1. Chemistry 100 Chapter 8 Chemical Bonding Basic Concepts

2. Chemical Bonds  Three basic types of bonds Ionic  Electrostatic attraction between ions. Covalent  Sharing of electrons. Metallic  Metal atoms bonded to several other atoms.

3. 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  These two elements combine to form an ionic compound 2 Na (s) + Cl 2 (g)  2 NaCl (s)

4. What’s Happening? [Ne]3s 1 [Ne]3s 2 3p 5 (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

5. The NaCl Crystal

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

7. Stability of Ionic Compounds  The stability of ionic compounds depends on two main factors 1. The electron affinity of one of the elements 2. 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?

8. 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 ( lat H). The energy required to completely separate one mole of the solid ionic compound into its gas-phase ions.

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

10. 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. Covalent bonds H 2  H-H, F 2  F-F, Cl 2  Cl-Cl  For many electron atoms (like F and Cl), we again to worry only about the outermost (valence) electrons.

11. Covalent Bonding  In covalent bonds, atoms share electrons.  There are several electrostatic interactions in these bonds: Attractions between electrons and nuclei, Repulsions between electrons, Repulsions between nuclei.

12. Examples of Covalent Bonding  Let’s look at the Cl 2 example.  Each Cl atom has 7 valence shell electrons 3 Lone pairs and one unpaired electron Lone pairs Unshared electron

13. The Cl 2 Molecule  The structure we have just drawn are called Lewis structures.  The dash between the atomic centres represents bonding electrons  Redraw F 2 lone pairs (non bonding) bonding electrons

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

15.  check out N 2  :NN: (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.

16. Covalent Compounds  Compounds that contain only covalent bonds are called covalent compounds.  There are two main of covalent compounds, Molecular covalent compounds (CO 2, C 2 H 4 ) Network covalent compounds (SiO 2, BeCl 2 ). The network covalent compound are characterized by an extensive “3-D” network bonding

17. 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 CCl 4

18. 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.

19. 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

20.  Electronegativity is related to the electron affinity and the ionization energy.  Compare the following elements. Na  low I 1, small negative E.A.  low  F  high I 1, large, negative E.A.,  high 

21. 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

22. Plot of  Values Plot of  Values

23. Electronegativity and Bond Type  Can we use the electronegativity values to help us deduce the type of bonding in compounds?  values bond type 0.0 <  < 0.5 non-polar covalent 0.5    1.9 polar covalent 2.0    3.3 Ionic bond

24. 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.

25.  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).

26.  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.

27. 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?

28. Rules for Formal Charges  Neutral molecules   formal charges = 0  Ions  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

29. Resonance This is the Lewis structure we would draw for ozone, O 3.

30. Resonance  Note the true, observed structure of ozone… …both O—O bonds are the same length. …both outer oxygens have a charge of −1/2.

31. Resonance  One Lewis structure cannot accurately depict a molecule like ozone.  We use multiple structures, resonance structures, to describe the molecule.

32. Resonance Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.

33. Experimental Evidence for Resonance.  The resonance structures for benzene C 6 H 6  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

34. Exceptions to the Filled Valence Shell Rule  Be compounds  BeH 2, BeCl 2,  Boron and Al compounds  BF 3, AlCl 3, BCl 3  BF 3 is stable  The B central atom has a tendency to pick up an unshared e- pair from another compound BF 3 + NH 3  BF 3 NH 3  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.

35. Odd e- molecules  These molecules have uneven numbers of electrons  no way that they can form octets.  Examples NO and NO 2. These species have an odd number of electrons.

36.  Look at the dimerization reaction of NO 2. 2 NO 2 (g) ⇄ N 2 O 4 (g) K eq = 210

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

38.  Look at HClO 3 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.

39.  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 Note: the final three structures reduce the formal charges

40. Bond Energies and Thermochemistry  Look at the energy required to break 1 mole of gaseous diatomic molecules into their constituent gaseous atoms. H 2 (g)  H (g) + H (g)H° = 436.4 kJ Cl 2 (g)  Cl (g) + Cl (g) H° = 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).

41. For Polyatomic Molecules. CO 2 (g)  C (g) + 2 O (g)H = 745 kJ  Denote the H of this reaction D(C=O)  What about dissociating methane into C + 4 H’s? CH 4 (g)  C(g) + 4 H (g) H° = 1650 kJ  Note 4 C-H bonds in CH 4  D (C-H) = 412 kJ/mol

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

43.  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.