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309 Havemeyer

Have you seen this hall in the movies??? But before we start the show…. 309 Havemeyer. G1403 Fall 2004. Thank You!. Please Turn Off Cellular Phones. Tentative Course Material To Be Covered for Exam 1 Exam 1: Wednesday, September 29 (6 lectures covered).

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309 Havemeyer

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  1. Have you seen this hall in the movies??? But before we start the show… 309 Havemeyer

  2. G1403 Fall 2004 Thank You! Please Turn Off Cellular Phones

  3. Tentative Course Material To Be Covered for Exam 1 Exam 1: Wednesday, September 29 (6 lectures covered). Chapter 1. The Atomic Nature of Matter. Chapter 2. Stoichiometry. Chapter 3. Chemical Periodicity and the Formation of Simple Compounds. Exam 2: Wednesday, October 27 (7 lectures covered) Chapter 15. Nuclear Chemistry. Chapter 16. Quantum Mechanics and the Hydrogen Atom. Chapter 17. Many-electron Atoms and Chemical Bonding. Exam 3: Wednesday, December 1 (8 lectures covered) Chapter 18. Molecular Orbitals, Spectroscopy, and Atmospheric Chemistry. Chapter 19. Coordination Complexes. Chapter 24. From Petroleum to Pharmaceuticals. Period before final (3 lectures covered) Chapter 25. Synthetic and Biological Polymers (plus spectroscopy of organic molecules).. Final exam December 20, 2004, 1:10 PM-4 PM (probably in 309 Havemeyer)

  4. Chapter 1: Learning goals: (1) Distinguish between elements, compounds and mixtures. (2) Understand the atomic interpretation of elements, compounds and mixtures. (3) Laws leading to the Dalton’s atomic theory of matter. (4) The principles of Dalton’s atomic theory of matter. (5) Law of combining volumes of gases. (6) Avogadro’s hypothesis. (7) The mole concept. (8) Distinguish between empirical formula, molecular formula, structural formula. (9) Perform calculations employing the mole (the chemical amount) concept. (10) Interconvert units using dimensional analysis (recitation section).

  5. Chapter 1. The Atomic Nature of Matter Reading. This chapter is a review of high school chemistry, especially Sections 1-1 through 1-6. The exam will stress Sections 1-7 and 1-8. Homework assignment: Section 1-7: 37, 41, 43 a. Section 1-8: 49, 55, 61, 67. Additional problems: 89, 91, 101, 103, 107. Chapter 2. Stoichiometry Reading. You are expected to read all sections. Homework assignment: Section 2-1: 1, 3 b, c, f, h, 5, 7. Section 2-2: 11, 15, 19, 25. Section 2-3: 33, 41. Section 2-4: 45, 47. Additional problems: 75, 77, 81, 83. Chapter 3. Chemical Periodicity. Formation of Simple Compounds Reading. You are expected to read all sections. Homework assignment:Section 3-3: 17, 19. Section 3-4: 25, 27, 29, 31. Section 3-5: 33, 35, 41,43, 47. Section 3-7: 57, 59, 61, 63, 65, 67, 69. Section 3-8: 73, 81. Additional problems: 87, 89, 91, 93, 95, 97, 105.

  6. Section 1.7: The Mole Concept: Counting and Weighting Atoms and Molecules (1) The mole is a number: 6.02 x 1023. (2) In chemistry the mole refers to the number of atoms or molecules in a substance. (3) The “chemical amount” of a substance is given in moles. The amount of a substance is its mass. (4) One mole of a substance equals the amount of the substance that contains Avogadro’s number of atoms for an element or Avogadro’s number of molecules for a compound. (5) Instead of counting atoms or molecules, chemists weight substances and then compute the number of atoms or molecules. (6) By definition the mass of one mole of carbon atoms is defined as weighing 12 grams exactly. (7) The atomic weight of all other elements are based on carbon.

  7. Catalogs of Materials • Heterogeneous at the level of an optical microscope: wood, most rocks, blood, milk. • Homogeneous mixture: solutions, mixed gasses, many glasses. Looks uniform under a microscope, but can be separated into various components using chromatography, crystallization or distillation.

  8. “Pay attention, therefore, while I demonstrate that there exist certain bodies that are absolutely solid and indestructible, namely those atoms which according to our teaching are the seeds or prime units of things from which the whole universe is constructed.” Lucretius, De Rerum Natura, (Book IV)

  9. All States of Matter Are Made of Atoms and Molecules Densities of Molecules: Rare, or Low: Gas Dense but Fluid: Liquid Dense, Well-Packed and Mostly Stationary: Solid

  10. The Atom: Nucleus Surrounded by Electrons 198Au

  11. Section 1.2 The Composition of Matter and Section 1.3: The Atomic Theory of Matter. Matter: Matter is any material that occupies space and has mass. Atomic interpretation: Any material that contains atoms, which occupy space and have mass. Substance (idealization): A substance is a chemically pure sample of matter. This means the sample does not change its characteristics upon further attempts of purification. Element: An element is a substance which cannot be decomposed into simpler substances by chemical processes. Examples: hydrogen, carbon, oxygen. Atomic interpretation: An element is a substance that contains only one kind of atom. Hydrogen (H) atoms, carbon atoms (C), oxygen atoms (O). Compound: A compound is a substance that can be decomposed into simpler substances by chemical processes. Examples: water, carbon dioxide, carbon monoxide. Atomic interpretation: A compound is a substance that contains atoms of two or more chemical elements. H, C, O. Mixture: A sample of matter that is not a pure substance.

