1 / 33

Lecture 1

Lecture 1. Brooklyn College Inorganic Chemistry (Spring 2006). Prof. James M. Howell Room 359NE (718) 951 5458; jhowell@brooklyn.cuny.edu Office hours : Mon. & Thu. 10:00 am-10:50 am & Wed. 5 pm-6 pm Textbook: Inorganic Chemistry, Miessler & Tarr,

hope-munoz
Download Presentation

Lecture 1

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Lecture 1

  2. Brooklyn CollegeInorganic Chemistry(Spring 2006) • Prof. James M. Howell • Room 359NE (718) 951 5458; jhowell@brooklyn.cuny.edu Office hours: Mon. & Thu. 10:00 am-10:50 am & Wed. 5 pm-6 pm • Textbook: Inorganic Chemistry, Miessler & Tarr, 3rd. Ed., Pearson-Prentice Hall (2004)

  3. What is inorganic chemistry? Organic chemistry is: the chemistry of life the chemistry of hydrocarbon compounds C, H, N, O Inorganic chemistry is: The chemistry of everything else The chemistry of the whole periodic Table (including carbon)

  4. Single and multiple bonds in organic and inorganic compounds

  5. Unusual coordination numbers for H, C

  6. Typical geometries of inorganic compounds

  7. Inorganic chemistry has always been relevant in human history • Ancient gold, silver and copper objects, ceramics, glasses (3,000-1,500 BC) • Alchemy (attempts to “transmute” base metals into gold led to many discoveries) • Common acids (HCl, HNO3, H2SO4) were known by the 17th century • By the end of the 19th Century the Periodic Table was proposed and the early atomic theories were laid out • Coordination chemistry began to be developed at the beginning of the 20th century • Great expansion during World War II and immediately after • Crystal field and ligand field theories developed in the 1950’s • Organometallic compounds are discovered and defined in the mid-1950’s (ferrocene) • Ti-based polymerization catalysts are discovered in 1955, opening the “plastic era” • Bio-inorganic chemistry is recognized as a major component of life

  8. Nano-technology

  9. Hemoglobin

  10. The hole in the ozone layer (O3) as seen in the Antarctica http://www.atm.ch.cam.ac.uk/tour/

  11. Some examples of current important uses of inorganic compounds Catalysts: oxides, sulfides, zeolites, metal complexes, metal particles and colloids Semiconductors: Si, Ge, GaAs, InP Polymers: silicones, (SiR2)n, polyphosphazenes, organometallic catalysts for polyolefins Superconductors: NbN, YBa2Cu3O7-x, Bi2Sr2CaCu2Oz Magnetic Materials: Fe, SmCo5, Nd2Fe14B Lubricants: graphite, MoS2 Nano-structured materials: nanoclusters, nanowires and nanotubes Fertilizers: NH4NO3, (NH4)2SO4 Paints: TiO2 Disinfectants/oxidants: Cl2, Br2, I2, MnO4- Water treatment: Ca(OH)2, Al2(SO4)3 Industrial chemicals: H2SO4, NaOH, CO2 Organic synthesis and pharmaceuticals: catalysts, Pt anti-cancer drugs Biology: Vitamin B12 coenzyme, hemoglobin, Fe-S proteins, chlorophyll (Mg)

  12. Atomic structure A revision of basic concepts

  13. Energy levels in the hydrogen atom Energy of transitions in the hydrogen atom Atomic spectra of the hydrogen atom Paschen series (IR) Balmer series (vis) Bohr’s theory of circular orbits fine for H but fails for larger atoms …elliptical orbits eventually also failed0 Lyman series (UV)

  14. Planck quantization of energy h = Planck’s constant n = frequency E = hn • = wavelength h = Planck’s constant m = mass of particle v = velocity of particle de Broglie wave-particle duality l = h/mv Heisenberg uncertainty principle Dx uncertainty in position Dpx uncertainty in momentum Dx Dpx h/4p • H: Hamiltonian operator • : wave function E : Energy Schrödinger wave functions The fundamentals of quantum mechanics

  15. Quantum mechanics provides explanations for many experimental observations From precise orbits to orbitals: mathematical functions describing the probable location and characteristics of electrons electron density: probability of finding the electron in a particular portion of space

  16. Characteristics of a well behaved wave function • Single valued at a particular point (x, y, z). • Continuous, no sudden jumps. • Normalizable. Given that the square of the absolute value of the eave function represents the probability of finding the electron then sum of probabilities over all space is unity. It is these requirements that introduce quantization.

