1. Chapter 8: Covalent Bonding Because no one wants to be alone.

2. Objectives 1.Compare and contrast covalent and ionic bonding. 2.Use electronegativity difference to determine the type of a chemical bond that will form between 2 elements. 3.Draw Lewis structures for covalent compounds. 4.Use the VSEPR theory to describe molecule shape, bond polarity, and molecule polarity. 5.Describe the 3 types of intermolecular forces.

3. Covalent Chemical Bond: a strong attractive force between atoms that form molecules. Molecule – a neutral group of atoms joined together by covalent bonds.

4. Covalent Bonds Chemical bond between 2 or more atoms sharing electrons 2 atoms = diatomic molecule

5. Covalent Bonds Nonpolar covalent bonds –Electrons are shared equally –Small difference in electronegativity Polar covalent bonds –Electrons are shared unequally –Large difference in electronegativity

6. LEARNING CHECK: 1.What are the 2 main differences between ionic and covalent bonds? 2.What are the 2 types of covalent bonds? 3.Which type of bond exhibits higher electronegativity difference?

7. Ionic Vs. Covalent Bonds

8. Molecular Compounds and the Octet Rule –Covalent bonds occur by atoms sharing electrons with the goal of getting 8 electrons in their outer energy level. –This forms molecular compounds a.k.a. MOLECULES –Molecular formulas are chemical formulas for covalently-bonded atoms. – END OF PART 1

9. Lewis Structures & Structural Formulas –The Lewis structure shows both how electrons are shared within the molecule as well as the unshared valence electrons –A structural formula shows us how the atoms are shared in a molecule but does not show the unshared pairs of electrons Example: F–F and H–Cl –Single bonds are one shared pair, a double bond two shared pairs, and a triple bond is three shared pairs of electrons

10. Drawing Lewis Structures Let’s use the compound C 2 H 2 The NASL Method: 1.Write the electron-dot notation for each type of atom. 2.Chose a central atom. Your choice will be the atom with the greatest number of unpaired electrons. If carbon is present it is in the center. Hydrogen is never in the center.

11. Drawing Lewis Structures 3. Calculate N (Needed) as the sum of electrons needed for all atoms by the octet rule. Exceptions: H=2, Be=4, B=6. 4.Calculate A (Available) as the sum of all valence electrons. 5.Calculate S (Shared) as the difference between N – A. 6.Divide S by 2 to obtain the number of bonds to be extended from the central atom. 7.Calculate L (Lone-pair electrons or simply “dots”) as the difference between A – S. Place the L dots into the skeleton as to fill the octet of every atom except hydrogen. Remember that hydrogen has only one bond and NO dots. 8.Check that the total number of used electrons is equal to A.

12. Drawing Lewis Structures Example: H 2 is written as H:H or H–H –Try H 2 O, O 2, and F 2

13. Resonance Structures Anytime a molecule or ion cannot be correctly represented by a single Lewis structure, it has resonance structures –E–Examples: O 3 SO 4 - NO 3 - –D–Draw a Lewis structure for O 3 –N–Now draw its resonance structure. –G–Got it???

14. LEARNING CHECK: 1.What is a molecule? 2.How is the octet rule satisfied in molecular compounds (molecules)? END OF PART 2

15. VSEPR Valence Shell Electron Pair Repulsion –Repulsions between electron pairs makes them orient as far as possible from one another Count the electron pairs on the central atom –If 2 then linearBeCl 2 –If 3 then trigonal planarBF 3 –If 4 then tetrahedralCH 4 –If 5 (rare) then trigonal bipyramidalPCl 5 –If 6 (rare) then octahedralSF 6

16. LEARNING CHECK: What does the VSEPR theory mean?

18. VSEPR Unshared electron pairs on the central do play a role in the 3-D geometry of the molecule –W–When we describe the shape of a molecule we refer only to the relative positions of the atoms –E–Examples: NH 3 and H 2 O Trigonal pyramidalBent or Angular

19. END OF PART 3

20. Intermolecular Forces Forces of attraction between molecules –Typically much weaker than atomic bonds –Three types: Dipole-Dipole Forces: occur between polar molecules (direction from + to – ) Hydrogen bonds: is like a dipole-dipole, only hydrogen is bonded to a highly electronegative atom (F, O, Cl) London Dispersion Forces: movement of electrons which cause momentary dipole like attractions

21. Molecular Dipole Moments Dipole moments have positive and negative ends (arrow is negative, tail is positive)Dipole moments have positive and negative ends (arrow is negative, tail is positive) The arrow head is attracted to the tail of another arrowThe arrow head is attracted to the tail of another arrow

22. Hydrogen Bonding Hydrogen bonds are very similar to dipole-dipole interactions but always involve the positive hydrogen side of a molecule and the negative side of another molecule

23. LEARNING CHECK 1.How do multiple molecules stay together? 2.Describe how water (H 2 O) molecules stay together.