1. Chapter 12 Chemical Bonding

2. Chemical Bond A bond is an electrostatic force of attraction holding two atoms together. Electrostatic forces can be either attractive or repulsive (opposite charges attract whereas like charges repel). “Perhaps you gentlemen can tell me what it is that you find so attractive”.

3. Chemical Bond Bonds are formed by valence electrons. The type of bond that forms between atoms is based on what is occurring with the valence electrons.

4. Ionic Bond Ionic bonds are formed when electrons are transferred (lost and gained). Ionic bonds tend to occur when a metal bonds with a nonmetal.

5. Covalent Bond Covalent bonds are formed when electrons are shared. Covalent bonds tend to occur when a nonmetal bonds with another nonmetal.

6. Polar Covalent and Nonpolar Covalent Covalent bonds can be nonpolar covalent which result when electrons are shared equally between atoms or polar covalent when electrons are shared unequally between atoms.

7. Polar Covalent Bond A a

8. Metallic Bond Metallic bonds occur when valence electrons drift from one metal atom to another. These “delocalized” electrons therefore do not belong to any particular metal atom but instead can be thought of as belonging to all the metal atoms.

9. Delocalized Electrons

10. The delocalized electrons are what give metals some of their unique properties such as malleability, conductivity and luster. The more delocalized electrons a metal has the more enhanced these metallic properties are.

11. Malleable vs. Brittle Does the environment of the substance change? Metallic: Metal ions in a sea of delocalized electrons Ionic: Ions surrounded by oppositely charged ions.

12. Metals are malleable Ionic compounds are brittle

13. Conductivity of Metals

14. 4 Types of Bonds

15. Electronegativity Electronegativity is the tendency of an atom to attract electrons to itself when it is bonded to another atom. It takes into account some of the concepts that we discussed in earlier sections such as atomic size, ionization energy, electron arrangement, etc. It can also be used to determine bond type.

16. Electronegativity Table

17. Bond Polarity Most bonds are a blend of ionic and covalent characteristics. Difference in electronegativity determines bond type.

18. Electronegativity Difference Examples: HCl 2.83 – 2.20 = 0.63 ( polar covalent) MgO 3.50 – 1.23 = 2.27 (ionic) H 2 = 2.20 –2.20 = (nonpolar covalent) Electronegativity Difference Electrons are primarily: Type of Bond Less than or equal to 0.3 shared equallynonpolar covalent Greater than 0.3 but less than or equal to 1.7 shared unequallypolar covalent Greater than 1.7transferredionic

19. Electronegativity Difference The difference in electronegativity can also be used to predict bond strength. The greater the difference in electronegativity the stronger the bond. Examples: HCl 2.83 – 2.20 = 0.63 MgO 3.50 – 1.23 = 2.27 H 2 = 2.20 –2.20 = 0 Mg – O is the strongest bond and H – H is the weakest

20. Properties of Polar Covalent bonds Form molecules Unequal sharing of electrons One atom has a stronger pull on the electrons than the other atom does. Intermediate strength between ionic are nonpolar covalent bonds. 88% of all bonds are polar covalent (Ionic bonds make up about 10% of all bonds)

21. ++ -- Bond Polarity Polar Covalent Bond –e - are shared unequally –nonsymmetrical e - density –results in partial charges (dipole)

22. Properties of Non-Polar Bonds Equal sharing of electrons Weakest bond type About 2% of all bonds are nonpolar covalent. Usually identical atoms or nonmetals very close on periodic table.

23. zIonic zPolar zNonpolar Bond Polarity

24. Chemistry Video Project

25. Homework Chapter 12 Homework Worksheet (9 questions).