Chapter 8 Concepts of Chemical Bonding (8-5 to 8-8)

1. Chapter 8Concepts of Chemical Bonding(8-5 to 8-8)

2. 8.5 Drawing Lewis Structures • Lewis structures are representations of covalent bonding in molecules. Though simple, they are useful for better understanding of the bonds between atoms in many molecules. • Follow the procedures below. PCl3 • Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. • If it is an anion, add one electron for each negative charge. • If it is a cation, subtract one electron for each positive charge. 5 + 3(7) = 26

3. The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds of shared electrons. Each single bond contains two electrons. Keep track of the electrons: 26  6 = 20

4. Complete the octets of the outer atoms with pairs of unshared electrons. Keep track of the electrons: 26  6 = 20  18 = 2

5. Complete the octet of the central atom with unshared (lone) electrons pairs. • Place any remaining electrons around the central atom even though it violates the octet rule. The octet rule does not always apply for certain atoms. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

6. If you run out of electrons before the central atom has an octet, form multiple bonds by taking unshared electrons and making them shared (bonding) electrons. Remember, shared electrons count towards fulfilling the octet for both bonded atoms. In the above structure, both C and N share the electrons in the triple bond between them to achieve an octet of electrons. The single pair between C and H satisfies the requirement for H to have a duet of electrons. Also remember that H can only have a duet (one shared pair) of electrons. It can never have more than one shared bond and can, therefore, never be a central atom.

7. Lewis Structures for Polyatomic Ions • Add one electron for each negative charge • Subtract one electron for each positive charge • Draw the structure following the rules for neutral compounds. • Put the Lewis structure in brackets with the charge outside the brackets at the upper right. • See sample exercise 8.8 and the practice exercise

8. Formal Charges • Sometimes, it is possible to draw two or more Lewis structures for a particular molecule or ion. • Calculating formal charges for each atom in a molecule allows you to determine which structure is the most reasonable. • The formal charge is the number of valence electrons in an isolated atom minus the number of electrons assigned to it in a Lewis structure. • The formal charge for an atom in a molecule is the charge the atom would have if all the atoms in the molecule had the same electronegativity (i.e., they were equally shared). • Formal charges do not represent real charges on atoms. They are “bookkeeping” conventions.

9. Assigning formal charges • For each atom, count the electrons in lone pairs and halve the electrons it shares with other atoms. • Subtract that from the number of valence electrons for that atom: The difference is its formal charge.

10. The best Lewis structure… • …is the one with the fewest charges. • …puts a negative charge on the most electronegative atom. • Formal charge is a tool for estimating the distribution of electric charge within a molecule. • The concept of oxidation states constitutes a competing method to assess the distribution of electrons in molecules. • If the formal charges and oxidation states of the atoms in carbon are compared, the following values are arrived at:

11. Neither oxidation number nor formal charge gives an accurate description of actual charges on atoms. • Oxidation numbers assign all shared electrons to the most electronegative atom, whereas formal charges divides all shared electrons equally between atoms. • The images below shows the difference. The image on the far right shows the calculated distribution of electrons in HCl. Regions of more negative charge are red.

12. 8.6 Resonance Structures • Sometimes the experimentally determined arrangement of atoms is not adequately described by a single Lewis structure • This is the Lewis structure we would draw for ozone, O3. • We would expect the double bond to be shorter than the single bond. • The negative charge is assigned to one of the oxygens. + -

13. But this is at odds with the true, observed structure of ozone, in which… • …both O—O bonds are the same length. • …both outer oxygens have a charge of 1/2.

14. One Lewis structure cannot accurately depict a molecule such as ozone. • The double bond could just as easily been put between the other O=O bond. • We use multiple structures, called resonance structures, to describe the molecule.

15. Just as green is a synthesis of blue and yellow, ozone is a synthesis of these two resonance structures. • The real molecule is described as an average, or blend, of the two Lewis structures. • Ozone cannot be depicted by a single Lewis structure in which the electrons are locked into position. • A single Lewis structure would have to show the electrons not localized, as shown to the left, but delocalized, i.e., spread over the two bonds.

16. Consider the Lewis structures for the formate ion, below • In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. • They are not localized, but rather are delocalized.

17. The organic compound benzene, C6H6, has two resonance structures. • It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

18. 8.7 Exceptions to the Octet Rule • There are three types of ions or molecules that do not follow the octet rule: • Ions or molecules with an odd number of electrons. • Ions or molecules with less than an octet. • Ions or molecules with more than eight valence electrons (an expanded octet).

19. Odd Number of Electrons • Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

20. Less than an Octet of Valence Electrons • Consider BF3: • Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. • This would not be an accurate picture of the distribution of electrons in BF3.

21. Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

22. If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

23. More than an Octet of Valence Electrons • The only way PCl5 can exist is if phosphorus has 10 electrons around it. • It is allowed to expand the octet of atoms on the 3rd row or below. • Presumably d orbitals in these atoms participate in bonding.

24. Sometimes an expanded valence shell is shown even if a structure can be drawn with an octet. • Consider the case with the PO43- ion. • Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, formal charges show that the better structure puts a double bond between the phosphorus and one of the oxygens, thus eliminating the charge on the phosphorus and one of the oxygens. • The best representation is a series of Lewis structures in resonance.

25. 8.8 Covalent Bond Strength • The strength of a bond is measured by determining how much energy is required to break the bond. • The bond enthalpy is the enthalpy change, ΔH, for the breaking of a particular bond in one mole of a gaseous substance. • The bond enthalpy for a Cl—Cl bond, D(Cl—Cl), is measured to be 242 kJ/mol. • This is the enthalpy change when 1 mol of Cl2 is dissociated into Cl atoms.

26. Average Bond Enthalpies • This table lists the average bond enthalpies for many different types of bonds. • Average bond enthalpies are positive, because bond breaking is an endothermic process. • These are average bond enthalpies, not absolute bond enthalpies. • The C—H bonds in methane, CH4, will be a bit different than the C—H bond in chloroform, CHCl3.

27. Enthalpies of Reaction • Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed. • In other words, Hrxn = (bond enthalpies of bonds broken)  (bond enthalpies of bonds formed)

28. CH4(g) + Cl2(g)  CH3Cl(g) + HCl(g) • In this example, one C—H bond and one Cl—Cl bond are broken. • One C—Cl and one H—Cl bond are formed.

29. So, Hrxn = [D(C—H) + D(Cl—Cl)  [D(C—Cl) + D(H—Cl) = [(413 kJ) + (242 kJ)]  [(328 kJ) + (431 kJ)] = (655 kJ)  (759 kJ) = 104 kJ

30. Bond Enthalpy and Bond Length • We can also measure an average bond length for different bond types. • As the number of bonds between two atoms increases, the bond length decreases.