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Chemistry 4362

Chemistry 4362. Advanced Inorganic Chemistry. Instructor: Dr. Byron K. Christmas Class Time: Tue & Thur - 5:29 to 6:49 p.m. Classroom: N-936  C-320 Phone: (713) 221-8169 FAX: (713) 221-8528 E-Mail: ChristmasB@uhd.edu. Introduction.

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Chemistry 4362

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  1. Chemistry 4362 Advanced Inorganic Chemistry Instructor: Dr. Byron K. Christmas Class Time: Tue & Thur - 5:29 to 6:49 p.m. Classroom: N-936  C-320 Phone: (713) 221-8169 FAX: (713) 221-8528 E-Mail: ChristmasB@uhd.edu

  2. Introduction • “The chemistry of everything that is NOT organic…” What is Inorganic Chemistry ? • “The chemistry of all of the elements and their compounds • except for the hydrocarbons and their derivatives.” • “The branch of chemistry falling between and overlapping • with physical chemistry and organic chemistry.” • “What Inorganic Chemists Do!” • Your Personal Definition??

  3. What Do Inorganic Chemists Do ? • Synthesize and characterize substances other than those that are clearly “organic”. • Determine the structures of inorganic substances. • Investigate the chemical reactions of inorganic substances. • Investigate the physical properties of inorganic substances. . • Develop hypotheses and theories to explain and systematize • the empirical data collected.

  4. Why Should You Study Inorganic Chemistry ? Elemental Composition of the Sun and the Universe Sun Universe Hydrogen92.5 % 90.87 % Helium7.3 % 9.08 % All Others 0.2 % 0.05 % • Essentially the entire universe is Inorganic. • The Earth is predominantly Inorganic. Elemental Composition of the Earth’s Crust Oxygen 45.5 % Iron 6.20 % Silicon 27.2 % Calcium 4.66 % Aluminum 8.30 % All Others 8.14 %

  5. Inorganic materials are an essential part of our national economy. U.S. Production of Top 10 Chemicals (x 109 lb.) - 1997 Sulfuric Acid 95.58 Nitrogen 82.88* Oxygen 64.84* Ethylene 51.08 Lime 42.56 Ammonia 38.39 Propylene 27.53 Phosphoric Acid 26.83 Ethylene Dichloride 26.29 Sulfur 26.24 From C&EN, June 29, 1998 *Calculated from “billion cubic feet at STP”

  6. http://scifun.chem.wisc.edu/chemweek/Sulf&top/Sulf&Top.html

  7. U.S. Production of Top 50 Chemicals (x 109 lb.) - 1994 Total Organics 279.17 Total Inorganics 450.19 Grand Total 729.36 • Inorganics are essential to life. • Water is essential for all life. • About 30 different elements are believed to be essential to life - 28 in addition to carbon and hydrogen. • For all practical purposes, Inorganic Chemistry IS • chemistry - the study of the properties, composition, and structure of matter, the physical and chemical changes it undergoes, and the energy liberated or absorbed during those changes.

  8. Approaches to the Study of Inorganic Chemistry • Empirical Approach (Descriptive Chemistry) • Historically this was the way it was taught. • It involves essentially all memorization. • It is necessary for a complete understanding of Chemistry. • Theoretical Approach • It provides a framework for understanding the “why” of descriptive chemistry. • It can provide “intellectual satisfaction”. • It is limited in its ability to give explanations for all observed phenomena. • It has dominated the teaching of Inorganic Chemistry for over 30 years.

  9. Introduction to Descriptive Chemistry Definition: “…the study of the composition, structure, and properties of matter….” - thefacts about the elements and their compounds. Comments from the Experts: “…evident differences between this and previous editions…is the absence of much theoretical material previously included…the continuing rapid growth of chemistry…required the addition of impor- tant new factsto all of the descriptive material…over the years, become less persuaded of the value of certain types of theorizing….Thus, we felt obliged to make space forfacts at the expense of theoretical material.” Cotton and Wilkinson, Advanced Inorganic Chemistry, 5th Edition, 1988.