  17. Electron in One Dimensional Box Definition of the Potential, V(x) V(x) = 0 inside the box 0 <x<l V(x) = infinite outside box; x <0 or x> l

  18. Q.M. solution in atomic units • ½ d2/dx2 X(x) = E X(x) Standard technique: assume a form of the solution. Assume X(x) = a ekx Where both a and k will be determined from auxiliary conditions. Recipe: substitute into the DE and see what you get.

  19. Substitution yields ½ k2 ekx = E ekx or k = +/- i (2E)0.5 General solution becomes X (x) = a ei sqrt(2E)x + b e –i sqrt(2E)x where a and b are arbitrary consants Using the Cauchy equality eiz = cos(z) + i sin(z) Substsitution yields X(x) = a cos (sqrt(2E)x) + b (cos(-sqrt(2E)x) + i a sin (sqrt(2E)x) + i b(sin(-sqrt(2E)x)

  20. Regrouping X(x) = (a + b) cos (sqrt(2E)x) + i (a - b) sin(sqrt(2E)x) Or X(x) = c cos (sqrt(2E)x) + d sin(sqrt(2E)x) We can verify the solution as follows ½ d2/dx2X(x) = E X(x) (??) - ½ d2/dx2 (c cos (sqrt(2E)x) + d sin (sqrt(2E)x) ) = - ½ ((2E)(- c cos (sqrt(2E)x) – d sin (sqrt(2E)x) = E (c cos (sqrt(2E)x + d sin(sqrt(2E)x)) = E X(x)

  21. We have simply solved the DE; no quantum effects have been introduced. Introduction of constraints: -Wave function must be continuous at x = 0 or x = l X(x) must equal 0 Thus c = 0, since cos (0) = 1 and second constraint requires that sin(sqrt(2E) l ) = 0 Which is achieved by (sqrt(2E) l ) = n p Or

  22. In normalized form

  23. Atomic problem, even for only one electron, is much more complex. • Three dimensions, polar spherical coordinates: r, q, f • Non-zero potential • Attraction to nucleus • For more than one electron, electron-electron repulsion. The solution of Schrödinger’s equations for a one electron atom in 3D produces 3 quantum numbers Relativistic corrections define a fourth quantum number

  24. Symbol Name Values Role n Principal 1, 2, 3, ... Determines most of the energy l 0 1 2 3 4 5 l Angular momentum 0, 1, 2, ..., n-1 Describes the angular dependence (shape) and contributes to the energy for multi-electron atoms orbital s p d f g ... ml Magnetic 0, ± 1, ± 2,..., ± l Describes the orientation in space ms Spin ± 1/2 Describes the orientation of the spin of the electron in space Quantum numbers Orbitals are named according to the l value:

  25. Principal quantum number n = 1, 2, 3, 4 …. determines the energy of the electron in a one electron atom indicates approximately the orbital’s effective volume n = 1 2 3

  26. s Angular momentum quantum number l = 0, 1, 2, 3, 4, …, (n-1) s, p, d, f, g, ….. determines the shape of the orbital

  27. Magnetic quantum number • Determines the spatial orientation of the orbital ml = -l,…, 0 , …, +l l = 2 ml = -2, -1, 0, +1, +2 l = 0 ml = 0 l = 1; ml = -1, 0, +1 See: http://www.orbital.com

  28. Electrons in polyelectronic atoms (the Aufbau principle) • Electrons are placed in orbitals to give the minimum possible energy to the atom • Orbitals are filled from lowest energy up • Each electron has a different set of quantum numbers (Pauli’s exclusion principle) • Since ms =  1/2, no more than 2 electrons may be accommodated in one orbital • Electrons are placed in orbitals to give the maximum possible total spin (Hund’s Rule) • Electrons within a subshell prefer to be unpaired in different orbitals, if possible

  29. Placing electrons in orbitals

  30. 5p E Approximate order of filling orbitals with electrons 4d 5s 3d 4s 4p 3p 3s 2p 2s 1s

  31. 5p E 4d 5s 3d 4s 4p 3p 3s 2p 2s 1s

More Related