Revised: 2017-11-24 Name: Date: Per: Nomenclature Review Chemical Names and Formulas Part A – Naming Binary Ionic Compounds (compounds with only two elements – one metal, one non-metal) Rules: 1. Name the first element (the name of the metal). 2. Name the second element changing the ending to –ide. Examples: NaCl = sodium chloride Al2O3 = aluminum oxide Symbol Binary Name Symbol Binary Name Symbol Binary Name H hydride F fluoride As arsenide B boride Si silicide Se selinide C carbide P phosphide Br iodide N nitride S sulfide Te telluride O oxide Cl chloride I iodide DIRECTIONS: Name the following binary compounds. 1. Al4C3 aluminum carbide 2. NaF sodium fluoride 3. ZnS zinc sulfide 4. Al2S3 aluminum sulfide 5. Ca3P4 calcium phosphide 6. BaCl2 barium chloride 7. MgO magnesium oxide 8. KI potassium iodide 9. Ag2S silver sulfide 10. LiH lithium hydride DIRECTIONS: Write the formulas for the following binary compounds. 11. Sodium bromide NaBr . 12. Potassium fluoride KF 13. Cesium iodide CsI 14. Aluminum nitride AlN 15. Calcium oxide CaO 16. Magnesium sulfide MgS 17. Lithium iodide LiI 18. Hydrogen chloride HCl 19. Silicon carbide SiC 20. Barium oxide BaO Part B – Assigning Oxidation Numbers Rules: 1. The oxidation number of atoms in their elemental form is equal to 0 (zero). 2. The oxidation number of a monatomic ion equals the ionic charge. 3. Hydrogen has an oxidation number of +1 in most compounds (-1 in hydrides). 4. Oxygen has an oxidation number of -2 in most compounds. 5. In binary compounds in which they are the positive element (cation): • Group 1 (IA) elements are always +1 • Group 2 (IIA) elements are always +2 • Group 13 (IIIA) elements are usually +3. • Group 14 (IVA) elements are usually +4 6. In binary compounds in which they are the negative element (anion): • Group 14 (IVA) elements are usually -4 • Group 15 (VA) elements are usually -3 • Group 16 (VIA) elements are usually -2 • Group 17 (VIIA) elements are usually -1 7. The sum of all oxidation numbers of all ions (atoms) in a compound must be zero. Example: Fe2O3 Rule #4 says that O (oxygen) has an oxidation number of -2. There is no rule for Fe (iron), so apply Rule #7. 2 (Fe) + 3 (O) = 0 2 (Fe) + 3 (-2) = 0 2 (Fe) + -6 = 0 2 (Fe) = 6 (Fe) = +3Revised: 2017-11-24 Name: Date: Per: Nomenclature Review Chemical Names and Formulas DIRECTIONS: Assign an oxidation number to the underlined element in each of the following compounds. Remember – an oxidation number is for one atom of the atom. 1. FeO +2 . 2. Al4C3 +3 . 3. CuI +1 . 4. FeBr3 +3 . 5. CoCl3 +3 . 6. Cu2O +1 . 7. SnBr4 +4 . 8. PbO2 +4 . 9. AgCl +1 . 10. Cr2O3 +3 . Part C – Naming Binary Ionic Compounds with Multiple Oxidation State Metals Note: Most metals have multiple possible oxidation states that allow them to become stable. Usually the metal will have one oxidation state that that is better than the others, but the metal will form other oxidation states if the best one cannot be reached. The most common oxidation states of metals are listed on the common ion chart. Rules: 1. Name the first element (the name of the metal). 2. Check the metal for multiple oxidation states using the periodic table or common ion chart. If more than one oxidation state is possible, calculate the correct oxidation state using the process in Part B. Indicate the oxidation number of the metal by writing it in parentheses following the metal’s name using Roman numerals. 3. Name the second element changing the ending to –ide. Examples: CuBr2 = copper (II) bromide CuBr = copper (I) bromide DIRECTIONS: Name the following binary compounds. Remember to include the oxidation state in the name only if the metal has more than one oxidation state. 1. AuCl3 gold (III) chloride 2. CdS cadmium (II) sulfide 3. BaBr2 barium bromide 4. Fe2O3 iron (III) oxide 5. K2O potassium oxide 6. Cr2O3 chromium (III) oxide 7. MnCl2 manganese (II) chloride 8. Cu2O copper (I) oxide 9. Ag2S silver sulfide 10. PbCl2 lead (II) chloride DIRECTIONS: Write the formula for the following compounds. 11. Iron (II) oxide FeO 12. Barium hydride Ba(OH)2 13. Tin (IV) chloride SnCl4 14. Silver nitride Ag3N 15. Calcium chloride CaCl2 16. Manganese (II) fluoride MnF2 17. Strontium nitride Sr3N2 18. Gold (I) sulfide Au2S 19. Mercury (II) oxide HgO 20. Chromium (III) sulfide Cr2S3 21. Potassium oxide K2O 22. Cobalt (II) iodide CoI2 23. Thallium (I) selenide Tl2Se 24. Aluminum bromide AlBr3 25. Copper (II) iodide CuI2 26. Zinc phosphide Zn3P2 27. Magnesium hydride MgH2 28. Nickel (II) bromide NiBr2 29. Iron (III) oxide Fe2O3 30. Lithium oxide Li2ORevised: 2017-11-24 Name: Date: Per: Nomenclature Review Chemical Names and Formulas Part D – Identifying Polyatomic Ions in Formulas Note: A polyatomic ion is a charged particle made up of atoms of more than one element. The atoms in the ion are bonded covalently as the whole group gain or lose electrons (causing the charge). Usually polyatomic ions are anions (negatively charged), but polyatomic cations (positively charged) do occur. Ammonium (NH4+) is the most common example of a polyatomic cation while mercury (I) (Hg22+) is a less common example. Polyatomic ions have unique names. The most common polyatomic ions are listed on the common ion chart. Rules: 1. Polyatomic ions are discrete units of matter with unique formulas and names. 2. If a compound contains a multiple of a polyatomic ion (two nitrate ions, three hydroxide ions, etc.), the formula of the polyatomic ion is placed in parentheses and the multiple is written as a subscript following the parentheses. Examples: Na2CO3 = “CO3” represents the polyatomic ion “carbonate” MgC2O4 = “C2O4” represents the polyatomic ion “oxalate” Al(OH)3 = “OH” represents the polyatomic ion “hydroxide” – the “( )3” indicate there are 3 of them DIRECTIONS:
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