  10. Comments from the Experts: “The facts concerning the properties and reactions of substances are the very essence of chemistry. Facts undergo little if any change in contrast to constantly changing theories. Moreover, …a chemist needs a solid background of facts in order to appreciate the need for theories….” R. J. Gillespie in the Forward of Chemistry of the Elements, Greenwood and Earnshaw, 1st Edition, 1984.

  11. “Over the years, the theoretical part tended to grow at the expense of the descriptive material….The theoretical part tended to become the end rather than the means….By the 1970’s many teachers had to abandon any attempt to cover descriptive inorganic chemistry in the traditional sense. Thus we can encounter the student who can write an erudite account of structural minutiae in copper(II) chemistry, ligand field spectra and…,but who knows little about the more mundane compounds of the transition elements and would be hard pressed to locate indium in the Periodic Table, let alone venture anything about its chemistry.” Derek W. Smith, Inorganic Substances: A Prelude to the Study of Descriptive Inorganic Chemistry, 1st Edition, 1990.

  12. “Chemistry has always been, and still is, a practical subject….An American professor told me he divided inorganic chemistry books into two types: theoretical and practical. In deciding how to classify any particular book, he first looked to see if the extraction of the two most produced metals (Fe and Al) was adequately covered, what impurities were likely to be present, and how the processing was adapted to re- move them. Second, he looked to see if the treatment of the bonding in xenon compounds and ferrocene was longer than that of the pro- duction of ammonia. Third, he looked to see if the production and uses of phosphates were covered adequately….For some years there has been a trend for chemistry teaching to become more theoretical. There is always theoretical interest in another interesting oxidation state or another unusual complex, but the balance of this book is tilted to ensure that they do not exclude the commonplace, the mun- dane and the commercially important.”J. D. Lee, Concise Inorganic Chemistry, 5th Edition,1996.

  13. Industrial Applications Approach • Few schools other than chemical engineering programs have used this approach. • It is of great “practical” importance for students preparing for industry. • It is of limited utility in preparing for graduate work in chemistry. • Balanced Approach • Provides a balance among all approaches. • Applicable to “survey-type” course. • Useful for either graduate school or industry preparation. • Used in THIS COURSE!!

  14. Course Overview • Introduction to Inorganic Polymers • Theoretical Concepts • Atomic Structure & the Periodic Table • Properties of the Elements • Introduction to Chemical Bonding • The Covalent Bond • The Metallic Bond • The Ionic Bond • Intermolecular Attractive Forces • Inorganic Thermodynamics and Kinetics • Solvent Systems and Acids and Bases • Oxidation/Reduction • Descriptive Chemistry and Industrial Applications • Student Presentations

  15. Introduction to Inorganic Polymers (Page 338 in Text) Questions to Ponder: 1. Would you know an Inorganic Polymer if you saw one? 2. How could you determine if an inorganic material was, in fact, “polymeric”? 3. List important types of Inorganic Polymers. 4. How would you determine what is and what is NOT an Inorganic Polymer? Is NaCl a Polymer? Is Graphite a Polymer? What about Diamond? Is Aluminum a Polymer? What about Window Glass? 5. What general principles of chemical bonding, atomic size, etc. lead to effective polymer formation for different types of elements? 6. What are commercially important inorganic polymers?

  16. Introduction to Inorganic Polymers Catenation – What are the requirements? Valence of two or more? Bond energies? What else? Homocatenation Heterocatenation Bond Energies – kJ/mole C-C 356 Si-Si 222 Ge-Ge 188 Sn-Sn 167 Pb-Pb 87 S-S 251 P-P 201 O-O 142 Si-O 460 Sn-O 544 Al-O 586 Si-N 355 B-N 460 http://chemviz.ncsa.uiuc.edu/content/doc-resources-bond.html

  17. Introduction to Inorganic Polymers Polysulfide Demonstration/Experiment Questions? Assignments! Study Hand-outs and your text on Inorganic Polymers Find three to five ADDITIONAL references on the Web and study them Prepare for next Thursday’s Silicone laboratory (Page 176 in Lab Manual)

  18. Theoretical Concepts Chapter 1 Atomic Structure & the Periodic Table Properties of the Elements Introduction to Chemical Bonding The Ionic Bond The Covalent Bond The Metallic Bond Intermolecular Attractive Forces Thermodynamics Acids and Bases Oxidation/Reduction

