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

Learn how to write and name chemical formulas for ionic and molecular compounds, and understand oxidation numbers. This chapter covers the rules and examples for naming and writing formulas, and assigning oxidation numbers to elements.

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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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