1. Chapter 121 Chemical Bonding Chapter 12

2. 2Introduction The properties of many materials can be understood in terms of their microscopic properties. Microscopic properties of molecules include: the connectivity between atoms and the 3D shape of the molecule.

3. Chapter 123 We consider three bonds “within molecules” (intramolecular forces): ionic bond (electrostatic forces which hold ions together, e.g. NaCl) covalent bond (results from sharing electrons between atoms, e.g. Cl 2 ) metallic bonding (refers to metal nuclei floating in a sea of electrons, e.g. Na).

4. Chapter 124 In all chemical bonds, electrons are shared and transferred between atoms. In bonding, electrons are involved

5. Chapter 125 The electrons involved in bonding are called valence electrons Valence electrons are found in the incomplete, outermost orbital of an atom.

6. Chapter 126 Ionic Bonding Consider the reaction between sodium and chlorine: Na(s) + ½Cl 2 (g)  NaCl(s)

7. Chapter 127 Ionic Bonding Na(s) + ½Cl 2 (g)  NaCl(s)  H° f = -410.9 kJ The reaction is violently exothermic. We infer that the NaCl is more stable than its constituent elements. Why? Na has lost an electron to become Na + and chlorine has gained the electron to become Cl . Note: Na + has an Ne electron configuration and Cl  has an Ar configuration. That is, both Na + and Cl  have an octet of electrons surrounding the central ion.

8. Chapter 128 Ionic Bonding NaCl forms a very regular structure in which each Na + ion is surrounded by 6 Cl  ions. Similarly, each Cl  ion is surrounded by six Na + ions. There is a regular arrangement of Na + and Cl  in 3D. Note that the ions are packed as closely as possible. Note that it is not easy to find a molecular formula to describe the ionic lattice.

9. Chapter 129 Ionic Bonding

10. Chapter 1210 Covalent Bonding In ionic bonding one atom completely loses an electron while the other gains the electron. When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. When similar atoms bond, they share pairs of electrons to each obtain an octet. Each pair of shared electrons constitutes one chemical bond. Example: H + H  H 2 has electrons on a line connecting the two H nuclei.

11. Chapter 1211 Covalent Bonding

12. Chapter 1212 Covalent Bonding Multiple Bonds It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): One shared pair of electrons = single bond (e.g. H 2 ); Two shared pairs of electrons = double bond (e.g. O 2 ); Three shared pairs of electrons = triple bond (e.g. N 2 ). Generally, bond distances decrease as we move from single through double to triple bonds.

13. Chapter 1213 In a covalent bond, electrons are shared. Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. There are some covalent bonds in which the electrons are located closer to one atom than the other. Unequal sharing of electrons results in polar bonds.

14. Chapter 1214 Electronegativity The preference one atom in a chemical bond for electrons is called electronegativity. Electronegativity is a scale from 0.7 (Cs) to 4.0 (F). Electronegativity increases across a period up a group.

15. Chapter 1215 Electronegativity

16. Chapter 1216 Difference in electronegativity is a gauge of bond polarity: electronegativity differences of 0-0.2 result in non- polar covalent bonds (equal or almost equal sharing of electrons); electronegativity differences around 0.3-1.6 result in polar covalent bonds (unequal sharing of electrons); electronegativity differences between 1.7-4.0 result in ionic bonds (transfer of electrons). There is no sharp distinction between bonding types.

17. Chapter 1217 The positive end (or pole) in a polar bond is represented  + and the negative pole  -.

18. Chapter 1218 Lewis Symbols Also known as electron dot diagrams A way of keeping track of valence electrons. How to write them 1) Write the symbol. 2) Put one dot for each valence electron 3) Don’t pair up until they have to X

19. Chapter 1219 We generally place the electrons one four sides of a square around the element symbol. Octet rule: we know that s 2 p 6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs).

20. Chapter 1220 Lewis symbol for nitrogen Nitrogen has 5 valence electrons. First we write the symbol. Then add 1 electron at a time to each side. Until they are forced to pair up. The number of electrons available for bonding are indicated by unpaired dots Nitrogen would have 3 bonding electrons N

21. Chapter 1221 Drawing Lewis Structures of Molecules Add the valence electrons. Identify the central atom (usually the one with the highest molecular mass and closest to the center of the periodic table). Place the central atom in the center of the molecule and add all other atoms around it. Place one bond (two electrons) between each pair of atoms. Complete the octet for the central atom. Complete the octets for all other atoms. Use double bonds if necessary.

22. Chapter 1222 To draw the electron dot structure of an ion, you must add the charges to the number of electrons (for a negative ion) See p.342 in Heath

23. Chapter 1223 Formal Charge It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms. To determine which structure is most reasonable, we use formal charge. Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity. To calculate formal charge, electrons are assigned as follows: All nonbonding electrons are assigned to the atom on which they are found. Half the bonding electrons are assigned to each atom in a bond.

24. Chapter 1224 Drawing Lewis Structures Formal Charge Formal charge is: valence electrons - number of bonds - lone pair electrons Consider: For C: There are 4 valence electrons (from periodic table). In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge: 4 - 5 = -1.

25. Chapter 1225 Drawing Lewis Structures Formal Charge Consider: For N: There are 5 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge = 5 - 5 = 0. We write:

26. Chapter 1226 Drawing Lewis Structures Formal Charge The most stable structure has: the smallest formal charge on each atom, the most negative formal charge on the most electronegative atoms. Resonance Structures Some molecules are not well described by Lewis Structures. Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms.

27. Chapter 1227 More on Resonance Structures Check this out:

28. Chapter 1228 Drawing Lewis Structures Resonance Structures Example: experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter).

29. Chapter 1229 Drawing Lewis Structures Resonance Structures Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities.

30. Chapter 1230 Drawing Lewis Structures Resonance Structures Example: in ozone the extreme possibilities have one double and one single bond. The resonance structure has two identical bonds of intermediate character. Common examples: O 3, NO 3 -, SO 4 2-, NO 2, and benzene.

31. Chapter 1231 Drawing Lewis Structures Resonance in Benzene Benzene consists of 6 carbon atoms in a hexagon. Each C atom is attached to two other C atoms and one hydrogen atom. There are alternating double and single bonds between the C atoms. Experimentally, the C-C bonds in benzene are all the same length.

32. Chapter 1232 Drawing Lewis Structures Resonance in Benzene We write resonance structures for benzene in which there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring: Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor).

33. Chapter 1233 Drawing Lewis Structures Resonance in Benzene

34. Chapter 1234 Exceptions to the Octet Rule There are three classes of exceptions to the octet rule: Molecules with an odd number of electrons; Molecules in which one atom has less than an octet; Molecules in which one atom has more than an octet. 1) Odd Number of Electrons Few examples. Generally molecules such as ClO 2, NO, and NO 2 have an odd number of electrons.

35. Chapter 1235 NO has 5+6=11 valence electrons to place: Nitrogen ends up with less than an octet.

36. Chapter 1236 Exceptions to the Octet Rule Less than an Octet Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. Most typical example is BF 3. BF 3 has (7*3)+3=24 valence electrons. Boron ends up with less than an octet.

37. Chapter 1237 More than an Octet This is the largest class of exceptions. Atoms from the 3 rd period onwards can accommodate more than an octet.Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.