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Chemical Names and Formulas: Writing, Naming, and Oxidation Numbers

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  1. Chapter 7 Chemical Formulas and Chemical Compounds

  2. 7.1 Chemical Names and Formulas • Chemical formulas show the relative #’s of atoms in a chemical compound • Examples • C12H22O11 C = 12 H = 22 O = 11 • Pb(NO3)4 Pb = 1 N = 4 O = 12 • (NH4)2CrO4 N = 2 H = 8 Cr = 1 O = 4

  3. Chemical Names and Formulas • Naming monatomic ions • End in -ide • Look up the charges on the periodic table • F vs F-1 • F • F-1 • S vs S-2 • S • S-2 Fluorine Fluoride Sulfur Sulfide

  4. Writing Formulas for Ionic Compounds • Cations – positive ions (metal) • metal • Anions – negative ions • nonmetal • Charge on compound = 0

  5. Writing Formulas for Ionic Compounds • Binary ionic compounds • Rules • if the charges on the ions are the same, drop ‘em • if the charges are different, criss-cross • Same charges – • Na+1 Cl-1 - • Mg+2 O-2+- • Different charges- • Na+1 S-2- • Mg+2 Cl-1- NaCl Sodium Chloride MgO Magnesium Oxide Na2S Sodium Sulfide MgCl2 Magnesium Chloride

  6. Writing Formulas for Ionic Compounds • Ternary Ionic compounds • Metal + (Polyatomic ion) • When naming, do not use the ending –ide • Sodium Nitrate • Sodium Carbonate • Aluminum Nitrate Na+1 NO3-1 Na(NO3) Na+1 CO3-2 Na2(CO3) Al+3 NO3-1 Al(NO3)3

  7. Writing Formulas for Ionic Compounds • Aluminum Phosphate • Aluminum Bihypophosphite • Aluminum Carbonate Al+3 PO4-3 Al(PO4) Al2(HPO2)3 Al+3 HPO2-2 Al2(CO3)3 Al+3 CO3-2

  8. Writing Names for Ionic Compounds • Front name – positive (cation – metal) • Back name – negative (anion – nonmetal) • Binary ionic compounds – composed of only 2 types of elements ( M + NM) – end in -ide • NaCl • MgCl2 • Al2O3 • NaH Sodium Chloride Magnesium Chloride Aluminum Oxide Sodium Hydride

  9. The BIG Lie • Stock System – use Roman Numerals for naming compounds with metals that have multiple charges (the transitions!) Cory Matthews Loves Topanga Intensely

  10. The BIG Lie • More exceptions to the Lie! • Ag is always = +1 charge • DON’T write Ag I • Zn always = +2 charge • DON’T write Zn II

  11. Practice! • Sn3N2 • AgOH • PbCO3 • Zn(OH)2 • Fe2(SO4)3 Tin (II) Nitride Silver Hydroxide Lead (II) Carbonate Zinc Hydroxide Iron (III) Sulfate

  12. More Practice!! • Cu(HSO2)2 • CuSO2 • CuHSO2 Copper (II) Bihyposulfite Copper (II) Hyposulfite Copper (I) Bihyposulfite

  13. Practice! • LiClO3 • LiClO2 • CaCO3 • Ca(HCO2)2 • Fe(NO3)3 Lithium Chlorate Lithium Chlorite Calcium Carbonate Calcium Bicarbonite Iron (III) Nitrate

  14. Writing Names for Molecular Compounds • Molecular Compounds – covalent compounds • 2 nonmetals • To name, we use prefixes don’t use the prefix Mono on the first atom

  15. Writing Names for Molecular Compounds • Prefix-name prefix-name-ide • CO • CO2 • PCl3 • CBr4 • N2O5 • SF6 Carbon Monoxide Carbon Dioxide Phosphorous trichloride Calcium tetrabromide Dinitrogen pentoxide Sulfur hexafluoride

