Chemistry 100 Chapter 8

1. Chemistry 100 Chapter 8 Chemical Bonding Basic Concepts

2. The Valance Electrons When atoms interact to form chemical bonds, only the outer (valance) electrons take part. Need a tool for keeping track of valence electrons, e.g., The Lewis dot symbol 1 v.E. 7 v.E’s When these two elements combine to form a compound 2 Na (s) + Cl2 (g) ® 2 NaCl (s)

3. What’s Happening? [Ne]3s1 [Ne]3s23p5 (g) ® Na+ (g) + e- (ionizes, loses e-) • an electron configuration of [Ne] • (g) + e- ® Cl- (g) • an electron configuration of [Ar] • In the crystal lattice, • Na+ and Cl- ions; strong electrostatic attractions

4. The NaCl Crystal

5. Ionic Bonding • q1 magnitude of charge 1 • q2 magnitude of charge 2 • r  distance between the ionic • centres Electrostatic attractions that hold ions together in an ionic compound. The strength of interaction depends on charge magnitude and distance between them.

6. Stability of Ionic Compounds • The stability of ionic compounds depends on two main factors • The electron affinity of one of the elements • The ionization energy of the other • Note • electron affinities and ionization potentials are gas-phase reactions? • How are they related to the stability of solid materials?

7. The Lattice Energy • A quantitative measure of just how strong the interaction is between the ionic centres (i.e., a measure of the strength of the ionic bond) • For the reaction KCl (s)  K+ (g) + Cl- (g) H = 718 kJ/mol • Lattice energy (DlatH). • The energy required to completely separate one mole of the solid ionic compound into its gas-phase ions.

8. Lattice Energies of Various Ionic Compounds Determined using a thermochemical cycle - the Born-Haber cycle (a Hess’s Law application)

9. Covalent Bonding In a wide variety of molecules, the bonding atoms fulfill their valance shell requirements by sharing electrons between them. Covalent bonds - a bond in which the electrons are shared by two atoms. H2 ® H-H, F2® F-F, Cl2 Cl-Cl For many electron atoms (like F and Cl), we again to worry only about the outermost (valence) electrons.

10. Covalent Bonding

11. Examples of Covalent Bonding Lone pairs Unshared electron Let’s look at the Cl2 example. Each Cl atom has 7 valence shell electrons 3 Lone pairs and one unpaired electron

12. The Cl2 Molecule lone pairs (non bonding) bonding electrons The structure we have just drawn are called Lewis structures. The dash between the atomic centres represents bonding electrons Redraw F2

13. Valence shell requirements are satisfied by the formation of a double bond. Note both Cl2 and F2 satisfy their valence shell requirements by the formation of a single bond. What about O2? How can we satisfy the octet rule for 2 O atoms?

14. check out N2  :NºN: (triple bond) • Note that the octet rule works mainly for the second row elements. • Filled valence shells can have more than 8 electrons after Z=14 (Si). This is generally termed octet expansion.

15. Covalent Compounds • Compounds that contain only covalent bonds are called covalent compounds. • There are two main of covalent compounds, • Molecular covalent compounds (CO2, C2H4) • Network covalent compounds (SiO2, BeCl2). • The network covalent compound are characterized by an extensive “3-D” network bonding

16. Comparison between Ionic and Covalent Compounds • Ionic Compounds • usually solids with very high melting points • conduct electricity when molten (melted) • usually quite water soluble and they are electrolytes in aqueous solution • NaCl • Covalent Compounds • usually low melting solids, gases or liquids • don’t conduct electricity when molten • aren’t very soluble in water and are non electrolytes • CCl4

17. The Filled Valence Shell rule • Filled Valence Shell rule • Atoms participate in the formation of bonds (either ionic or covalent) in order to satisfy their valence shell requirements. • Atoms other than H tend to form bonds until they end up being surrounded by 8 valence electrons (the noble gas configuration). • Your text calls this the “octet” rule.

18. Electronegativity • Electronegativity is defined as the ability of an atom to attract electrons towards itself in a molecule ( (pronounced ‘chi’)) • Examine the H-F covalent bond +H-F •  denotes a partial “+” charge on the H atom • - denotes a partial “-“ charge on F atom

19. Electronegativity is related to the electron affinity and the ionization energy. • Compare the following elements. • Na  low I1, small negative E.A.  low  • F  high I1, large, negative E.A.,  high 

20. Trends in the  Values • Across a row • The  values generally increase as we proceed from left to right in the periodic table. • Down a group • The  values generally decrease as we descend the group. • Transition metals • Essentially constant  values

21. Plot of  Values

22. Electronegativity and Bond Type Can we use the electronegativity values to help us deduce the type of bonding in compounds?

