1. Chemical Bonding

2. Naming Ions Recall: oxidation numbers tell us which ion is formed during chemical reactions Naming ions (name of metal) ion (name of nonmetal) –ide ion Examples Na 1+ sodium ion O 2- oxide ion

3. Naming Ions Name the following ions: Mg 2+ F 1- Cl 1- magnesium ion fluoride ion chloride ion

4. Ionic Bonding occurs between a metal and a nonmetal involves a transfer of electrons ions are separate, but are attracted to one another because of opposite charges ions form a compound

5. Ionic Bonding

6. Covalent Bonding Covalent bonding occurs between two nonmetals involves a sharing of electrons combine to form a molecule

7. Physical Properties of Ionic and Covalent Compounds Property:Ionic Compound Covalent Compound Phase at room temperature Usually solidUsually gas DensityMore denseLess dense Melting pointHigher Temps.Lower Temps. Boiling pointHigher Temps.Lower Temps. ConductivityConductsNonconductive

8. Metallic Bonding Occurs between atoms of a metal Valence electrons are shared between the cations

9. Ionic Bonding Electron dot diagrams Sodium & Phosphorus

10. Writing Chemical Formulas Use the criss-cross method: Example: sodium chlorine potassiumoxygen

11. Writing Chemical Formulas Write the chemical formula: magnesiumchlorine Mg 2+ Cl 1- MgCl 2

12. Writing Chemical Formulas Write the chemical formula: lithiumsulfur Li 1+ S 2- Li 2 S

13. Writing Chemical Formulas Make sure you simplify your formula so that it has the smallest subscripts possible. Example: Calcium reacts with Sulfur Ca 2+ S 2- Ca 2 S 2 Simplify to CaS

14. Writing Chemical Formulas Example: BoronNitrogen

15. Writing Chemical Formulas Transition metals many have several oxidation numbers Example: Copper commonly forms ions with a +1 or +2 charge Copper (I), Copper (II) distinguishes between oxidation numbers

16. Writing Chemical Formulas Write the chemical formula: iron (III)oxygen Fe 3+ O 2- Fe 2 O 3

17. Writing Chemical Formulas Write the chemical formula: copper (II)chlorine Cu 2+ Cl 1- CuCl 2

18. Diatomic Molecules Naturally occurring diatomic molecules: Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2 - these occur as gases at room temperature * except Br 2 (liquid at room temp) * except I 2 (solid at room temp) - naming: ex: O 2 oxygen gas N 2 nitrogen gas Br 2 liquid bromine I 2 solid iodine

19. Lewis Dot Diagrams Steps: 1. Find each element’s symbol 2. Find the number of valence electrons for each element (recall: count across the periodic table) 3. Draw a dot beside the symbol for each element to represent valence electrons as you position atoms so that they all fulfill the octet rule. **Helpful tip: when you have more than two atoms bonded, put the element with the least number of atoms in the center**

20. Lewis Dot Diagrams: Covalent Bonding Examples: HF Br 2 CCl 4 H 2 O

21. Lewis Dot Diagrams: Covalent Bonding Some molecules can also form double bonds or triple bonds Example of a double bond: CO 2 Example of a triple bond: N 2

22. Exceptions to the octet rule: Hydrogen (H)2 Beryllium (Be)4 Boron (B)6 Gallium (Ga)6

23. Lewis Dot Diagrams: Ionic Bonding Examples: NaCl MgO

24. Naming Compounds (Ionic Bonding Occurring) ** We use the IUPAC naming system** (International Union of Pure and Applied Chemistry) Steps 1. Write the name of the cation (positive ion) 2. If there is a transition metal, use a roman numeral to indicate the charge. 3. Write the name of the anion (negative ion), but change the ending to –ide Example NaClsodium chloride MgBr 2 magnesium bromide Fe 2 O 3 iron (III) oxide

25. Naming Compounds (Ionic Bonding Occurring) Name the following: MgCl 2 NiBr Li 2 O * What about molecules? (covalent bonding) CO 2

26. Prefixes used in Molecules (When covalent bonding occurs) mono-1 di-2 tri-3 tetra-4 penta-5 hexa-6 hepta-7 octa-8 nona-9 deca-10

27. Naming Molecules (Covalent bonding occurring) Same as naming ionic compounds, except we must use prefixes to indicate the number of atoms of each element. Note: metalloids behave as nonmetals when bonding. Example: N 2 O 5 N 3 O 7 PCl 5 NO BF 3 P 2 O trinitrogen heptoxide phosphorus pentachloride nitrogen monoxide boron trifluoride diphosphorus monoxide dinitrogen pentoxide

28. Common Substances to Recognize H 2 Owater CH 4 methane NH 3 ammonia

29. Summary: Ionic Bonding - Metal and nonmetal - Naming: the metal, then the nonmetal (ending changed to –ide) - Electrons transferred - Smallest unit: ions Covalent Bonding -Two nonmetals -Naming: the first nonmetal, then the second nonmetal (ending changed to – ide). Use prefixes. -Electrons shared -Smallest unit: molecule

30. Polyatomic Ions Polyatomic ions many atoms bonded together with a single charge on the entire molecule When bonding with another ion, polyatomic ions form ionic bonds

31. Common Polyatomic Ions (Chem I) ammoniumNH 4 + hydroxideOH - acetateC 2 H 3 O 2 - nitrateNO 3 - chlorateClO 3 - carbonateCO 3 2- sulfateSO 4 2- phosphatePO 4 3-

32. Common Polyatomic Ions (Pre AP) IonName NH 4 + ammonium NO 2 - nitrite NO 3 - nitrate SO 3 2- sulfite SO 4 2- sulfate HSO 4 - hydrogen sulfate (bisulfate) OH - hydroxide PO 4 3- phosphate IonName NCS - thiocyanate CO 3 2- carbonate HCO 3 - Hydrogen carbonate (bicarbonate) ClO 3 - chlorate C2H3O2-C2H3O2- acetate MnO 4 - permanganate CrO 4 2- chromate O 2 2- peroxide

33. Polyatomic Ions: Writing Formulas Before using the criss-cross method, place parentheses around the polyatomic ion Example calcium and nitrate Ca 2+ NO 3 - Ca(NO 3 ) 2 calcium nitrate

34. Polyatomic Ions in Ionic Bonding Examples: berylliumhydroxide sodiumacetate magnesiumphosphate bariumsulfate ammoniumoxygen silver (II)nitrate

35. Naming Compounds with Polyatomic Ions Examples: Ca(OH) 2 (NH 4 ) 2 O Li 3 PO 4 calcium hydroxide ammonium oxide lithium phosphate

36. Types of Formulas Molecular formula shows the true number of atoms of each element in a compound Empirical formula shows the lowest (smallest) ratio of atoms in a compound Structural formula shows the physical arrangement of atoms within a compound

37. Types of Formulas Example:Benzene Molecular formula C 6 H 6 Empirical formula CH Structural Formula

38. Types of Formulas Example: What type of formula is this? Give the molecular formula and the empirical formula for this molecule.

39. Types of Formulas Note: If the molecular formula cannot be reduced, then the empirical formula is the same as the molecular formula. Example: If the molecular formula is C 3 H 7, what is the empirical formula?

40. Resources http://en.wikipedia.org/wiki/Transition_metal http://fr.wikipedia.org/wiki/Hydrocarbure_aromatique http://chemed.chem.purdue.edu/genchem/topicreview/b p/ch8/valenceframe.html http://chemed.chem.purdue.edu/genchem/topicreview/b p/ch8/valenceframe.html

41. SOL covered during lessons CH 3 a, c, d