  16. Writing Formulas for Molecular Compounds • The prefixes = the subscripts. • Do NOT look at the charges. • Sulfur Dioxide • Disulfur Trioxide • Dinitrogen pentoxide SO2 S2O3 N2O5

  17. Naming Acids • Acid - when a solution yields H+ ions in solution • 2 types • Binary • H and one other type of atom • ternary (sometimes called oxy) • acids that have H with a polyatomic ion

  18. Naming Binary Acids • Rules • Hydro__(begninng of name)__ic acid • Ex. HCl • Hydrochloric acid • HBr • HF • H2S • H3P Hydrobromic acid Hydrofluoric acid Hydrosulfuric acid Hydrophosphoric acid

  19. Writing Formulas from Names for Acids • Do the criss-cross • Ex. Hydronitric acid • H+1 N-3 H3N • Hydroiodic acid • H+1 I-1 HI • Hydrosulfuric acid • H+1 S-2 H2S

  20. Naming Ternary Acids • H + polyatomic • Rules • Do NOT start with hydro- • If the ending of polyatomic is –ate -ic + acid • If the ending of polyatomic is –ite -ous + acid • Ate/ite ic/ous • Example: • H2SO4 • H+ and SO4-2 – sulfate • sulfuric acid

  21. Naming Ternary Acids • H2SO3 • HClO4 • HClO3 • HClO2 • HClO

  22. Formulas for Ternary Acids • Use the criss-cross method • Nitric acid • Phosphorous acid

  23. 7.2 Oxidation Numbers • Since electrons are shared, there is no definite charge - we assign the more electronegative element the “apparent” negative charge - this is known as the oxidation # • oxidation numbers can also be positive. • oxidation # - a number assigned to an atom to show the distribution of elements

  24. Oxidation Numbers • Rules • Free elements = 0 • Ex. Mg = O • Ions= charges • Ex. F = -1 S = -2 • Oxygen (0) = -2 • except in peroxides (H2O2) O = -1 • H = +1 • except in metal hydrides (MgH2, NaH) H = -1

  25. Oxidation Numbers • ….Rules • More electronegative atom gets a (-) charge • Ox #’s add up to 0 in compounds • Ox #’s = the charge in polyatomic ions

  26. Oxidation # Practice • FeO (Iron II Oxide) O = -2 Fe = ? • Fe2O3 (Iron III oxide) O = -2 Fe = ? -2 + x = 0 x = 2 3(-2) + 2x = 0 x = 3

  27. Oxidation # Practice • H2SO4 (Hydrogen Sulfate or Sulfuric Acid) O = -2 H = +1 S = SO4-2 X + 4(-2) = -2 X = 6

  28. Oxidation # Practice • H2SO3 (Hydrogen sulfite or sulfurous acid) • H = • S = • O = +1 +4 -2 SO3-2 X + 3(-2) = -2 X = 4

  29. Oxidation # Practice • H2Cr2O7 (Hydrogen Dichromate or Dichromic Acid • H = +1 • Cr = • O = -2 • NO3-1 (Nitrate) • N = • O = -2 • MgH2 (Magnesium hydride) • Mg = • H = -1 +6 +4 +2

  30. 7.3 Using Chemical Formulas • Step 1 – be able to calculate molar mass (aka – formula mass, molecular weight, atomic weight, atomic mass, gram formula weight, etc.) • Add atomic weights from the periodic table • round to the nearest 10th place • Examples • CH4 • MgSO4· 7H2O 12.0 + 4(1.0) = 16.0 g/mol 24.3 + 32.0 + 4(16.0) + 14(1.0) + 7(16.0) = 246.3 g/mol

  31. Liters 22.4 L Mole Molar Mass 6.022 x1023 Atoms, molecules, particles Grams 7.3 Using Chemical Formulas • Step 2 – be able to convert between grams, moles, particles, and liters

  32. Using Chemical Formulas • Convert 32.0 g of CH4 to moles, liters, molecules, total atoms, atoms of H • Moles • Liters • Molecules • Atoms • Atoms H

