Covalent Chemical Bonds

1. Covalent Chemical Bonds • The electronic structures of atoms and molecules have many features in common. Individual atoms often possess unpaired electrons. These atoms are usually chemically unstable. Two such atoms can come together to form a molecule with no unpaired electrons. This process can involve the formation of covalent chemical bonds and is highly exothermic.

2. Covalent Chemical Bonds • In isolated atoms most of the electrons will be found in pairs in a number of different atomic orbitals. When atoms combine to form molecules, valence shell electrons are rearranged. Two electrons from different atoms can “pair up” to form a single covalent bond where the bonding molecular orbital is associated with more than one atom.

3. Covalent and Ionic Bonds • Before moving on to Lewis structures lets practice (review) looking at chemical formulas and identifying whether covalent bonds, ionic bonds or bothare important for the compounds under consideration.

4. Covalent and Ionic Bonds

5. Covalent Chemical Bonds • We will consider molecular orbitals in more detail shortly. Before doing so we will use Lewis electron dot structures to represent bonding in both covalent and ionic compounds. In these structures valence shell electrons are represented by dots. The rest of the atom – the nucleus and the core (non-valence) electrons are represented using the chemical symbol for the element.

6. Lewis Theory: An Overview • Valence e- play a fundamental role in chemical bonding. • e- transfer leads to ionic bonds. • Sharing of e- leads to a covalent bond. • e- are transferred or shared to give each atom a noble gas configuration, the octet. Gilbert Newton Lewis (1875-1946) General Chemistry: Chapter 10

7. •• •• •• •• •• As Bi P Sb N • • • • • • • • • • • • • • • •• •• •• Al Ar • Se •• I • •• • • •• • • •• • •• Lewis Symbols and Lewis Structures • Si • • • A chemical symbol represents the nucleus and the coree-. Dots around the symbol represent valence e-. General Chemistry: Chapter 10

8. Lewis structures General Chemistry: Chapter 10

9. Covalent Bonding: An Introduction General Chemistry: Chapter 10

10. + H H - •• N H Cl •• •• H N H •• •• H H Coordinate Covalent Bonds Cl H FIGURE 10-2 • Formation of the ammonium ion, NH4+ General Chemistry: Chapter 10

11. Multiple Covalent Bonds • In most cases Main Group elements are surrounded by eight valence shell electrons when molecules are formed – octet rule (H is an impt exception!). Such molecules can contain single bonds (one pair of bonding electrons), double bonds (two pairs of bonding electrons and triple bonds as well as lone pairs of electrons.

12. Lewis Structures • Class Examples: Construct Lewis structures for H2, F2,O2, N2, CO and CN-. In which species are multiple bonds required to satisfy the octet rule? What type(s) of experimental evidence would suggest that the Lewis structures for these species have physical meaning?

13. Electron Dot Formulas -Shortcomings • Electron dot structures account for the fact that the covalent bond in hydrogen (single pair of bonding electrons) is weaker than the bond in oxygen (two pairs of bonding electrons) which, in turn, is weaker than the bond in nitrogen (three pairs of bonding electrons). For oxygen, the expected Lewis structure does not predict that the O2 molecule is paramagnetic. The magnetic properties of O2(l) are easily demonstrated.

14. Paramagnetism of Oxygen Paramagnetism in O2 will be explained when we consider molecular orbitals. General Chemistry: Chapter 10

15. Molecules Are Often Polar • Lewis structures also do not account satisfactorily for the fact that covalently bonded molecules can be electrically polar. Electrical polarity is important for many reasons including the fact that polar and nonpolar molecules have very different physical properties. In polar molecules the centers of positive charge and negative charge do not coincide. As a result, polar molecules have a non-zero electric dipole moment.

16. Molecules Can Be Polar – cont’d: • Molecules with electrical polarity have a tendency to orient in an electric field. The size of the electric dipole moment can be measured by studying rotational energies of molecules (gas phase) perturbed by an electric field (higher level courses). Further insight into the electronic structure of molecules can be gained using electrostatic potential maps which reflect Coulomb’s Law but plot electrical work rather than Coulombic force.

17. Electrostatic Potential Maps • We will use electrostatic potential maps in a qualitative way to get some insight into the electron distribution in covalently bonded molecules. In the text diagrams (for relatively simple molecules) the colour blue corresponds to a “positively charged part” of a molecule and red to a “negatively charged” part of the molecule.

18. Polar Covalent Bonds and Electrostatic Potential Maps FIGURE 10-4 • Determination of the electrostatic potential map for ammonia General Chemistry: Chapter 10

19. The electrostatic potential maps for sodium chloride, hydrogen chloride and chlorine FIGURE 10-5 General Chemistry: Chapter 10

20. Electrostatic Potential Maps – cont’d: • The electrostatic map for the one homonuclear diatomic molecule, Cl2, shows little colour variation and, as well, symmetry. The left hand side of the map is the mirror image of the right hand side implying that electrons in the Cl2 molecule are equally shared between the two Cl atoms comprising the molecule. For HCl we see a greater range of colours implying increased molecular polarity.

21. Electrostatic Potential Maps – cont’d: • We see a large range of colours for the NaCl molecule implying extreme electrical polarity. The maps also show that the Na and Cl ends of the NaCl molecule are, respectively, positively and negatively charged. It will be important to distinguish between a sodium chloride crystal where the bonding is pure ionic and NaCl molecules in the gas phase where the bonding has significant covalent character.

22. Aside: Gas Phase Metal/Nonmetal Molecules • At very high temperatures even compounds with high melting points have a small vapour pressure. One can therefore get NaCl molecules in the gas phase at high temperature. Such molecules can also be made using laser ablation experiments where pulses of light from a high frequency laser can vaporize small amounts of metal which can react with gaseous nonmetals to form, e.g., NaCl(g).

23. Lewis Structures - Class Examples: • 1. Draw a Lewis structure for BrF. Does this molecule have a single, double or triple covalent bond? Why? Write chemical formulas for five other molecule having similar Lewis structures. • 2. Draw Lewis structures for CN- and for HCN. What is the strongest bond in the HCN molecule? How can a proton react with CN- to form a HCN molecule?

24. Lewis Structures - Class Examples: • 3. Draw Lewis structures for the CS2 and OCS molecules (central S atom). Do these structures obey the “octet rule”? Write chemical formulas for three other molecules with similar Lewis structures. Do these molecules have any common structural features? • 4. Draw Lewis structures for H2O, HF and NH3 and CH4. Are there common features and important differences here?