to the lone pair of electrons; as a result a coordinate bond is formed. This is a hydrogen bond. In liquid water, hydrogen bonds are constantly broken and reformed.

21 In solid water each water molecule can form hydrogen bonds with other 4 surrounding water molecules, creating a 3-D structure of ice. As a result, the boiling point of H 2 O is higher than that of NH 3 or HF.