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Chapter 7

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Chapter 7

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  1. Chapter 7 Chemical Formulas and Chemical Compounds

  2. Section 1: Chemical Names & Formulas • There are literally thousands of chemicals • Isn’t always best to use common names for chemicals (calcium carbonate is limestone, sodium chloride is salt, and hydrogen oxide is water) • Common names don’t give information about chemical composition.

  3. Section 1: Chemical Names & Formulas • Significance of chemical formulas: • Gives relative number of atoms of each kind of element. • Subscripts: small numbers to the right that tell the number of atoms • If no subscript then it is understood to be 1 • H2SO4 • 2 hydrogens, 1 sulfur, 4 oxygen

  4. Section 1: Chemical Names & Formulas • When parentheses are used you must multiply inside and out. • Al2(SO4)3 (2 Aluminums, 3 Sulfurs, 12 Oxygens)

  5. Section 1: Names of Binary Compounds • Binary compoundsare those formed from only 2 elements. • To write their formulas the positive ion is written first and then the negative. • To name them use the complete name of the positive ion and add the negative ion name but change the ending to “-ide.” (Sulfur becomes sulfide, oxygen becomes oxide, phosphorous becomes phosphide)

  6. Section 1: Formulas of Binary Compounds • To write the formula of a compound you must consider the charges and multiply by adding subscripts so that the overall charge on the compound is zero. • Ex: zinc is (2+) and sulfur is (2-) so: • ZnS • Name: Zinc Sulfide

  7. Section 1: Formulas of Binary Compounds • Ex: zinc is (2+) and iodine is (1-) so: • Zn 2+ I 1- • ZnI2 • Name: Zinc Iodide 2+ 2- ( = 0 ) Subscripts1 2

  8. Section 1: Formulas of Binary Compounds • How do you know the charge? • Use the valence electrons • Group 1 = 1+, Group 2 = 2+, 3+, 4± • Group 15 = 3-, Grp. 16 = 2-, Grp. 17 = 1- • May use charge chart (page 205) for transition metals.

  9. Assignment: • Worksheet: Writing formulas and names for binary compounds.

  10. Section 1: Stock System of Nomenclature • Some transition metals have more than one possible charge: • Ex. Copper: Cu+ and Cu2+ Iron: Fe2+ and Fe3+ Lead: Pb+3 and Pb+4 Tin: Sn+2 and Sn+4

  11. Section 1: Stock System of Nomenclature • The charges of these elements must be represented in the name of the compounds. • Charges are provided by using Roman numerals in the names • Ex: Iron (II) oxide and Iron (III) oxide • Formulas: FeOFe2O3

  12. Section 1: Stock System of Nomenclature • How do you know how to write the formula??? • Iron (II) combines with oxygen Fe2+ O2- (charges equal zero so FeO) • Iron (III) combines with oxygen Fe3+ O2- (add subscripts and multiply to equal zero 2 3

  13. Section 1: Stock System of Nomenclature • How do you know how to write the name if you only see the formula??? • CuBr2 • The name is Copper Bromide but is it Copper (I) Bromide or Copper (II) Bromide???

  14. Section 1: Stock System of Nomenclature Cu Br2 Then +1 -2 ≠ 0 If charges are +1 -1

  15. Section 1: Stock System of Nomenclature Cu Br2 Then +2 -2 = 0 If charges are +2 -1

  16. Assignment • Worksheet: Naming and Writing Formulas for Compounds Using the Stock System

  17. Section 1: Naming Binary Molecular Compounds • Molecular compounds are those in which the elements are close together on the periodic table. • Ex: • Nitrogen and Oxygen • Carbon and Oxygen • Sulfur and Oxygen • Phosphorus and Chlorine

  18. Section 1: Naming Binary Molecular Compounds • Ex: Compounds of Nitrogen and Oxygen • N2O • NO • NO2 • N2O3 • N2O5 • Newer method of naming is to use the stock system with Roman Numerals. • Old traditional method uses prefixes.

