1. Unit 9 Chemical Bonding

2. Electronegativity

3. Electronegativity is the relative tendency of an atom to attract electrons when it is bonded to another atom Trend for electronegativity –Increases as you move up a group or to the right across a period

4. Periodic Trend for Electronegativity

5. Assignment 1-2 page 231 5 points Due at the end of class

6. Bond Character The electronegativity difference between two bonded atoms can be used to determine what type of bond is present –If there is little difference it is a covalent bond –If there is a difference in electronegativities, you have a ionic bond

7. Bond character Assignment –3-4 page 232 (use chart on page 230) –Due at end of class (5 points)

8. Ionic Bonds Ionic bonds form because of an electrostatic attraction between ions Best explained as a complete transfer of electrons An ionic bond is the electrostatic force that holds together two ions due to their differing charges

9. Ionic Radii See table 10-1 in text (page 186) Internuclear distance – the distance between the central point of two ions (radius + radius)

10. Covalent Bonds A shared pair of electrons makes a covalent bond Resultant particle is called a molecule The line joining the nuclei is the bond axis

11. Covalent Bonds The angle between the two bond axes the bond angle The distance between the nuclei along the bond axis is the bond length

12. Characteristics of Compounds with Covalent Bonds Low melting points Brittle Do not conduct electricity

13. Examples of covalent bonds

14. Covalent Radii Approximate bond length can be determined by adding the covalent radii of 2 bonded atoms The covalent radii are averages obtained from a number of different compounds Complete exercise 5 on page 236

15. Polyatomic Ions An ionic compound made of two or more elements Made up of covalently bonded atoms See table 4-4

16. Van Der Waals Radii Van Der Waals Radius is the minimum distance maintained between nonbonded atoms because of electron cloud repulsion See table 12-6 pg 236

17. Special Properties of Metals Metals form a crystalline structure where outer level orbitals of the atoms overlap, leading to the existence of delocalized electrons –The delocalized electrons holding the metallic atoms together make a metallic bond The metallic bond is responsible for the strength of metals The more outer electrons, the stronger the metal –The delocalized electrons are the reason for metals high conductivity

18. Linus Pauling Says:

19. Chapter Review Complete Questions