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Chapter 3

Chapter 3. Molecules and Compounds. C + 4 H = CH 4. Molecular model. Molecules and Compounds - Chemical Formulas. 1:1 1:2 1:3 2:3 1:4 etc CO H 2 O NH 3 Al 2 O 3 CH 4. Ethanol. Molecular Models. Molecular Models.

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Chapter 3

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  1. Chapter 3 Molecules and Compounds

  2. C + 4 H = CH4 Molecular model Molecules and Compounds -Chemical Formulas 1:1 1:2 1:3 2:3 1:4 etcCO H2O NH3 Al2O3 CH4

  3. Ethanol Molecular Models

  4. Molecular Models

  5. Some compounds are IONIC - electrons are TRANSFERRED from a metal to a nonmetal; the compound is held together ELECTROSTATICALLY Coulomb’s Law: force = k (n+)(n-)/d2

  6. e- + + e- E. g. carbon dioxide Some compounds are COVALENT - electrons are SHARED between two atoms This commonly occurs for two or more NONMETALS

  7. The Covalent Bond

  8. Ions • An atom or group of atoms with a net charge caused by the net loss or gain of electrons • CATION = positively charged ion • ANION = negatively charged ion

  9. Formation of Cations & Anions

  10. Predicting Whether an Atom Will Form a Cation or Anion in Order to Make an Ionic Compound Metals LOSE electrons to form CATIONS, nonmetals GAIN electrons to form ANIONS

  11. Valence Electrons in Ionic Compounds • The A-group (representative) elements follow the OCTET RULE; they obtain an inert gas valence (outer) shell that contains 8 electrons • Metals - lose # electrons = group numbere.g. Ca  Ca2+ + 2e- (Ar outer shell) • Nonmetals - gain electrons = 8 - group #e.g. N + 3e-  N3- (Ne outer shell)

  12. Valence Electrons for Covalent Compounds • Covalent compounds form between two or more nonmetals • In this case the nonmetals can either LOSE all of their valence electrons,or - • GAIN enough electrons to obtain an OCTET

  13. Examples - SO3 - oxygen (VIA) gains 8-6 = 2 O2- ion forms - sulfur (VIA) loses all 6 S6+ ion forms NOTE: There is a rule that states that oxygen is ALWAYS -2. These rules are coming up!

  14. Example 3.3 - Predicting Ion Charges When Forming Ionic Compounds • Metals lose electrons, Nonmetals gain them • Al = group IIIA metal, so LOSES 3 ELECTRONS Al Al3+ + 3e- • S = group VIA nonmetal, so GAINS 8 - 6 = 2 ELECTRONS S + 2e-  S2-

  15. Polyatomic Ions • Contains 2 or more atoms COVALENTLY bonded, and the complete unit contains a net charge, e.g. nitrate, NO3-

  16. Polyatomic Ion Examples NO2-, CO32-, SO42-, PO43- NO2-ion CO32-ion

  17. Compound held together electrostatically Very strong forces hold the lattice together, so ionic cmpd’s have very high melting points NaCl crystal latticem.p. = 800 oC Ionic Compound = Metal + Nonmetal or a Metal + Polyatomic Ion

  18. Predicting Formulas of Ionic Compounds Balance positive and negative charges to produce a neutral molecule Ca2+ + Cl- Ca2+ + CO32-Ca2+ + PO43-Al3+ + O2-

  19. Oxidation Numbers • A number assigned to each element in a compound in order to keep track of the electrons during a reaction Mg2+ = +2 Cl- = -1 O2- = -2 N3- = -3

  20. Rules for Assigning Oxidation Numbers(Chap. 5, p. 207) Rules higher up take precedence over lower rules • The O.N. for an atom in its pure, uncombined state = 0. • The sum of the O.N.’s for a neutral molecule = 0. For a polyatomic ion, the sum = charge.