  12. Antoine Lavoisier: 1743-1794 He established the Law of Conservation of Mass byburning phosphorus and sulfur in air, and proving that the products weighed more than the original. The weight gained was lost from the air. Repeating the experiments of Priestley, he demonstrated that air is composed of two parts, one of which combines with metals to form calxes. He named this portion oxygen (Greek for acid-former), and the other azote (Greek for no life). He also discovered that the inflammable air of Cavendish which he termed hydrogen (Greek for water-former), combined with oxygen to produce a dew, as Priestley had reported, which appeared to be water. He begins therefore a first systematic list of the elements.

  13. “Babel” of Chemical Nomenclature 1770 • Chemistry Symbols from the 18th century encyclopedia published by philosopher Denis Diderot and mathemetician Jan Le Rond D’Alembert • These symbols describe Alloys, compounds, chemical operations, quantities, glassware, etc.

  14. A Page from Lavoisier’s Book Nomenclature and Symbols Similar to the Modern He proposed chemical reactions could be written as “equations”, similar to algebra He proposed that we keep track of mass conservation… forerunner to the “balanced reaction”, the subject of the next lecture, Monday September 13.

  15. Laws which lead to the atomic theory of matter: (1) Law of conservation of mass during a chemical reaction: Mass is neither created nor destroyed in a chemical change. (2) Law of definite proportions: For a given compound, the proportions by mass of the elements that make up the compound are in a fixed ratio. This ratio does not depend the origin of the compound or its method of preparation. (3) Law of multiple proportions: When two elements form more than one compound, the masses of one of the elements that combines with a fixed mass of the second element are in the ratio of small whole numbers. (Problem 1_19).

  16. John Dalton Dalton AtomicTheory of Matter (1808) Quaker, son of a poor weaver, working instead of school, until he teaches school at age 12. Interested in meteorology and the atmosphere. He described regular patterns in the chemical combining weights of substances, suggesting the existence of discrete particles of characteristic mass and combining properties. See related concepts of law of combining volumes and law of definite proportions. However indirect this intellectual connection, it was apparently correct and put chemistry on a far more sure footing. 2g H + 16g O -> 18g H2O 14g N + 16g O -> 30g NO 3g H + 14g N -> 17g NH3

  17. Dalton’s Atomic Theory of Matter. (1) All matter consists of solid and individual atoms. (2) All atoms of a given element are identical in mass and all other properties. (3) Different elements have different kinds of atoms. The mass of different atoms are different. (4) The properties of atoms are unchanged when they are involved in chemical transformations (Law of conservation of mass = conservation of atoms). (5) Compounds are formed from elements when atoms of unlike elements combine (Law of definite proportions = when a fixed number of different atoms combine their numbers and weights are in definite relative proportions to each other). (6) More than one compound can be formed by the combination of two elements (Law of multiple proportions = combinations of different atoms sometimes occur in small whole numbers).

  18. Section 1.4 Chemical Formulas and Relative Atomic Masses Empirical formula: A chemical formula which displays the relative, not necessarily the absolute, number of atoms in a compound. Examples: CH = C2H2 = C3H3 etc. Only the ratio of different atoms is known in an empirical formula, not the exact number of each. Chemical formula (composition): A chemical formula displays the symbols for the exact numbers and kinds of elements which compose the compound, with numerical subscripts that state the number of atoms of each element in the compound. Examples: water = H2O, ammonia = NH3

  19. Molecular Structure: The key intellectual framework of all of chemistry. Composition: Number and kinds of atoms Constitution: How all the atoms are connected by bonds to one another Configuration: How all the atoms are distributed in space relative to one another in three dimensions

  20. Paradigm: A characteristic set of beliefs and/or preconceptions (theoretical, instrumental, procedural and metaphysical) that is shared by a community of practitioners. In a global sense the paradigm embraces all of the shared commitments of a scientific group. A paradigm is what defines the scientific community. Thomas Kuhn. 1923-1996.