  19. ATOMIC STRUCTURE Definition of Chemistry: The study of the properties, composition, and STRUCTURE of matter, the physical and chemical changes it undergoes, and the energy liberated or absorbed during those changes. The foundation for theSTRUCTUREof inorganic materials is found in theSTRUCTUREof the atom. Material Properties Bulk Structure Molecular Structure Atomic Structure

  20. ATOMIC STRUCTURE Historical Development: • Greek Concepts of Matter • Aristotle - Matter is continuous, infinitely • divisible, and is composed of only 4 elements: • Earth, Air, Fire, and Water • Won the philosophical/political battle. • Dominated Western Thought for Centuries. • Seemed very “logical”. • Was totally WRONG!!

  21. ATOMIC STRUCTURE The “Atomists” (Democritus, Lucippus, Epicurus, et. al.) - Matter consists ultimately of “indivisible” particles called “atomos” that canNOT be further subdivided or simplified. If these “atoms” had space between them, nothing was in that space - the “void”. • Lost the philosophical/political battle. • Lost to Western Thought until 1417. • Incapable of being tested or verified. • Believed the “four elements” consisted of “transmutable” atoms. • Was a far more accurate, though quite imperfect “picture” of reality.

  22. ATOMIC STRUCTURE Modern Concepts of Matter John Dalton (1803)- An atomist who formalized the idea of the atom into a viable scientific theory in order to explain a large amount of empirical data that could not be explained otherwise. • Matter is composed of small “indivisible”particles called “atoms”. • The atoms of each element are identical to each other in mass but different from the atoms of other elements. • A compound contains atoms of two or more elements bound together in fixed proportions by mass.

  23. ATOMIC STRUCTURE • A chemical reaction involves a rearrangement of of atoms but atoms are not created nor destroyed during such reactions. Present Concepts - An atom is an electrically neutral entity consisting of negatively charged electrons (e-) situated outside of a dense, posi- tively charged nucleus consisting of positively charged protons (p+) and neutral neutrons (n0). ParticleChargeMass Electron - 1 9.109 x 10 -28 g Proton +1 1.673 x 10 -24 g Neutron 0 1.675 x 10 -24 g

  24. ATOMIC STRUCTURE Nucleus Model of a Helium-4 (4He) atom p+no e- e- no p+ Electron Cloud How did we get this concept? - This portion of our program is brought to you by: Democritus, Dalton, Thompson, Planck, Einstein, Millikan, Rutherford, Bohr, de Broglie, Heisenberg, Schrödinger, Chadwick, and many others.

  25. ATOMIC STRUCTURE • Democritus - First atomic ideas • Dalton - 1803 - First Atomic Theory • J. J. Thompson - 1890s - Measured the charge/mass • ratio of the electron (Cathode Rays) Fluorescent Material _ Cathode + Anode Electric Field Source (Off) With the electric field off, the cathode ray is not deflected.

  26. ATOMIC STRUCTURE - Fluorescent Material - Cathode + + Anode Electric Field Source (On) With the electric field on, the cathode ray is deflected away from the negative plate. The stronger the electric field, the greater the amount of deflection. - Cathode + Anode Magnet

  27. ATOMIC STRUCTURE With the magnetic field present, the cathode ray is deflected out of the magnetic field. The stronger the magnetic field, the greater the amount of deflection. e/m = E/H2r e = the charge on the electron m = the mass of the electron E = the electric field strength H = the magnetic field strength r = the radius of curvature of the electron beam Thompson, thus, measured the charge/mass ratio of the electron - 1.759 x 108 C/g

  28. ATOMIC STRUCTURE • Summary of Thompson’s Findings: • Cathode rays had the same properties no matter what metal was being used. • Cathode rays appeared to be a constituent of all matter and, thus, appeared to be a “sub-atomic” particle. • Cathode rays had a negative charge. • Cathode rays have a charge-to-mass ratio of 1.7588 x 108 C/g.

  29. ATOMIC STRUCTURE R. A. Millikan - Measured the charge of the electron. In his famous “oil-drop” experiment, Millikan was able to determine the charge on the electron independently of its mass. Then using Thompson’s charge-to-mass ratio, he was able to calculate the mass of the electron. e = 1.602 10 x 10-19 coulomb e/m = 1.7588 x 108 coulomb/gram m = 9.1091 x 10-28 gram Goldstein - Conducted “positive” ray experiments that lead to the identification of the proton. The charge was found to be identical to that of the electron and the mass was found to be 1.6726 x 10-24 g.

