1. Nomenclature

2. Valence electrons  Valence electrons are the electrons that are in the highest energy level of an atom.  These electrons are involved in forming bonds with other atoms.  Elements (except helium) have the same # of valence electrons as their group #.  Electron dot structures are used to show valence electrons.

4. Octet rule  In order to become stable, atoms tend to either gain or lose valence electrons so that its highest energy level will become full with 8 electrons, similar to a noble gas. (except He, which has 2).  This is called the octet rule.

5. Ionic compounds  A cation (+) is formed when an atom loses electrons. Usually metals are cations.  An anion (-) is formed when an atom gains electrons. Usually nonmetals are anions  Cations and anions have opposite charges and are attracted to one another.  These attractive forces can hold the ions together in an ionic bond, forming a compound.  Ionic compounds are usually made up of a metal and nonmetal.

7. Properties of ionic compounds  High melting point  Low malleability – break and shatter easily  Can conduct electricity under certain conditions

8. Naming Ionic Compounds 1. Name the cation 1 st & the anion 2 nd. 2. Monatomic (1 element) cations use the element name without any change. 3. Monatomic anions use the root of their element name plus the suffix -ide 4. If the compound has a polyatomic ion, simply name that ion without any change. 5. If the ion has more than one “common” form, it has to be labeled with a Roman numeral. Ex. Iron (II) phosphate

10. Some metal ions with more than one common charge  Fe: 2+ and 3+  Cu: 1+ and 2+  Sn: 2+ and 4+  Pb: 2+ and 4+  Co: 2+ and 3+  Mn: 2+, 6+ and 7+

11. Polyatomic ions  NH 4 + Ammonium  OH - Hydroxide  NO 3 - Nitrate  CO 3 2- Carbonate  SO 4 2- Sulfate  PO 4 3- Phosphate  HCO 3 - Hydrogen carbonate / bicarbonate  C 2 H 3 O 2 - Acetate  SO 3 2- Sulfite

12. Hydrates  Hydrates are crystalline compounds which attract and hold water molecules.  The water is called the water of hydration and can be removed (evaporated) by heating.  After water is removed the crystal is said to be anhydrous.

13. Naming hydrates  To name hydrates simply name the compound (usually ionic) and then indicate the number of water molecules by using the same prefixes as in molecular compounds.  CuSO 4 5 H 2 O  Copper (II) sulfate pentahydrate  Sodium carbonate heptahydrate  Na 2 CO 3 7 H 2 O

14. Bonding in metals  Valence electrons of metal atoms can be modeled as a sea of electrons – they are mobile and can drift from one part of the metal to the other  Metallic Bond – the attraction of these “free-floating” electrons for the metal ions Metallic Bond  These bonds hold metals together and explain many of their physical properties

15. Molecular (covalent bonding) compounds  A covalent bond results from the sharing of electrons. The octet rule still applies  Covalent bonds generally occur when elements are close to each other on the periodic table.  The majority of covalent bonds form between nonmetallic elements.

16. Example of a covalent bond

17. Diatomic Elements (elements that exist in pairs)  HydrogenH 2  OxygenO 2  NitrogenN 2  FluorineF 2  ChlorineCl 2  BromineBr 2  IodineI 2

18. Properties of molecular compounds  Usually have a lower MP (than ionic, metallic)  Generally soft  Non-conductors (in any state & aqueous)  Molecular Compounds can exist in all 3 states:  Solids – sugar, ice, aspirin  Liquids – water, alcohols  Gases – O 2, CO 2, and N 2 O (laughing gas)  Very important to organic chemistry and pharmaceutics.

19. Naming Molecular (covalent bonding) Compounds 1. The 1 st element is named first, using the entire element name 2. The 2 nd element is named using the root of the element & adding the suffix –ide 3. Prefixes are used to indicate the # of atoms of each type that are present. *Exception: The 1 st element in a formula never uses the prefix mono-

20. Prefixes in Molecular (covalent) Compounds # of atoms Prefix# of atoms Prefix 1mono-6hexa- 2di-7hepta- 3tri-8octa- 4tetra-9nona- 5penta-10deca-

21. Ionic vs. Covalent  Look at the 1 st element:  1 st element metal: ionic compound  1 st element non-metal: covalent compound  Exception: ammonium (NH 4 + ): ionic compound

22. Naming acids 1. If the anion part normally ends in – ide (binary acid), then the acid name begins with the prefix hydro and ends with –ic. Ex. HCl is hydrochloric acid 2. If the anion part ends in –ate (polyatomic) then NO hydro is used and the ending is –ic. Ex. HNO 3 is nitric acid (notice – no hydro).

23. 3. If the anion part normally ends in –ite no hydro is used and the ending is –ous. Ex. HNO 2 is nitrous acid

24. 4. To write formulas for acids just use the number of H’s equal to the negative charge of the anion (since each H is +1). Ex. Carbonic acid – no hydro is used so the anion must be polyatomic. The acid name ends in –ic so the anion must end in –ate, i.e. carbonate. Since carbonate is CO 3 2- two H’s are necessary and the formula is H 2 CO 3.

25. NOTE: For a couple of elements only part, or even none, of the element ending is dropped before adding the acid ending. Ex. H 2 SO 4 is not sulfic acid it is sulfuric acid. H 3 PO 4 is phosphoric acid, no phosphic acid.