1. Ionic Compounds Ions are atoms that have gained or lost electron(s) Atoms tend to make ions with characteristic oxidation states (charges) Metals are losers (+ ions) Nonmetals are gainers (- ions) Some atoms do not readily make ions (C, Si, many metalloids)

2. Ionic Bonds An ionic bond is formed between two or more oppositely charged ions Ionic compounds are made of a metal (+) and a nonmetal (-) Ionic compounds are called salts The overall charge on an ionic compound is zero When a metal and nonmetal react, electrons are transferred

3. Making Ionic Bonds NaCl 11p + 11e - 17p + 17e - 10e - + 18e - -

4. Making More Ionic Bonds MgO 12 p + 12e - 8p + 8e - 10e - +2 10e - -2

5. Formulas and Names NaCl, sodium chloride MgO, magnesium oxide Binary salts –metal, then nonmetal –nonmetal ending changed to “ide” –no subscripts when ratio is 1:1

6. Unequal charges Na S 11 p + 11e - 16p + 16e - 10e - +1 Na 11 p + 11e - 10e - +1 18e - -2 Formula: Na 2 S Name: sodium sulfide Total + charge = 2, Total - charge = 2 Total charge overall = 1 + 1 + (-2) = 0

7. Unequal charges Determine the formula of calcium bromide. Ca Br Formula: CaBr 2 +2 -1 -1 +1

8. Polyatomic ions Transition metal salts Salts of polyatomic ions

9. Solubilities Salts are soluble in water if ion-water interactions can supply enough energy to break apart the crystal lattice Salts of lower-charged ions are more likely to be soluble (lower lattice energy) All alkali metal and ammonium salts are soluble All nitrates are soluble All oxides are insoluble (alkali metal oxides react to form hydroxides)

10. Ionic compound properties Made of metal and nonmetal (except ammonium and organic base salts) High MP (chemical bonds are broken in melting) Crystal lattice Brittle Form ions in water solution (ionization) NaCl  Na + + Cl - Conduct electricity when melted

11. Hydrates Water can get trapped in crystal lattice of a crystallized salt Na 2 CO 3. 10H 2 O Sodium carbonate decahydrate CuSO 4. 5H 2 O Copper (II) sulfate pentahydrate Sodium acetate trihydrate NaC 2 H 3 O 2. 3H 2 O

12. Hydrates Some salts take water out of the air to become hydrates: hygroscopic Example: Na 2 CO 3 Others take enough water to become solutions: deliquescent Example: CaCl 2

13. Crystal Lattices and Energy Regular repeating arrangement of ions is a crystal lattice Energy holding lattice together is the lattice energy Energy is released when lattice is formed (from gaseous ions) and absorbed when it is broken

14. Crystal Lattices and Energy Lattice energy is measured from the viewpoint of the system When gaseous ions come together to form a crystal energy leaves the system Since system energy is lower, lattice energy is always given as a negative value

15. Crystal Lattices and Energy Magnitude of lattice energy is directly proportional to charge density Charge density is related to charge magnitude and ion size Crystallization from gaseous ions is always negative; crystallization from solution can be negative or positive

16. Metallic Bonds Metals form molecular orbitals that cover the entire crystal Electrons can move anywhere in the orbital, so metals conduct heat and electricity well Metallic bonds are non-directional, so metals are malleable and ductile Strength of metallic bonds depends on the number of mobile electrons in the bond per atom Transition metals have mobile s and d electrons, so they are stronger and harder than alkali metals (only 1 s electron is mobile)

17. Metal Alloys Alloys are solid solutions of one or more metals Substitutional alloy: made by metals with atoms of similar size Interstitial alloy: made by metals with very different atomic sizes Adding nonmetals (such as carbon to iron) makes directional bonds Directional bonds make alloys harder, stronger and more brittle

18. Covalent Bonds Nonmetals of similar electronegativity cannot form ionic bonds These atoms share electrons to complete their octet Shared electrons “count” for both atoms Each atom’s nucleus attracts the other atom’s electrons

19. Forming Covalent Bonds H Cl Shared! Single bond, 2 electrons 8 e - ! 2 e - !

20. Multiple Bonds CO H H Needs 1e -, makes 1 bond Needs 4 e -, makes 4 bonds Needs 2 e -, Makes 2 bonds 2e - ! 8e - ! Sigma  bond Electron density between nuclei Pi (  ) bond electron density above and below nuclei Double bond

21. Molecular Dot Structures Count electrons – all valence electrons must appear in final structure Follow octet rule Remember how many bonds each type of atom makes (one for each extra electron needed)

22. Polyatomic Ion Dot Structures Same as molecular dot structures, except electrons must be added or subtracted to account for ion charge Subtract electrons for + charge, add for – charge Make all structures as symmetrical as possible

23. Carbonate (CO 3 -2 ) Dot Structure Symmetry! C O OO Count electrons! 6 + 6 + 6 + 4 + 2 = 24 -2

24. Molecular Substances Made of molecules, which are loosely held together – van der Waals or London Dispersion forces Tend to be liquids, gases or low melting solids Melting molecular solids involves separating molecules from each other Most are insulators

25. Formulas and Names of Small Molecules Many have common names (i.e. water, ammonia) Systematic names use prefixes for each element P 2 O 5 – diphosphorus pentoxide N 2 O – dinitrogen monoxide “mono” is not used for the first element in a compound

26. Formulas and Names of Small Molecules CO 2 – carbon dioxide CO – carbon monoxide SO 3 – sulfur trioxide CCl 4 – carbon tetrachloride