1. Naming Compounds Use handout for notes.

2.  Covalent – show sharing of e -  Ionic – show transfer of e -

4.  These have only 2 elements and are ionic (1 metal and 1 nonmetal).  Write the name of the cation (metal).  Then write the name of the anion (nonmetal) but end it with the suffix –ide.  Examples:  sodium chloride (NaCl)  Calcium oxide (CaO)

5.  Anion nonmetals form ions with negative charges.  Group number – 8 = charge  Examples:  nitrogen: 5-8=-3  Oxygen: 6-8=-2  Iodine: 7-8=-1

6. ElementIon NameIon SymbolIon Charge FluorineFluorideF- ChlorineChlorideCl- BromineBromideBr- IodineIodideI- OxygenOxideO2-O2--2 SulfurSulfideS2-S2--2 NitrogenNitrideN3-N3--3 PhosphorusPhosphideP3-P3--3

7.  The alkali metals (group 1A), alkaline earth metals (group 2A), and aluminum form ions with a positive charge equal to their number of valence electrons (or group number).  Examples: ElementIon symbolCharge PotassiumK++1 CalciumCa 2 ++2 AluminumAl 3 ++3

8.  The transition metals form cations differently.  They form more than 1 type of ion.  Use a Roman numeral next to the ion name to indicate its charge.  Examples: copper can be copper (I) with a +1 charge or copper (II) with a +2 charge

9. ElementIon NameIon Symbol CopperCopper (I)Cu+ CopperCopper (II)Cu 2 + IronIron (II)Fe 2 + IronIron (III)Fe 3 + LeadLead (II)Pb 2 + LeadLead (IV)Pb 4 + ChromiumChromium (II)Cr 2 + ChromiumChromium (III)Cr 3 + TitaniumTitanium (II)Ti 2 + TitaniumTitanium (III)Ti 3 + TitaniumTitanium (IV)Ti 4 + MercuryMercury (II)Hg 2 +

10. Common Ion Charges 1+ 2+3+NA3-2-1- 0

11.  The total (net) charge of an ionic compound must be 0 (zero).  The cation and anion charges must cancel each other out.  Example:  Copper (II) oxide– copper has a charge of +2 and oxygen has a charge of -2  +2 cancels -2 to equal 0

13. 1. Titanium (II) sulfide 2. Copper (I) oxide 3. Lead (II) oxide 4. Lead (IV) nitride 5. Mercury (II) iodide

14. 1. **Titanium (II) sulfide +2-2=0 2. Copper (I) oxide +1-2=-1 3. **Lead (II) oxide +2-2=0 4. Lead (IV) nitride +4-3=+1 5. Mercury (II) iodide +2-1=+1

15.  Copper (1) oxide  CuO  Does not balance to 0  +1-2=-1  But 2 atoms of copper (1) would make it equal 0  Cu 2 0  +1+1-2=0

16.  Mercury (II) iodide +2-1=+1  HgI  Does not balance to 0, but…  HgI 2  +2-1-1=0

17.  Lead (IV) nitride +4-3=+1  PbN  Does not balance to 0, but…  Pb 3 N 4  +4+4+4-3-3-3-3=0  +12-12=0

18. 1. Iron (II) bromide 2. Chromium (III) nitride 3. Lead (IV) oxide 4. Titanium (III) fluoride 5. Lead (IV) phosphide

19. 1. Iron (II) bromideFeBr 2 2. Chromium (III) nitrideCrN 3. Lead (IV) oxidePbO 2 4. Titanium (III) fluorideTiF 3 5. Lead (IV) phosphide Pb 3 P 4

20.  The atoms in polyatomic ions are joined by covalent bonds but have a net + or – charge.  The prefix poly- means many.  Polyatomic ions contain many ions.  Example:  Ammonium contains 1 nitrogen and 4 hydrogen atoms.  -3+1+1+1+1=-3+4=+1 (net charge)

21. Ion NameFormulaIon NameFormula AmmoniumNH 4 +AcetateC2H3O2-C2H3O2- HydroxideOH-PeroxideO22-O22- NitrateNO 3 -PermanganateMnO 4 - SulfateSO 4 2 -Hydrogen sulfateHSO 4 - CarbonateCO 3 2 -Hydrogen carbonate HCO 3 - PhosphatePO 4 3 -Hydrogen phosphate HPO 4 2 - ChromateCrO 4 2 -DichromateCr 2 O 7 2 - SilicateSiO 3 2 -HypochloriteOCl-

22.  I will give you polyatomic ion formulas on the test.  There is a chart in your book on page 173 to use on classwork and homework.

23.  Polyatomic ions can combine with metals to form compounds.  Put the polyatomic ion in parentheses when you write the formula.  Example:  Iron (III) hydroxideFe(OH) 3  Iron has a +3 charge and hydroxide has a -1 charge. You must have 3 hydroxide ions to balance the iron and make the net charge 0.

24. 1. Chromium (II) sulfate 2. Lead (IV) hydroxide 3. Mercury (II) acetate 4. Copper (I) sulfate 5. Titanium (IV) dichromate  See page 173 for the polyatomic ion formulas.

25. 1. Chromium (II) sulfateCr(SO 4 ) 2. Lead (IV) hydroxidePb(OH) 4 3. Mercury (II) acetate Hg(C 2 H 3 O 2 ) 2 4. Copper (I) sulfateCu 2 (SO 4 ) 5. Titanium (IV) dichromate Ti(Cr 2 O 7 ) 2

26.  Consider the following: ◦ Does it contain a polyatomic ion?  -ide, 2 elements  no  -ate, -ite, 3+ elements  yes ◦ Does it contain a Roman numeral?  Check the table for metals not in Groups 1 or 2. ◦ No prefixes!

27.  These are called molecules.  Name the most metallic element (the one that appears the furthest left on the periodic table) first.  If both elements are in the same group, name the one closest to the bottom first.  Add the suffix –ide to the second element.  Example: carbon dioxide

28.  Use the Greek prefixes to indicate how many of each atom are in the molecule. NumberPrefixNumberPrefix 1Mono-6Hexa- 2Di-7Hepta- 3Tri-8Octa- 4Tetra-9Nona- 5Penta-10Deca-

29.  N 2 O 4  Dinitrogen tetraoxide  Di = 2 nitrogens  Tetra = 4 oxygens  Because nitrogen is a group left of oxygen name it first  Because oxygen is named second, change it to end in -ide

30.  NO 2  Mononitrogen dioxide  If there is only 1 of the first element, it is not necessary to use the mono- prefix.  Nitrogen dioxide

31. 1. NO 2 2. P 2 F 4 3. P 2 O 5 4. CO 5. N 2 S 5

32. 1. NO 2 nitrogen dioxide 2. P 2 F 4 diphosphorus tetrafluoride 3. P 2 O 5 diphosphorus pentaoxide 4. COcarbon monooxide 5. N 2 S 5 dinitrogen pentasulfide

33. 1. Nitrogen dioxide 2. Diphosphorus tetrafluoride 3. Carbon dioxide 4. Dihydrogen oxide 5. Dinitrogen tetraoxide

34. 1. Nitrogen dioxideNO 2 2. Diphosphorus tetrafluorideP 2 F 4 3. Carbon dioxideCO 2 4. Dihydrogen oxideH 2 O 5. Dinitrogen tetraoxide N 2 O 4

35. NOF Cl Br I H  The Seven Diatomic Elements Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2