  19. Section 1: Naming Binary Molecular Compounds Prefixes indicate the number of atoms in the compound • 1 atom: Mono • 2 atoms: Di • 3 atoms: Tri • 4 atoms: Tetra • 5 atoms: Penta • 6 atoms: Hexa • 7 atoms: Hepta • 8 atoms: Octa • 9 atoms: Nona • 10 atoms:Deca

  20. Section 1: Naming Binary Molecular Compounds • The less electronegative element is written first and is given a prefix only if it has more than one atom in the formula. • Next element has a prefix indicating the number of atoms and ends typically with “ide.” • Examples: • N2O • NO • NO2 • N2O3 • N2O5 Dinitrogen Monoxide Nitrogen Monoxide Nitrogen Dioxide Dinitrogen Trioxide Dinitrogen Pentoxide

  21. Assignment • Worksheet: Naming and Writing Molecular Compounds Using Prefixes

  22. Section 1: Compounds with Polyatomic Ions • Many compounds are composed of polyatomic ions (a group of covalently bonded atoms that carry a charge). • Examples of polyatomic ions: • Sulfate (SO4)2- • Nitrate (NO3) – • Phosphate (PO4)3- • Carbonate (CO3)2- • Dichromate (Cr2O7)2- • Ammonium (NH4)+

  23. Section 1: Compounds with Polyatomic Ions • Most polyatomic ions end with “–ate” or “-ite” but there are a few exceptions: • Cyanide (CN)- • Hydroxide (OH)- Note of caution: Don’t confuse these with binary compounds since they end in “ide.”

  24. Section 1: Naming Compounds with Polyatomic Ions • Simply write the complete name of the positive element and the name of the polyatomic ion. • KNO3 = Potassium Nitrate • CaSO4 = Calcium Sulfate • Al(OH)3 = Aluminum Hydroxide

  25. Section 1: Writing Compounds with Polyatomic Ions • Writing the formulas for these compounds are a little trickier. • Make sure that you treat the polyatomic ion as a whole unit and do not change its subscripts! • (SO4)2-= 1 sulfate ion • (SO4)2-2 = 2 sulfate ions NOT… (S2O8)

  26. Section 1: Writing Compounds with Polyatomic Ions • Examples: • Potassium nitrate Totals: 1+ and 1- = 0 Charges: + - Symbols: K (NO3) Final Formula: KNO3 (no parenthesis needed since only 1 ion is required

  27. Section 1: Writing Compounds with Polyatomic Ions • Examples: • Aluminum Sulfate Totals: 6+ and 6- = 0 Charges: 3+ 2- Symbols: Al (SO4) Add Subscripts: 2 3 Final Formula: Al2(SO4)3 (parenthesis must be used to show 3 sulfate ions)

  28. Section 1: Writing Compounds with Polyatomic Ions • Polyatomic ions may be paired with transition metals that have multiple charges. • Ex: Copper (II) and sulfate = CuSO4 • But Copper (I) and sulfate = Cu2SO4 • When naming them the Roman numeral must be included. • Fe3(PO4)2 = Iron (II) Phosphate

  29. Assignment: • Worksheet: Naming and Writing Formulas for Compounds Containing Polyatomic Ions

  30. Section 1: Naming Acids and Salts • Memorize the formulas for the common acids. • All begin with one or more H atoms. • Sulfuric Acid H2SO4 • Hydrochloric Acid HCl • Nitric Acid HNO3 • Phosphoric Acid H3PO4 • Carbonic Acid H2CO3

  31. Section 1: Naming Acids and Salts • Binary acids contain only 2 elements • Example: Hydrochloric acid HCl • Oxyacids contain hydrogen, oxygen, and one other element • Example: Sulfuric acid H2SO4

  32. Section 1: Naming Acids and Salts • When acids have less oxygen atoms than normal the names change: • Normal HClO3 is chloric acid • Loss of 1 oxygen atom HClO2 is chlorous acid • Loss of 2 oxygen atoms HClO is hypochlorous acid • An extra oxygen atom HClO4 is perchloric acid

