1. Covalent Compounds & Molecule Shapes 2014

2. I. IONIC COMPOUNDS >How are Ionic Compounds Made? >Made of metal and nonmetal (or sometimes, polyatomic ions may be present) >Electrons are transferred >Individual particles are called formula units Examples: (salts) aluminum chloride, copper II nitrate

3. >Writing Formulas and Names of Ionic Compounds Which elements? > Metal names don’t change while nonmetals end in “-ide”. >Polyatomic ion names don’t change either. >Many transition metals require a roman numeral to indicate their oxidation number >Polyatomic ions are written inside parenthesis >Oxidation numbers are used to determine subscripts.

4. >Writing Formulas and Names of Ionic Compounds >Oxidation numbers are used to determine subscripts. >Examples: Aluminum chloride Tin IV Oxide Al +3 Cl -1 Sn +4 O -2 x1 x3 x2 x4 Al Cl 3 Sn 2 O 4 -- > reduces to SnO 2 really Sodium sulfate Na +1 SO 4 -2 barium chlorate Ba +2 ClO 3 -1 x2 x1 x1 x2 Na 2 (SO 4 ) Ba(ClO 3 ) 2

5. II. COVALENT COMPOUNDS >How are Covalent Compounds Made? >Made of 2 nonmetals or sometimes hydrogen and a nonmetal >Electrons are shared >Individual particles are called molecules >Examples: water, sugar, carbon dioxide, dinitrogen pentoxide

6. > STEPS for Writing Names of Covalent Compounds a. Write the name of the first element (IF there is MORE THAN ONE atom of the first element, use a prefix!) b.Identify the number of atoms of the second element, write the prefix that indicates how many of that element, then…. c. Write the name of the second element and change its ending to “-ide” PREFIXES for step a and b…. 1 = mono- 4 = tetra- 7 = hepta- (some use septa-) 2 = di- 5 = penta- 8 = octa- 3 = tri- 6 = hexa- 9 = nona- 10 – deca- That’s enough!

7. EXAMPLES: SiO 2 = silicon dioxide (one silicon on first name….NO Prefix) P 4 O 10 = tetraphosphorus decoxide (drop the prefix’s “a” when using oxygen) N 2 F 5 = dinitrogen pentafluoride (the first name’s ending doesn’t change)

8. >Writing Formulas for Covalent Compounds Use the prefixes to find how many atoms of the element Do NOT reduce the formula, even if the subscripts “could” be reduced! (hydrogen peroxide for cuts) dihydrogen dioxide = H 2 O 2

9. III. COMPARING PROPERTIES OF IONIC AND COVALENT BONDS / COMPOUNDS Ionic Bonds / Compounds Made of metal & nonmetal generally Electrons transferred (e- lost by one, gained by the other) HIGH boiling point HIGH melting point (can’t melt it in the kitchen; it’s solid) CONDUCT electricity (when melted or as a solid) Crystalline structure Ex. stalactites and stalagmites in caves Often is a SALT Nearly always dissolves in water (Electrolyte) 1 particle is called a Formula Unit

10. III. COMPARING PROPERTIES OF IONIC AND COVALENT BONDS / COMPOUNDS Covalent Bonds / Compounds Made of 2 nonmetals Electrons are shared LOW boiling point (turn to gas easily or IS a gas) LOW melting point Do NOT conduct electricity Many do NOT dissolve in water (nonelectrolyte) 1 particle is called a Molecule

11. . IV. COVALENT BONDS >Why do they form? Nonmetals try to get 8 valence electrons like noble gases have. Exceptions: Hydrogen actually tries to get 2 valence electrons, not 8. Beryllium actually tries to get 4 valence electrons, not 8. Boron actually tries to get 6 valence electrons, not 8.

12. . IV. COVALENT BONDS >Why do they form? Nonmetals try to get 8 valence electrons like noble gases have. Exceptions: Hydrogen actually tries to get 2 total valence electrons, not 8. Beryllium actually tries to get 4 total valence electrons, not 8. Boron actually tries to get 6 total valence electrons, not 8. 2 nonmetals can’t both “gain” electrons, so they “share” electrons instead As a compound, the elements have less energy and, therefore, are more stable.

13. Get ready to learn how to predict the shapes of molecules: Lewis Dot Structures of Elements--- dots represent only the valence electrons Be Li B CNO F Ne 1 2 3636 4 7 5858

14. . Types of Covalent Bonds that can form Polar = unequal sharing of electrons One element has a stronger pull on the electrons ( more pull on e- = more electronegative) Nonpolar = equal sharing of electrons The elements making the bond have equal (or very nearly equal) pull on the electrons

15. Covalent Bonds can be: A single bond is when 2 electrons are shared. A double bond is when 4 electrons are shared A triple bond is when 6 electrons are shared H-H 

16. . Molecular Models Models help us demonstrate the 3-D image of a compound Even though the atom is too small to see, we can predict its molecular geometry (shape). What molecule shapes can happen? What do they look like? ………………………….

17. Linear  2 or 3 bonded atoms  No lone electrons  180 o Bond Angle  Symmetrical shape  Example: BeCl 2

18. Bent 3 bonded atoms 1 or 2 lone pairs of electrons 104.5 o Bond Angle Though not visible in the model, a lone pair of electrons takes up more space than a single atom (Electrons repel each other!) Asymmetrical Example: SO 2 (left picture--not common!), H 2 O (right picture 2 lone pairs = more common)

19. Trigonal Planar  4 bonded atoms  No lone electrons  120 o Bond Angle  Symmetrical shape  Example: BCl 3

20. Trigonal Pyramidal (or Triangular Pyramidal) 4 bonded atoms 1 lone pair of electrons 107.3 o Bond Angle Though not visible in the model, a lone pair of electrons takes up more space than a single atom (Electrons repel each other!) Asymmetrical shape Example: PH 3

21. Tetrahedral  5 bonded atoms  No lone electrons  109.5 o Bond Angle  Symmetrical shape  Example: CH 4

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