  30. ATOMIC STRUCTURE Ernest Rutherford - Developed the “nuclear” model of the atom. The Plum Pudding Model of the atom: + + + A smeared out “pudding” of positive charge with negative electron “plums” imbedded in it. - - - - - - - - - - - - - - - + + + + + + + The Metal Foil Experiments: Fluorescent Screen a-particles Radioactive Material in Pb box. Metal Foil

  31. ATOMIC STRUCTURE If the plum pudding model is correct, then all of the massive a-particles should pass right through without being deflected. In fact, most of the a - particles DID pass right through. However, a few of them were deflected at high angles, disproving the “plum pudding” model. Rutherford concluded from this that the atom con- sisted of a very dense nucleus containing all of the positive charge and most of the mass surrounded by electrons that orbited around the nucleus much as the planets orbit around the sun.

  32. ATOMIC STRUCTURE Assignment: Assume the diameter of the nucleus of a hydrogen atom is 1 x 10 -13 cm and the diameter of the atom is 1 x 10 -8 cm. 1. Calculate the volume of the nucleus and the volume of the atom in cm3 . 2. Calculate the volume of empty space in the atom. 3. Calculate the ratio of the volume of the nucleus to volume of the whole atom. 4. Calculate the density of the nucleus if the proton’s mass is1.6726 x 10-24 g

  33. Problems with the Rutherford Model: It was known from experiment and electromagnetic theory that when charges are accelerated, they continuously emit radiation, i.e., they loose energy continuously. The “orbiting” electrons in the atom were, obviously, not doing this. ATOMIC STRUCTURE Planck • Atomic spectra and blackbody radiation were known to be DIScontinuous. Bohr • The atoms were NOT collapsing.

  34. ATOMIC STRUCTURE Atomic Spectra - Since the 19th century, it had been known that when elements and compounds are heated until they emit light (glow) they emit that light only at discrete frequencies, giving a line spectrum. - + Hydrogen Gas Line Spectrum

  35. ATOMIC STRUCTURE When white light is passed through a sample of the vapor of a substance, only discrete frequencies are absorbed, giving an absorption ban spectrum. These frequencies are identical to those of the line spectrum of the same element or compound. For hydrogen, the spectroscopists of the 19th Century found that the lines were related by the Rydberg equation: n/c = R[(1/m2) - (1/n2)] n = frequency R = Rydberg Constant c = speed of light m = 1, 2, 3, …. n = (m+1), (m+2), (m+3), ….

  36. ATOMIC STRUCTURE Max Planck - In 1900 he was investigating the nature of black body radiation and tried to interpret his findings using accepted theories of electromagnetic radiation (light). He was NOT successful since these theories were based on the assumption that light had WAVE characteristics. To solve the problem he postulated that light was emitted from black bodies in discrete packets he called “quanta”. Einstein later called them “photons”. By assuming that the atoms of the black body emitted energy only at discrete frequencies, he was able to explain black body radiation. E = hn

  37. Both spectroscopy and black body radiation indicated that atoms emitted energy only at discrete frequencies or energies rather than continuously. ATOMIC STRUCTURE Is light a particle or a wave?? Why do atoms emit only discrete energies? What actually happens when light interacts with matter? What was wrong with Rutherford’s Model?

  38. Niels Bohr - Bohr corrected Rutherford’s model • of the atom by formulating the following postulates: • Electrons in atoms move only in discrete orbits around the nucleus. • When in an orbit, the electron does NOT emit energy. • They may move from one orbit to another but are NEVER residing in between orbits. • When an electron moves from one orbit to another, it absorbs or emits a photon of light with a specific energy that depends on the difference in energy between the two orbits. ATOMIC STRUCTURE