  33. Section 1: Naming Salts • Any ionic compound composed of a cation and the anion from an acid is referred to as a salt. • Example: • NaCl (anion from hydrochloric acid) • CaSO4 (anion from sulfuric acid)

  34. Section 2 Oxidation Numbers

  35. Section 2: Oxidation Numbers • Oxidation numbers are numbers assigned to the atoms in a molecular compound or ion that indicates the general distribution of electrons among bonded atoms. • Oxidation numbers are not actual charges. • Oxidation numbers are useful in naming compounds and writing formulas. -1 +1 +2 -2 +3

  36. Section 2: Oxidation Numbers • Rules for assigning oxidation numbers: • Atoms in a pure element have an oxidation number of zero – O2 Ox. # = 0 • Fluorine always has ox. # of -1 • Oxygen almost always has ox. # of -2 except in peroxides such as H2O2 – then it is a -1.

  37. Section 2: Oxidation Numbers • (Rules continued): • Hydrogen’s ox. # is +1 unless it is with metals – then it is -1 • The sum of the ox. # in molecules must be zero, but in polyatomic ions, it is equal to the ions charge.

  38. Section 2: Oxidation Numbers • What are the oxidation numbers for each atom in these compounds? UF6 : Fluorine is -1 x 6 = -6 Uranium +6 {+6 + (-6)} = 0 H2SO4 : Oxygen is -2 (x 4 = -8) Hydrogen is +1 (x 2 = +2) so Sulfur has to be +6 { (+6) + (+2) + (-8) }= 0

  39. Section 2: Oxidation Numbers • What are the oxidation numbers for the chlorate polyatomic ion? ClO3- : Oxygen is -2 x 3 = -6 Chlorine must be +5 { (+5) + (-6)} = -1 (the ion’s charge)

  40. Section 2: Oxidation Numbers • Assignment: • Page 219, question 1, A-K

  41. Section 3: Using Chemical Formulas

  42. Section 3: Using Chemical Formulas • With a chemical formula, you can calculate many characteristic values for a compound. • Formula Mass: • Compounds have masses – just like elements.

  43. Section 3: Using Chemical Formulas • Formula Mass: • The formula mass of any molecule, formula unit, or ion is the sum of the average atomic masses of all the atoms represented in its formula. • To find the mass of a compound simply add the masses of the atoms that make up the compound. Units are amu’s.

  44. Section 3: Using Chemical Formulas • To find the formula mass of sulfuric acid (H2SO4): element # of atoms x mass (to 2 decimals) H 2 1.01 = 2.02 amu S 1 32.01=32.01 amu O 4 16.00=64.00 amu 98.03 amu

  45. Section 3: Using Chemical Formulas • To find the formula mass of Calcium Nitrate Ca(NO3)2 element # of atoms x mass = Ca 1 40.08 =40.08 amu N 2 14.01 =28.02 amu O 6 16.00= 96.00 amu 164.10 amu

  46. Section 3: Using Chemical Formulas • Molar Mass • The mass of a mole of any substance is equal to its formula mass – except instead of amu’s it is in grams. • Formula mass of sulfuric acid = 98.03 amu • Molar mass of sulfuric acid = 98.03 grams

  47. Percentage Composition • It is sometimes useful to know what the percentage of a compound is an element. • What percentage of water is oxygen? H: 1.01 x 2 = 2.02 O: 16.0 x 1 = 16.0 Molar Mass= 18.02 g 16.0 ÷18.02 = 88.79%

  48. Section 3: Using Chemical Formulas • Molar Mass can be used as a conversion factor. 98.03 grams 1 mole H2SO4 or 98.03 grams 1 mole H2SO4

  49. Section 3: Using Chemical Formulas • How many moles are there in 25 g of H2SO4? 1 mole H2SO4 25 g H2SO4 x = 0.255 mol 98.03 grams

  50. Section 3: Using Chemical Formulas • What is the mass of 4.2 moles of H2SO4? 98.03 g H2SO4 4.2 mol H2SO4 x = 411.73 mol 1 mol H2SO4