  21. Rules cont’d • Group IA = +1Group IIA = +2 • H = +1 UNLESS combined with IA or IIA, then = -1 • Oxygen = -2 • For binary ionic compounds only -Group VA = -3Group VIA = -2Group VIIA = -1

  22. Examples • P4 • Al2O3 • MnO4- • NaH • Na2SO3 • Mg3N2

  23. Chemical Nomenclature Examples (More Detail in Lab) • Ionic Compounds NaCl sodium chloride Al2S3 aluminum sulfide FeSO4 iron(II) sulfate KClO3 potassium chlorate • Covalent Compounds SO2 sulfur dioxide P2O5 diphosphorus pentaoxide N2O dinitrogen oxide

  24. The Mole - The mole is the chemist’s counting unit pair = 2 Dozen = 12 Gross = 144 Ream = 500 Avogadro’s Number (NA) = 6.022 X 1023

  25. By definition, 12C = 12.000 amu How many particles does it take to have 12.000 grams of 12C ? NA = 6.022 X 1023 (as determined by experiment) Where Does Avogadro’s Number Come From?

  26. Significance of the Mole Mass in amu’s Mass in grams/mole NA of carbon atoms weighs NA of iron atoms weighs

  27. Molar Mass - the mass in grams of one mole of any element • Molar mass of sodium (Na) = mass of 1 mol of Na atoms = 22.99 g/mol = mass of 6.022 X 1023 Na atoms • Molar mass of lead (Pb) = mass of 1 mol of Pb atoms = 207.2 g/mol = mass of 6.022 X 1023 Pb atoms

  28. Moles to Mass Mass to Moles moles • grams = grams 1 mole grams • 1 mole = moles grams Molar mass 1 / Molar mass Mass Moles Conversion

  29. Example 3.6 - Mass to Moles How many moles are represented by 125 g of silicon, an element used in semiconductors?

  30. Example 3.7 - Moles to Mass What mass, in grams, is equivalent to 2.50 mol of lead (Pb) ?

  31. Mole Calculation Using Density The graduated cylinder in the photograph contains 25.0 cm3 of Hg. If the density of Hg = 13.534 g/cm3 at 25 oC, how many moles of Hg are in the cylinder? How many atoms of Hg are there?

  32. Molar Mass of a Compound Sum up the molar masses of each atom in the compound HC2H3O2

  33. Example 3.9 - Molar Mass & Moles You have 16.5 g of the common compound oxalic acid, H2C2O4. Calculate - 1. The number of moles2. The number of molecules3. The number of C atoms4. The mass of one molecule

  34. C2H5OHMW = 46.07 Other Fun Stuff 46.07 g contains - 2(12.01) = 24.02 g of carbon 1(16.00) = 16.00 g of oxygen 6(1.008) = 6.05 g of hydrogen 1 molecule contains - 2 carbon atoms 1 oxygen atom 6 hydrogen atoms 1 mole contains - 2 moles of carbon atoms 1 mole of oxygen atoms 6 moles hydrogen atoms

  35. Conversion factors for C2H5OH - • 2(12.01) g C/ 46.07 g C2H5OH OR24.02 g C/ 46.07 g C2H5OH • 6 moles H/ mole C2H5OH • 1 mole oxygen/ 2 moles C

  36. More Problems - How many grams of Na are there in 200. g of Na2CO3 ? How many moles of oxygen are there in 25.0 mol of SO2 ?

  37. More Problems - How many aluminum atoms are there in 150. g of Al2O3 ? How many oxygen atoms are there in 500. mL of a 30.0 % solution of H2SO4 with a density of 1.250 g/cm3 ? (MW = 98.1)

  38. Percent Composition from a Known Formula NH3 MW= 17.03 g/mol % N = % H =

  39. Empirical & Molecular Formulas Empirical = simplest ratio of atoms in the molecules Molecular = actual ratio

  40. Calculating Empirical Formulas Formulas of unknown compounds are determined from the percent composition of each element by mass. Assume 100 g and divide by atomic weight Divide by fewest number of moles

  41. Calculating Molecular Formulas The molecular weight must be known. It is obtained from a separate experiment Benzene empirical formula = CHformula weight = 12.01 + 1.008 = 13.018 If the MW = 78.11, then what is the molecular formula?

  42. Example 3.10 Eugenol is the active component of oil of cloves. It has a MW of 164.2 g/mol and is 73.14 %C and 7.37 %H; the remainder is oxygen. What are the empirical and molecular formulas?

  43. Another example - Vanillin is a common flavoring agent. It has a molar mass of 152 g/mol and is 63.15 %C and 5.30 %H; the rest is oxygen. What are the empirical and molecular formulas?

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