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Understanding Molecules and Compounds in Chemistry

Learn about molecular models, ionic and covalent compounds, formation of ions, predicting charges, ionic compound formulas, oxidation numbers, chemical nomenclature, Avogadro’s Number, mole calculations, molar mass, conversion of mass to moles, and more.

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Understanding Molecules and Compounds in Chemistry

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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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