1. Chapter 7 Chemical Formulas & Naming(Nomenclature)

2. Chemical Formulas Chemical formulas indicate the relative # of each type of atom in a compound Ex: C 8 H 18 (octane)Al 2 (SO 4 ) 3 8 Carbons18 Hydrogens 2 Aluminum1 Sulfur & 4 Oxygen 3x everything in ( )

3. Monatomic Ions Monatomic Ions are formed from a single atom (+) positive ions = cations = lose e - (-) negative ions = anions = gain e - Atoms gain or lose electrons in an attempt to achieve a full valence shell (octet)

4. Naming Monatomic Ions Cations(+) K + : potassium cation Mg 2+ : magnesium cation D-Block Cations Cu + : copper(I) Fe 3+ : iron(III) * Use roman # to indicate pos. charge Anions(-) F - : fluoride N 3- : nitride O 2- : oxide * Element name ending in ide

5. Binary Ionic Compounds Binary ionic compounds are made from the combination of cations and anions The charges of the positive and negative ions must balance out to zero(neutral). Ex: Mg 2+ + Br -  MgBr 2

6. Crossing Over Method 1)Write the cation followed by the anion 2) Cross over charges to form subscripts 3) Check to make sure charges are balanced

7. Naming Binary Ionic Compounds Combine the names of both ions Ex: Al 2 O 3 = aluminum oxide AgCl = silver chloride CaBr 2 = calcium bromide NaCl = sodium chloride

8. Stock System of Nomenclature Use Roman #’s to indicate the charge of the cation(+) Ex: CuCl 2 = copper(II) chloride FeO = iron(II) oxide Fe 2 O 3 = iron(III) oxide SnF 4 = tin(IV) fluoride ZnBr = zinc(I) bromide

9. Polyatomic Ions Polyatomic ions are molecules with an overall positive or negative charge Oxyanions- polyatomic ions containing oxygen Ex: NO 2 - = nitrite NO 3 - = nitrate ClO 2 - = chloriteClO - = hypochlorite ClO 3 - = chlorateClO 4 - = hyperchlorate

10. Naming Binary Molecular Compounds (covalent bonds) Use prefixes to indicate the # of each type of element in the molecule 1)Write the least electronegative element 1 st 2)Use a prefix is there is more than 1 of the 1 st element 3)Write the second element always using a prefix to indicate the # of atoms present

11. Prefixes 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca

12. Molecular Compounds: Ex: P 4 O 10 = tetraphosphorus decaoxide N 2 O = dinitrogen monoxide SO 3 = sulfur trioxide

13. Acids & Salts Binary Acids: contain H & a halogen(17) Ex: HCL = hydrochloric acid Oxyacids: contain H, O, & a 3 rd element Ex: H 2 SO 4 = sulfuric acid HNO 3 = nitric acid Salts: formed from a cation and an anion of an acid Ex: HCl + NaOH  NaCl + H 2 O (Acid)(Base)(Salt)

14. Oxidation #’s (states) Rules: 1)The oxidation number of a free(single) element is always 0(zero) 2)The oxidation number of a monatomic ion equals the charge of the ion 3)The usual oxidation number of hydrogen(H) is +1 4)The oxidation number of oxygen(O) in compounds is usually -2 5)The oxidation # of other elements is based on the elements group # 6)The sum of the oxidation numbers of all of the atoms in a neutral compound is 0 7)The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion

15. Examples

16. Chemical Formulas Review Formula Mass- sum of the atomic masses of all elements in a compound, units (amu) Ex: H 2 O Molar Mass- mass of one mole of a substance, numerically equivalent to formula mass, different units (g/mol) Ex: H 2 SO 4

17. % Composition The percentage by mass of each element in a compound Mass of Element -------------------------- x 100 = % Composition Mass of Compound Ex: H 2 SO 4

18. Deriving Chemical Formulas Empirical Formula- the symbols of the elements in a compound w/ subscripts showing the smallest whole # ratio between the atoms –Ex: C 6 H 12 O 6  CH 2 O Chemical Formula Empirical Formula

19. Determining Empirical Formulas 1)Assume 100.0 g sample mass (if no sample mass is given) 2) Convert % Composition to Mass for each element 71 % A = 71 g of “A” 29 % B = 29 g of “B”

20. Determining Empirical Formulas 3) Determine the # of Moles of each element using the Molar Mass 71.0 g of A / 10.0 g/mol = 7.10 mol of A 29.0 g of B / 2.01 g/mol = 14.5 mol of B 4) Determine Mole Ratio by dividing both values from step 3 by the lower # A: 7.10 / 7.10 = 1.0 B: 14.5 / 7.10 = 2.0

21. *** If the mole ratio is NOT a whole # ratio after dividing then… Multiply both #’s by the same factor to get a whole # ratio –Ex: A = 1.5 x 2 = 3.0 B = 1.0 x 2 = 2.0 A = 2.25 x 4 = 9.0 B = 1.0 x 4 = 4.0

22. Practice Empirical Formulas 1)32.38% Sodium(Na) 22.65% Sulfur(S) 44.97% Oxygen(O) 2) 10.15 gram sample Contains 4.43 g of Phosphorus(P)

23. Determining Empirical Formulas 1)Convert % of each Element to Grams(g) 2)Convert Grams(g) to Moles(mol) 3)Divide Moles by Lowest # 4)Determine Whole # Ratio Ex: 47.3% Carbon 10.6% Hydrogen 42.1% Oxygen

24. Calculating Molecular Formulas 1)Divide: Actual Mass / Empirical Formula Mass 2) Multiply: Multiply the # from step 1 by the subscript of each element in the empirical formula Ex: BH 3 Actual Mass = 27.67 g/mol

25. END OF CHAPTER 7 NOTES