23. An Outline for Drawing Lewis Structures • Predict arrangement of atoms (i.e., predict the skeletal arrangement of the molecule or ion). • The H is always a terminal atom, bonded to ONE OTHER ATOM ONLY. A halogen atom is usually a terminal atom. • Note that the central atom usually has the least negative electron affinity.

24. Count total number of valence shell electrons (include ionic charges). Place 1 pair electrons (sigma bond, ) between each pair of bonded atoms (i.e., the central atom and each one of the terminal atoms). Place remaining electrons around the terminal atoms to satisfy the filled valence shell rule. (lone pairs).

25. All remaining electrons are assigned to the central atom. Atoms in the 3rd or higher row can have more than eight electrons around them. • If a central atom does not have a filled valence shell, use a lone pair of electrons from a terminal atom to make a pi () bond.

26. Formal Charges • Definition: formal charge on atom = number of valence electrons – number of non-bonding - ½ the number of bonding electrons. • Formal charge in a Lewis Structure is a bookkeeping “device” • keeps track of the electrons “associated” with certain atoms in the molecule vs. the valence e-‘s in the isolated atom! • How does it work?

27. Rules for Formal Charges • Neutral molecules ®S formal charges = 0 • Ions ® S formal charges = charge of ion • For molecules where the possibility of multiple Lewis Structures with different formal charges exist • Neutral molecule - choose the structure with the fewest formal charges. • Structures with large formal charges are less likely than ones with small formal charges • Two Lewis Structures with similar formal charge distribution ® negative formal charges on more electronegative atom

28. Resonance Structures • The structures differ in the location of the N=O double bond. • They are said to be resonance structures. • The actual structure of the molecule is a combination of three resonance structures (the resonance hybrid). Examine the NO3- anion.

29. Experimental Evidence for Resonance. • We would expect to find two different bond lengths in benzene (C=C and C-C bonds). • C= C ® bond length = 133 pm = 0.133 nm • C-C ® bond length = 0.154 nm • Experimentally, all benzene carbon-carbon bond lengths are equivalent at 0.140 nm The resonance structures for benzene C6H6

30. Exceptions to the Filled Valence Shell Rule Be compounds  BeH2, BeCl2, Boron and Al compounds  BF3, AlCl3, BCl3 BF3 is stable Þ The B central atom has a tendency to pick up an unshared e- pair from another compound BF3 + NH3® BF3NH3 the B-N bond is an example of a coordinate covalent bond, or a “dative” bond ® i.e. a bond in which one of the atoms donates both bonding electrons.

31. Odd e- molecules • These molecules have uneven numbers of electrons \ no way that they can form octets. • Examples • NO and NO2. These species have an odd number of electrons.

32. Look at the dimerization reaction of NO2. 2 NO2 (g) ⇄ N2O4 (g) Keq = 210

33. Valence Shells having more than 8 Electrons (Expanded Octets) Reason - elements in this category can use the energetically low-lying d orbitals to accommodate extra electrons A central atom having more than 8 valance shell electrons is possible with atomic number 14 and above.

34. High formal charge on the electronegative Cl atom (f.c.(Cl) = 7-2-1/2 (6) = +2) • This resonance structure would make a very small contribution to the overall resonance hybrid. Look at HClO3

35. Note: the final three structures reduce the formal charges With the possibility of using the low lying d-orbitals on the Cl atom to accommodate extra electron pairs, we may write other Lewis structures

36. Bond Energies and Thermochemistry Look at the energy required to break 1 mole of gaseous diatomic molecules into their constituent gaseous atoms. H2 (g) ® H (g) + H (g) DH° = 436.4 kJ Cl2 (g) ® Cl (g) + Cl (g) DH° = 242 kJ These enthalpy changes are called bond dissociation energies. In the above examples, the enthalpy changes are designated D (H-H) and D (Cl-Cl).

37. For Polyatomic Molecules. CO2 (g) ® C (g) + 2 O (g) DH = 745 kJ Denote the DH of this reaction D(C=O) What about dissociating methane into C + 4 H’s? CH4 (g) ® C(g) + 4 H (g) DH° = 1650 kJ Note 4 C-H bonds in CH4\ D (C-H) = 412 kJ/mol

38. H2O (g) ® 2 H (g) + O (g) DH° = 929 kJ/mol H2O • It takes more energy to break the first O-H bond. H2O (g) ® H (g) + OH (g) DH° = 502 kJ/mol H2O HO (g) ® H (g) + O (g) DH= 427 kJ/mol H2O • Note: we realize that all chemical reactions involve the breaking and reforming of chemical bonds. • Break bonds  add energy. • Make bonds  energy is released. • rxnH° S D(bonds broken) - S D(bonds formed)

39. These are close but not quite exact. Why? The bond energies we use are averaged bond energies, i.e., This is a good approximate for equations involving diatomic species. We can only use the above procedure for GAS PHASE REACTIONS ONLY.