  33. Conversions • Moles CH4 • 32.0 g 1 mole 1 16.0g • Liters • 32.0g 1 mole 22.4L 1 16.0g 1mole 2.0moles 44.8 Liters

  34. Conversions • Molecules • 32.0g 1 mole 6.022 x 1023molecules 1 16.0g 1mole • Atoms • 32.0g 1 mole 6.022 x 1023atoms 5 1 16.0g 1mole • Atoms H • 32.0g 1 mole 6.022 x 1023atoms 4 1 16.0g 1mole 1.20 x 1024molecules 6.00 x 1024molecules 4.80 x 1024molecules

  35. Percent Composition • Percentage Composition - every compound has a certain percentage of each type of atom (we measure it by mass) • Formula • % composition = mass element mass compound X 100 =

  36. Practice - % Composition • Calculate % composition if a compound contains 24 g of Carbon and 64 g of Oxygen • % composition = mass element mass compound • Total mass compound = 24 + 64 = 88 g • % Composition C = 24 x 100 = 27 % C 88 • % Composition O = 64 x 100 = 73% O 88 X 100 =

  37. Practice - % Composition • What is the % composition of Ba(OH)2? • Ba = 137.3 g • O = 2 (16.0) = 32.0g • H = 2 (1.0) = 2.0g • Ba(OH)2 = 171.3 g %Ba = 137.3 171.3 = 80.2% %O = 32.0 171.3 =18.7% %H = 2.0 171.3 =1.1%

  38. Practice - % Composition • What is the % composition of C6H12O6 • C = 6 (12.0) = 72.0g • H = 12 (1.0) = 12.0g • O = 6 (16.0) = 96.0g • C6H12O6 = 180.0g 40.0% 6.7% 53.3%

  39. 7.4 Determining a compound’s empirical and molecular formula • Empirical formula - the lowest whole number ratio of atoms in a compound (simplest formula) • 4 rules to find empirical formula • Cross out the % → change to grams • Divide each by own molar mass • Divide by the smallest number • If needed, multiply by 2 or 3 ONLYif a whole number ratio isn’t the result of step 3

  40. Determining a compound’s empirical formula • Ex1 Calculate the empirical formula if there is 52.17 % C, 13.04% H, and 34.78 % O • 52.17 g C 1 mole = 4.3 = 1 12.0 g • 13.04 g H 1 mole = 13.04 = 1 1.0 g • 34.78 g O 1 mole = 2.17 = 1 16.0 g 2 2.17 6 2.17 1 2.17 C2H6O

  41. Determining a compound’s empirical formula • Ex2 Calculate the empirical formula is there is 26.56 % K, 35.41 % Cr, and 38.03 % O • 26.56 g K 1 mole = 39.1 g • 35.41 g Cr 1 mole = 52.o g • 38.03 g O 1 mole = 16.0 g .68 = 1 x2 .68 .68 = 1 x2 .68 2.38 = 3.5 x2 .68 K2Cr2O7

  42. Determining a compound’s empirical formula • Ex3 Find the empirical formula if a sample contains 5.6 g N and 12.8 g O

  43. Determining a compound’s molecular formula • Rules • Same steps as empirical formula + 3 more • Find the mass of the empirical compound • Divide this mass by the given molecular weight • Multiply the empirical formula by this number

  44. Determining a compound’s molecular formula • Find the molecular formula of a compound (MW = 144.0 g) with 66.67 % C, 11.11 % H, and 22.22 % O • Empirical = C4H8O • Mass empirical = 72.0 g • 144.0 g/72.0 g = 2 • Molecular Formula = C8H16O2

  45. Determining a compound’s molecular formula • Find the molecular formula of that compound that contains 19.80% C, 2.50% H, 66.10% O, 11.60% N and MW = 242.0 • Empirical Formula = C2H3O5N • Mass Empirical = 121 g • 242 g / 121 g = 2 • C4H3O10N2

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