1. Chapters 8 and 9 Ionic and Covalent Bonding

2. Forming Chemical Bonds Chemical Bond  Force that holds 2 atoms together  Attraction between + nucleus and e- or between + and – ions  Only valence e- involved in bonding  Bonding occurs to have complete outermost energy levels – to become like noble gases

3. Forming Ions Ion  An atom or bonded group of atoms that have a + or – charge because valence e- have been lost or gained Cation  + charged ion; one that has lost e- Anion  - charged ion; one that has gained e-

4. Formation and Nature of Ionic Bonds Ionic Bond  Electrostatic force holds oppositely charged particles together  Usually occurs between a metal and nonmetal  #e- lost by one particle = #e- gained by another  Overall charge on a compound is always 0

5. Subscript  Shows how many atoms of an element are in a compound  Applies to the element to its immediate left  If there is no subscript, it means there is only one atom of that element.  What is the number of atoms in each of these compounds: H 2 O CO 2 H 2 SO 4 CO

6. Example: Na +1 and F -1 combine to form NaF (1:1 ratio) Na will lose 1e- and F will gain 1e- Overall charge on NaF is (+1 + -1 = 0) Swap Charges: Na +1 F -1  NaF

7. Types Oxide  Ionic compound with metal and oxygen Salt  Most other ionic compounds, usually between a metal and a halogen Binary Compounds  Contains 1 metallic and 1 nonmetallic ion Ternary Compounds  Contains 1 polyatomic and 1 monatomic ion

8. Important Vocab Formula Unit  Simplest ratio of ions in an ionic compound  The formula of the compound  Ex: NaCl; KBr Oxidation Number  Charge on an ion  #e- atom must gain/lose/share to complete its outermost energy level  Metal = #valence e-  Non-metal = #valence e- - 8

9. Properties of Ionic Compounds Form crystals: + and – ions are packed into a regular repeating pattern. Crystal shape depends on size and # of ions bonded IB are strong because particles are strongly attracted to each other  Bonds require much E to break apart  Compounds have high melting and boiling points  Mostly solids at room temp.  Hard, rigid, and brittle

10. When in aqueous solution (dissolved in water) ions move about and conduct electricity (electrolyte) Forming compounds from ions is exothermic  Gives off Energy Breaking down compounds into ions is endothermic  Requires or takes in Energy Properties of Ionic Compounds

11. Writing Formula Units 1. Identify the cation and write it’s symbol. 1. It will be the first element written in it’s entirety. 2. Identify the anion and write it’s symbol. 1. It will be the second element written as the root of the name + “ide” 3. Find the oxidation number for each element and write it above each element without the sign. 4. Swap the oxidation numbers and write the formula.

12. Examples: Calcium Chloride Ca +2 Cl -1 -> Ca 2 Cl 1 Ca 1 Cl 2 ----  CaCl 2 Sodium Oxide Na +1 O -2 -> Na 1 O 2 Na 2 O 1 -> Na 2 O

13. Covalent Bonds Chemical bond resulting from sharing valence e-  Usually between elements close on periodic table (usually nonmetals)  Not as strong as ionic bonds  Forms molecules or diatomic molecules  Atoms are too far away from each other to have a strong attraction. 1 atoms + nucleus attracts another’s e- cloud. But both clouds repel each other. Distance is right then to share e-, and not transfer them.

14. Molecules and Diatomic Molecules Molecule  2 different elements bonded covalently Diatomic Molecules  2 of the same element bonded covalently  H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2

15. Representing Molecules Molecular Formula  H 2 Electron Dot Structure H Lewis Structure  H – H Each dot represents an e-. A line represents a bonding pair of shared e-

16. Examples of Single Bonds H + OH  H – O – H 2H 2 + O 2  H 2 O H H + H + NH  H – N – H H 2  H – H 22 22 22

17. Properties of Covalent Compounds Melting and boiling points are lower than those of ionic compounds Many are gases at room temperature If in the solid form, molecules are soft or brittle

18. Examples: O + O  O = O N + N  N = N