1. 1 Naming and Formula Writing for Ionic and Covalent Compounds

2. 2 Objectives…  List differences among ionic and covalently bonded compounds  Identify polyatomic and monatomic ions and name them properly  Write the chemical formulas and names for ionic compounds  Write the chemical names and formulas of acids  Write the chemical names and formulas of molecules

3. 3 Review of Ionic Compounds (Crystals)  Transfer of electrons  Made from a metal and a non-metal  Metals lose electrons, nonmetals gain electrons  Chemical formula is arranged in the smallest whole number ratio (empirical formula)  Formula unit: the smallest repeating pattern within a crystal  Even though ionic compounds are made up of ions, they ARE ELECTRICALLY NEUTRAL!

4. 4 Ions can be monatomic or polyatomic:  Monatomic: Made up of a single atom  Polyatomic: Made up of multiple atoms Remember… Group #: 1 2 13 14 15 16 17 18 Charge: +1 +2 +3 X -3 -2 -1 X

5. 5  Monatomic cations have the same name as the element  Example: Na +1 = Sodium ion, Ca +2 = Calcium ion  Monatomic anions have the ending of the element name changed to“-ide”  Example: Cl -1 =chloride ion, O -2 =oxide ion

6. 6 Polyatomic ions: (See polyatomic ion sheet) are made up of two or more elements covalently bonded together with an overall positive or negative charge.

7. 7 Review Molecular Compounds (Molecules) SSharing of electrons MMade from nonmetals only MMolecules of the same compound are IDENTICAL and INDEPENDENT of each other CChemical formula indicates the exact makeup of one molecule (molecular formula)

8. 8 Rules for Formula Writing (Ionic Compounds) Since all compounds are neutral, figure out how many of each ion is needed to make a neutral compound. (Neutral means having a net zero charge.)

9. 9 Examples Mg and Cl Mg +2 Cl -1 Cl -1 Al and O Al +3 O -2 O -2 MgCl 2 Al 2 O 3

10. 10 More Examples Ca and S Ca +2 S -2 Ca and (AsO 4 ) Ca +2 (AsO 4 ) -3 Ca +2 CaS Ca 3 (AsO 4 ) 2 Parentheses are needed if there is more than one of the polyatomic ions!

11. 11 Aluminum Sulfite Al +3 (SO 3 ) -2 (SO 3 ) -2 Chemical Name to Formula Al 2 (SO 3 ) 3

12. 12 More Examples Silver Sulfate Ag ? (SO 4 ) -2 Ag +1 Ag 2 (SO 4 ) Ag +1 Use your Periodic Table to determine the charge on transition metals!

13. 13 More Examples Nickel (II) Nitrate Ni +2 (NO 3 ) -1 (NO 3 ) -1 Ni(NO 3 ) 2 The Roman Numeral will always tell you the charge and it will always be positive

14. 14 More Examples Iron (III) Chloride Fe +3 Cl -1 Cl -1 FeCl 3

15. 15 Chemical Formula to Name Recall the format for chemical formulas… (name of cation – metal) (name of anion – nonmetal) Examples: KBr CaI 2 Na(SO 4 ) Potassium Bromide Calcium Iodide Sodium Sulfate

16. 16 For elements with more than one possible charge (transition metals)… Use the charge on the anion to determine the charge on the cation! Example: Fe 2 (CrO 4 ) 3 -2(3)=-6+6/2=+3 Iron (III) Chromate

17. 17 Example: Sn(CO 3 ) 2 -2(2)=-4+4 Tin (IV) Carbonate

18. 18 Example: Cu 3 P -3+3/3=+1 Copper (I) Phosphide

19. 19 Criss Cross Short Hand Method Just use the number of the charge (not the charge itself) and criss cross as shown below. You will note that the overall charge is neutral, as the total positive charge of the metals offsets the total negative charge of the nonmetal. 19

20. 20 Naming Acids Acids are ionic compounds that contain H +1 as their cation. Acids are named based on their anion.

21. 21 If the anion ends in… “-ide”  hydro ___ ic acid Example: H 2 S Hydrosulfuric acid Sulfide (S -2 ) is the anion! The root name of the anion goes here (remove-ide)

22. 22 If the anion ends in… “-ate”  ___ ic acid Example: H 2 SO 4 Sulfuric acid Sulfate (SO 4 -2 ) is the anion!

23. 23 If the anion ends in… “-ite”  ___ ous acid Example: H 2 SO 3 Sulfurous acid Sulfite (SO 3 -2 ) is the anion!

24. 24 Mnemonic devices to help you remember Acid Nomenclatire #1) “eight is great”er number of oxygens “ite” is slight” less number of oxygens #2) Chemistry “Diseases” (joke..I am suffering from eight-ic-ite-ous or).... “ate” -ic “ite”-ous SO 4 = sulfate SO 3 = sulfite H 2 SO 4 H 2 SO 3 Sulfuric Acid Sulfurous Acid or lessous icmore NO 2 = nitrite NO 3 = nitrate HNO 2 nitrous acid HNO 3 nitric acid PO 3 = phosphite PO 4 = phosphate H 3 PO 3 = phoshporous Acid H 3 PO 4 = phosphoric Acid

25. 25 Formula Writing Use the reverse to determine the anion and balance out the charges for a neutral compound.

26. 26 Examples Hydrophosphoric acid H +1 P -3 H +1 All acids contain a H +1 charge as the cation! Ask yourself, “What was the original ending?” H3PH3P Original ending “ide”

27. 27 Examples Chromic acid H +1 (CrO 4 ) -2 H +1 Original ending “ate” H 2 CrO 4

28. 28 Naming Molecular Compounds Recall, covalently bonded molecules are made up of nonmetals only.

29. 29 Prefixes mono – 1 hexa – 6 di – 2 hepta – 7 tri – 3 octa – 8 tetra – 4 nona – 9 penta – 5deca - 10 Must be memorized!

30. 30 Rules  Use the prefixes to indicate how many of each element is in one molecule.  Change the ending on the second element to “-ide”

31. 31 Examples N2O4N2O4 CO SiO 2 If there is only one of the first element, do not write mono. dinitrogen tetraoxide carbon monoxide silicon dioxide

32. 32 Formula Writing  Use the prefixes to determine the subscripts Examples: trisulfur hexafluoride carbon pentaoxide tetraphosphrous dioxide S3F6S3F6 DO NOT REDUCE! CO 5 P4O2P4O2