1. Chapter 8: Covalent Bonding Vocabulary: Leave enough space for definition and example 1.Covalent bond 2.Electron dot structure 3.Diatomic Molecules 4.Polar covalent bond 5.Polar and Non Polar Molecules 6.Chemical Bonding 7.Intermolecular forces 8.Intra-molecular forces 9.Dispersion Forces 10.Hydrogen Bond

2. The goal of all Chemical Bonds Elements bond to have a full outer shell like noble gases. Ionic bonds: lose and gain electrons Covalent bonds: share electrons

3. Covalent Bonds Occur between two nonmetals Molecular compounds have covalent bonds. A bond in which two atoms share ONE pair of electrons between them is a SINGLE COVALENT BOND

4. Hydrogen gas (H 2 ) Hydrogen gas wants to look like Helium. To do this it needs to get one more electron.

5. Diatomic Molecules Atoms that exist as a pair in nature There are 7 diatomic elements MEMORIZE!

6. Hydrogen = H 2 Nitrogen = N 2 Oxygen = O 2 Fluorine = F 2 Chlorine = Cl 2 Bromine = Br 2 Iodine = I 2 Hydrogen gas Nitrogen gas Oxygen gas Fluorine gas Chlorine gas Bromine Iodine

7. Naming Covalent Use prefixes to distinguish the number of each atom present in the compound Example: C 2 H 4 = dicarbon tetra hydride You still add the ending –ide to the last element

8. Prefixes 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca

9. Common Names for some… Water = H 2 O Ammonia = NH 3 Methane = CH 4 Ozone = O 3

10. How to draw Lewis Dot Structures 1.Determine the number of valence electrons for each atom by using the periodic table. 2.Calculate the total number of electrons 1.Be sure to multiply the number of valence electrons if there is more than one atom of the element in a compound 3.Write the symbols for each atom and the correct amount for each element

11. What order do I write the symbols in for a molecule? a.Generally the first element in the formula goes in the center. b.Good rule of thumb – the least electronegative element goes in the center. c.Some elements NEVER go in the center (like: H, halogens)

12. Remember… Each element only needs 8 electrons 1 dot = 1 electron Only two dots per side * Remember, H and He are exceptions, they can have just 2 dots

13. 4.Once you have the symbols written, draw a solid line between each element. This represents the SINGLE COVALENT BOND. 5.For each line, subtract 2 from your valence electron total. a.Each line = 2 electrons

14. 6. Distribute remaining dots remembering to satisfy the octet rule (8 electrons per symbol) `7. If you run out of dots, you might need to make a double or triple bond. Double bonds: atoms share TWO PAIRS of electrons. (4 electrons total) Triple Bonds: atoms share THREE pairs of electrons. (6 electrons total)

15. Double and Triple Covalent Bonds Double bonds: atoms share TWO PAIRS of electrons. (4 electrons total) Triple Bonds: atoms share THREE pairs of electrons. (6 electrons total)

16. N e.x. Nitrogen atom Examples of Lewis Dot diagrams S e.x. Sulfur atom Electron orbitals electron

17. study question 1 1.Draw a Lewis Dot Diagram for an oxygen atom 2.Draw the Lewis Dot Diagram for a chlorine atom

18. Example: HBr # Valence electrons = Write the symbols:

19. Study Question. 2 H 2 O NF 3

20. Study Question. 3 O 2 CO 2 CO N 2

21. Exceptions to the octet rule… P and S can hold more than 8 electrons in some molecules Ex) PBr 5 Boron only requires 6 electrons to be stable Ex) BH 3

22. Lewis Dot Structures for Ions Works the same way, BUT For each + charge subtract an electron. For each – charge add an electron Ex) Cl-  8 valence electrons

23. 16.3 – Bond Polarity Atoms participating in covalent bonds share electrons, but they do not always share equally. Recall: Electronegativity: the tendency for an atom to attract electrons to itself when chemically combined with another element.

24. Nonpolar covalent bonds Electrons are shared equally Examples: H 2, O 2, N 2

25. Both Hs have an electronegativity of 2.1 (table 14.2, p 405). Equal electronegativy = share electrons equally = nonpolar bond

26. Polar Bonds Don’t share electrons equally. Have a difference in electronegativity of.4 – 2 Ex) HCl H: 2.1 Cl: 3.0 Difference of.9 so polar bond.

27. Chemical Bonding Ionic BondsCovalent Bonds Non polar Polar >2 0-0.40.4-2

28. Non Polar Molecules Symmetric shapes with similar atoms. Ex) tetrahedral, linear, trigonal planar… To get this, most times you must draw lewis dot structures.

29. Polar Molecules Molecules that have partially positive end and one partially negative end. Dipole: molecule with two charged regions (or “poles”)

30. Remember…Just because it has polar bonds doesn’t make it polar… Polar bonds can “cancel” each other out and result in a non polar molecule.

31. study question 4 Based on the general trend for electronegativity, determine whether the following molecules are polar or non-polar: CO HCl O 2 NaCl

32. Intermolecular Forces (IMF) Attractions BETWEEN MOLECULES van der Waals Forces Dispersion Forces Dipole Interactions Weakest type of intermolecular interactions

33. Dispersion Forces Weakest of all IMFs Caused by motion of electrons Strength of Dispersion forces increases with increasing size (molecular weight).

34. Study question 5 Which atoms would probably create stronger London forces? Small atoms or large atoms. Why?

35. Dipole Interactions Occurs when polar molecules are attracted to each other

36. Hydrogen Bonding Very strong dipole interaction between H and F, N, or O. A hydrogen bond is 5% the strength of the average covalent bond.

37. study question 5 Which of the following molecules might hydrogen bond? CH 4, CH 2 O, HF, H 2 O, H 2 S

38. Ionic vs. Covalent ** See table 16.5 ** Melting Point: ionic > covalent Electrical Conductivity Ionic > covalent Makeup of elements: Solubility: Covalent depends on make up of molecule