1. Bonding

2. Think of some properties of salt Forms crystals Brittle Hard Solid High melting and boiling point Forms an electrolyte (conducts electricity in water)

3. Ionic Compounds NaCl or table salt is an ionic compound. Ionic compounds form from ionic bonds – bonds formed by electrostatic attraction of positive and negative ions Positive ions are called cations Negative ions are called anions

4. Ionic Compounds Form between metals and nonmetals Or between metals and polyatomic ions. A polyatomic ion is a covalently bonded compound that has lost or gained electron(s). The polyatomic ion acts as a unit – one thing.

5. Determining ionic compound formulas 1.Write the element symbols involved in bonding K 1+ 2- O K 2 O 1.Next, write the oxidation numbers above the symbols 2.Now, move the numbers only across to the other symbols lower right corner – they become the subscripts

6. Naming ionic compounds A binary ionic compound is composed of 2 different elements The cation is always written first and is just the elements name The anion is written next by adding the suffix “ide” to the root of the element name Example: sodium and chlorine becomes sodium chloride

7. Covalent compounds Formed when atoms share valence electrons Sharing between nonmetals Sharing due to similar electronegativites, so both atoms are attracting electrons

8. Properties of covalent compounds Generally lower melting and boiling points Soft compared to ionic compounds Tend to be more flammable Aren’t usually very soluble in water

9. Naming Covalent Compounds The first word in the name is the first elements name The second word in the name is the root of the second element with the suffix “ide” Prefixes are added to both the first and second element to indicate how many of each element are in the formula 1- mono; 2 – di; 3 – tri; 4 – tetra; 5 – penta; 6 – hexa; 7 – hepta; 8 – octa; 9 – nona; 10 – deca Exception if thereis only one of the first element do not add mono

10. Diatomic molecules Hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine and iodine Always exist as 2 atoms bonded together H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2

11. CO 2 – carbon dioxide P 2 O 5 – diphosphorus pentoxide

12. Metallic bonds Bonds formed between metals when the positive metal ions are attracted to free valence electrons These valence electrons are delocalized – Delocalized electrons are free to move throughout the metal and are not attached to any one atom

13. Metallic bond properties Varying melting points Malleable – can be hammered into sheets Ductile – can be drawn into a wire Lustrous - shiny Conductive – will conduct heat and/or electricity

14. Forces within and between particles Intramolecular forces – forces that stick atoms together in a compound – a chemical bond Intermolecular forces – forces that hold one covalent molecule to another

15. Intramolecular forces Ionic bonds – transfer valence electrons; metal and nonmetal or metal and polyatomic ion Metallic – sea of electrons; each ion is bonded to all others by a “sea” of valence electrons; attraction of the metallic cation for all delocalized electrons; metal and metal Covalent – sharing of valence electrons; 2 or more nonmetals

16. Covalent bonds Covalent bonds may be polar or nonpolar Nonpolar – equal sharing of valence electrons Polar – unequal sharing of valence electrons; one atom has a stronger attraction for electrons (electronegativity) than the other – Results in a molecule with a slightly positive end and a slightly negative end

17. Polar bond - H 2 O Nonpolar bond - Cl 2

18. Geometry and polarity A molecule can contain polar bonds and still be a nonpolar molecule and is related to the molecules geometry – Symmetrical molecules are generally nonpolar and asymmetrical molecules are generally polar

19. Intermolecular forces Forces that hold one covalent molecule to another Dipole-dipole force – the attraction between the oppositely charged ends of a polar covalent bond

20. Hydrogen bonds a dipole-dipole bond between molecules containing a H atom bonded to a small highly electronegative atom with at least one lone pair. example: water

21. Dispersion forces weak forces from temporary shifts in the density of electrons in the electron cloud also called London forces

22. Van der Waals forces Force of attractions between nonpolar molecules