1. Chemical Bonding

2. I. Introduction  A. Types of Chemical Bonds – forces that hold two atom together 1. Ionic Bonds – occur b/w a metal & a nonmetal 2. Covalent Bonds – occur b/w 2 nonmetals & in polyatomic ions  a. Polar Covalent Bonds -

3. II. Ionic Bonds  A. A strong bond caused by the transfer of electrons from a cation (metal) to an anion (nonmetal). 1. Why? The driving force behind this bonding is that all elements want to have a completely filled outermost energy level! [OCTET RULE]  a.) These outermost electrons are called the VALENCE ELECTRONS  b.) Metals LOSE valence electrons to be stable.  c.) Nonmetals GAIN valence electrons to be stable.

4. Valence Electrons 1e-8e- 2e-3e-4e-5e-6e-7e- 2e- 1e- 2e-

6. Let’s try it!  1. Na and O  2. Al and F  3. Ca and S  4. Mg and P

7.  B. Ionic Bonding And Structures of Ionic Compounds 1. Ionic compounds are  a. very stable, huge amounts of energy necessary to break them apart  b. high melting & boiling points NaCl has a melting point = ~800°C

8. 2. Structures of Ionic Compounds  a. When you write the formula for an ionic compound, you are writing its empirical formula.  b. In reality, the actual solid contains tremendous amounts & equal numbers of cations and anions packed together so that the attractions b/w them are maximized. 1.) Remember that cations are always smaller than anions. WHY?

10. III. Covalent Bonding  A. Sharing electrons! 1. All bonding involves valence electrons ONLY!!!!!! 2. Covalent bonds occur when 2 atoms (usually nonmetals) share electrons. 3. LEWIS STRUCTURE – a representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule.  Thought up by G.N. Lewis while teaching a chemistry class in 1902.

11. See attached page for writing Lewis Structures!

12.  B. Structures – VSEPR Model 1. Valence Shell Electron Pair Repulsion Model  a. Useful for predicting the geometric shape of molecules formed from nonmetals!  b. The structure around a given atom is determined by minimizing repulsions between electron pairs.

13.  c. Steps for predicting VSEPR models 1. Draw the Lewis structure for the molecule. 2. Count the electron pairs & arrange them in the way that minimizes repulsions (PUT THEM AS FAR APART AS POSSIBLE!) 3. Determine the positions of the atoms from the way the electron pairs are shared. 4. Determine the name of the structure from the positions!

14. Metallic Bonding  How atoms are held together in the solid.  Metals hold onto their valence electrons very weakly.  Think of them as positive ions floating in a sea of electrons!

15. Sea of Electrons!  Electrons are free to move through the solid.  Metals conduct electricity. ++++ ++++ ++++

16. Metals are malleable!  Hammered into shape (bend).  Ductile - drawn into wires. ++++ ++++ ++++

17. Malleable  Electrons allow atoms to slide by. ++++ +++ ++++

18. Alloys  Solutions made by dissolving metal into other elements- usually metals.  Melt them together and cool them.  If the atoms of the metals are about the same size, they substitute for each other  Called a substitutional alloy

19. Metal A Metal B +  Substitutional alloy Bronze – Copper and Tin Brass- 60 % Copper 39% Zinc and 1%Tin 18 carat gold- 75% gold, 25%Ag or Cu

20. Alloys  If they are different sizes the small one will fit into the spaces of the larger one  Called and interstitial alloy

21. Metal A +  Metal B Interstitial Alloy Steel – 99% iron 1 % C Cast iron- 96% Iron, 4%C

22. Alloys  Making an alloy is still just a mixture  Blend the properties  Still held together with metallic bonding  Most of the metals we use daily are alloys.  Designed for a purpose