  21. "In a recently published paper, I pointed out that one of the causes of the present regression of chemical research in Germany is the lack of general, and at the same time thorough chemical knowledge; no small number of our professors of chemistry, with great harm to our science, are laboring under this lack. A consequence of this is the spread of the weed of the apparently scholarly and clever, but actually trivial and stupid, natural philosophy, which was displaced fifty years ago by exact science, but which is now brought forth again, out of the store room harboring the errors of the human mind; by pseudoscientists who try to smuggle it, like a fashionably dressed and freshly rouged prostitute, into good society, where it does not belong." H. Kolbe, “A Sign of the Times” J. Prakt. Chem., 15, 474 (1877). 1818-1884

  22. “A Dr. J. H. van't Hoff, of the Veterinary School at Utrecht, has no liking, apparently, for exact chemical investigation. He has considered it more comfortable to mount Pegasus (apparently borrowed from the Veterinary School) and to proclaim in his book how the atoms appear to him to be arranged in space, when he is on the chemical Mt. Parnassus which he has reached by bold flight.” H. Kolbe, “A Sign of the Times J. Prakt. Chem., 15, 474 (1877). J. H. van't Hoff (1852-1911) First Nobel Prize, Chemistry, 1901

  23. The Law of Combining Volumes of Gases: When two gases react, the volumes that combine are in a ratio of small whole numbers. The ratio of the volume of each product, if a gas, is also in the ratio of small whole numbers. Example: 1 Liter of hydrogen + 1 Liter of chlorine = 2 Liters of Hydrogen Chloride 2 Liters of hydrogen + 1 Liter of oxygen = 2 Liters of Water 3 Liters of hydrogen + 1 Liter of nitrogen = 2 Liters of Ammonia Mass is always conserved; but the volume of a gas is not. Avogadro’s Law: Equal volumes of different gases contain the same number of particles. The particles of a gas may be atoms or molecules. One liter of hydrogen = one liter of chlorine = one liter of hydrogen chloride in terms of particles (read molecules)

  24. Learning Goals: Chapter 2 Stoichiometry (1) How to balance chemical equations by inspection. (2) How to translate the coefficients of the atoms involved in balanced chemical equations into moles of elements and compounds and how to translate moles into mass of elements and compounds. (3) For gases, how to translate the coefficients of the atoms involved in balanced chemical equations into moles of elements and compounds and how to translate moles into volumes of elements and compounds which are gases. (4) How to determine the limiting reagent in a reaction from the balanced chemical equation and the available masses of the reagents.

  25. Section 2.1 Balancing chemical equations Chemical equations: An algebraic representation of a chemical reaction. Balanced chemical equation: Number of moles of atoms on each side of the equation are identical ( Law of conservation of atoms in a chemical reaction). Balancing equations by inspection. Balancing algebraically (Section 2.1, pages 55-57) is possible, but not needed for simple equations. Strategy for balancing chemical equations: (a) Start by giving the coefficient 1 to the most complex formula. The one that contains the most different elements). (b) Inspect both sides of the equation for elements that appear in only one formula which the coefficient is unassigned and balance for that element. (c) Repeat balancing elements, until all are balanced. (d) By convention, balance equations have only integer coefficients. Eliminate fractional coefficients by multiplying all the formulae by the smallest integer that eliminates the fraction.

  26. Example: Find whole numbers for the ? which balance atoms. ? NaCl + ? SO2 + ? H2O + ? O2 -> ? Na2SO4 + ? HCl ? NaCl + ? SO2 + ? H2O + ? O2 -> 1 Na2SO4 + HCl 2 NaCl + 1 SO2 + ? H2O + ? O2 -> Na2SO4 + HCl 2 NaCl + SO2 + ? H2O + ? O2 -> Na2SO4+ 2 HCl 2 NaCl + SO2+ 1H2O + ? O2 -> Na2SO4 + 2 HCl 2 NaCl + SO2 + 1 H2O + 1/2 O2 -> Na2SO4 + 2 HCl Balanced Equation (remove fractional coefficients): 4 NaCl + 2 SO2 + 2 H2O + 1 O2 -> 2 Na2SO4 + 4 HCl

  27. (1) Translate coefficients of a balanced chemical equation into moles. Obtain relative number of moles of reactants and products independent of the actual chemical amounts. (2) Translate coefficients into molecules. Obtain relative number of molecules of reactants and products independent of the actual chemical amounts. (3) Mass relationships in chemical reactions. Translate moles into grams (weight) or moles into liters (volume of gas). (4) Volume relationships in chemical reactions. Translate liters into moles.

  28. Section 2.2: Using balanced chemical equations Combustion of a hydrocarbon (compounds that contain C and H atoms only): Hydrocarbon + Oxygen -> Carbon dioxide + Water ? CxHy + ? O2 -> ? CO2 + ? H2O What is the balanced chemical equation for complete combustion of C4H10? ? C4H10 + ? O2 -> ? CO2 + ? H2O Answer (by inspection): 2 C4H10 + 13 O2 -> 8 CO2 + 10 H2O 116.2 g C4H10 + 416.0 g O2-> 352.1 g CO2 + 180.1 g H2O Example of the law of mass balance. 116.2 g + 416.0 g = 352.1 g + 180.1 g = 532.2 g REACTANTS PRODUCTS SAME MASS

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