  39. ATOMIC STRUCTURE Balmer Series (Visible) Lyman Series + Paschen Series (UV) (IR) The Bohr Model of the Atom

  40. ATOMIC STRUCTURE • The lowest possible energy state for an electron is called the GROUND STATE. All other states are called EXCITED STATES. En = (- 2.179 x 10-18 J)/n2 Ephoton = Ehigh - Elow Ephoton = [(- 2.179 x 10-18 J)/n2high] -[(- 2.179 x 10-18 J)/n2low] = - 2.179 x 10-18 J[(1/n2high) - (1/n2low)] Does this equation look familiar? n/c = R[(1/m2) - (1/n2)]

  41. ATOMIC STRUCTURE Niels Bohr won the Nobel Prize for his work. However, the model only worked perfectly for hydrogen. What about all of those other elements?? Louis de Broglie - Thought that if light, which was thought to have wave characteristics, could also have particle characteristics, then perhaps electrons, which were thought to be particles, could have characteristics of waves. l = h/mv An electron in an atom was a “standing wave”!

  42. ATOMIC STRUCTURE Werner Heisenberg - Developed the “uncertainty” principle: It is impossible to make simultaneous and exact measurements of both the position (location) and the momentum of a sub-atomic particle such as an electron. (Dx)(Dp)  h/2p Our knowledge of the inner workings of atoms and molecules must be based on probabilities rather than on absolute certainties.

  43. ATOMIC STRUCTURE Erwin Schrödinger - Developed a form of quantum mechanics known as “wave mechanics”. Hy = Ey H = Hamiltonian operator E = Total energy of the system y = Wave function [(-h2)/(8p2m)]2 - [e2/r]  = E Kinetic Energy Term Potential Energy Term This is simply a quantum mechanical statement of the Law of Conservation of Energy

  44. Of the numerous solutions to the Schrödinger equation • for hydrogen, only certain ones are allowed due to the • following boundary conditions: • Y, the wave function, must be continuous and finite. • It must be single-valued at all points (There can’t be two different probabilities of finding an electron at one point in space). • The probability of finding the electron, Y2, somewhere in space must = 1. ATOMIC STRUCTURE +  Y2dxdydz = 1 -  Y has many values that meet these conditions. They are called “orbitals”.

  45. ATOMIC STRUCTURE • Wave Function - A mathematical function associated • with each possible state of an electron in an atom or • molecule. • It can be used to calculate the energy of an • electron in the state • the average and most probable distance from the nucleus • the probability of finding the electron in any specified region of space. Y Y Y Y Y

  46. ATOMIC STRUCTURE Quantum Numbers: Principle Quantum Number, n - An integer greater than zero that represents the principle energy level or “shell” that an electron occupies. Energy # of orbitals n Level Shell n2 1 1st K 1 2 2nd L 2 3 3rd M 9 4 4th N 16 etc. etc. etc. etc.

  47. ATOMIC STRUCTURE Azimuthal Quantum Number, l - The quantum number that designates the “subshell” an electron occupies. It is an indicator of the shape of an orbital in the subshell. It has integer values from 0 to n-1. l = 0, 1, 2, 3, …, n - 1 s p d f Magnetic Quantum Number, ml - The quantum number that determines the behavior of an electron in a magnetic field. It designates the orbitaland has integer values from -l to +l including 0. ml = -l, …, -3, -2, -1, 0, +1, +2, +3, …, +l

  48. Orbital # of n l Name ml Orbitals 1 0 1s 0 1 2 0 2s 0 1 1 2p -1, 0, +1 3 3 0 3s 0 1 1 3p -1, 0, +1 3 2 3d -2, -1, 0, +1, +2 5 etc. etc. etc. etc. etc. ATOMIC STRUCTURE Spin Quantum Number, ms - The quantum number that designates the orientation of an electron in a magnetic field. It has half-integer values, +½ or -½.

  49. ATOMIC STRUCTURE So what do atoms look like? A. Interpretation of Y: Theprobability of finding an electron in a small volume of space centered around some point is proportional to the value of Y2at that point. B. Electron Probability Density vs. r C. Dot Density Representation: Imagine super- imposing millions of photographs taken of an electron in rapid succession. D. Radial Densities

  50. ATOMIC STRUCTURE • Electron Configuration • A. Many-electron atom: An atom that contains • two or more electrons. • B. Problems with the Bohr model: • 1. It “assumed” quantization of the energy • levels in hydrogen. • 2. It failed to describe or predict the spectra • of more complicated atoms.

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