1. C h a p t e rC h a p t e r C h a p t e rC h a p t e r 7 7 Covalent Bonding Chemistry 5th Edition McMurry/Fay Chemistry 5th Edition McMurry/Fay

2. 2 Electronegativity χ = (E i + E ea )/2 Usually converted to a unitless number between 0 and 4. Low χ → form cation High χ → form anion

3. 3 Electronegativity Bond polarity is due to electronegativity differences between atoms. Pauling Electronegativity: is expressed on a scale where F = 4.0

4. 4 Electronegativity Pauling Electronegativities

5. 5 Electronegativity Predicting χ 1. Increases across periods 2. Decreases down groups 3. Very low for noble gases.

6. 6 Electronegativity % Ionic Character: As a general rule for two atoms in a bond, we can calculate an electronegativity difference (∆EN ): ∆EN = EN(Y) – EN(X) for X–Y bond. If ∆EN < 0.5 the bond is covalent. If ∆EN < 2.0 the bond is polar covalent. If ∆EN > 2.0 the bond is ionic.

7. 7 The Covalent Bond Covalent bonds are formed by sharing at least one pair of electrons.

8. 8 The Covalent Bond Every covalent bond has a characteristic length that leads to maximum stability. This is the bond length.

9. 9 The Covalent Bond Every covalent bond has a characteristic length that leads to maximum stability. This is the bond length.

10. 10 Polar Covalent Bonds Using electronegativity values, predict whether the following compounds are nonpolar covalent, polar covalent, or ionic: SiCl 4 CsBrFeBr 3 CH 4 HClCCl 4 NH 3 H 2 O

11. 11 Electron-Dot Structures Using electron-dot (Lewis) structures, the valence electrons in an element are represented by dots. Valence electrons are those electrons with the highest principal quantum number (n). Valence electrons normally occupy s and p subshells

12. 12 Electron-Dot Structures Use the chemical symbol of the element, and add dots for each valence electron. N → [He] 2s 2 2p 3 ↑↓ ↑ ↑ ↑ 2s 2p N

13. 13 Electron-Dot Structures Octet: The s 2 p 6 configuration. The complete octet is the most stable configuration. Ar → [Ne] 3s 2 3p 6 ↑↓ ↑↓ 3s 3p Ar

14. 14 Electron-Dot Structures Ion Formation Elements may form ions in order to complete their octets. s-block and the lower left p-block elements will lose electrons; elements on the right of the p-block will gain electrons.

15. 15 Electron-Dot Structures Ionic Bonds: Two atoms form octets for each by one donating electron(s) to the other. Covalent Bonds: A pair of electrons is shared between two atoms, giving both an octet. (predicted by electronegativity)

16. 16 Electron-Dot Structures The electron-dot structures provide a simple, but useful, way of representing chemical reactions. Ionic: Covalent:

17. 17 Electron-Dot Structures Double Bond: Two atoms share two electron pairs, again forming octets for each. Triple Bond: Three electron pairs are shared.

18. 18 Electron-Dot Structures Single Bonds: Double Bonds: Triple Bonds:

19. 19 Drawing Lewis Structures Lewis Structure: A diagram showing how electrons are shared between atoms in a molecule. Octet Rule: Atoms proceed as far as possible toward completing their octets by sharing electron pairs. H needs only 2 electrons (and can accept no more).

20. 20 Drawing Lewis Structures Exceptions to the Octet Rule: 1. Boron may form an incomplete octet with only 6 valence electrons. 2. Phosphorus may form an extended octet with 10 valence electrons. 3. Sulfur may form an extended octet with 12 valence electrons.

21. 21 Drawing Lewis Structures Lone Pair: A pair of electrons residing on a single element and thus not involved in bonding. Bonding Pair: A pair of electrons shared in a covalent bond. Often represented as a line between atoms.F

22. 22 Drawing Lewis Structures Step 1: Determine the total number of valence electrons in the molecule. For anions add 1 electron for each negative charge. For cations subtract 1 electron for each positive charge. Step 2: Find the total number of electron pairs by dividing by 2.

23. 23 Drawing Lewis Structures Step 3: Identify the central atom(s) and arrange the other atoms around it (them). Central atoms are usually: C, N, B, P, S Step 4: Connect the atoms with single electron pairs first, then attempt to complete each octet by adding lone pairs or multiple bonds until all electrons are used.

24. 24 Drawing Lewis Structures