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Chemistry NYA Class Notes and Exercises Part 2

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College Chemistry NYA Class Notes and Exercises Part 2
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  Email:  jrobichaud@champlaincollege.qc.caOffice:f-228Phone:x311 CHAMPLAIN COLLEGESAINT-LAMBERT GENERAL CHEMISTRY  202-NYA-205  Winter 2012  JoelRobichaud  Course Outline ã nomenclature ã empirical & molecular formulas ã stoichiometry  ã gas laws ã molarity  UNIT 1:BasicsUNIT 2: Atomic Theory ã history of atomic theory  ã the Bohr atom ã the modern approach (quantum theory) ã quantum numbers ã electron configurations ã electron affinity  UNIT 3: Periodicity & Chemical Reactions ã electron configuration & chemical properties of elements ã ionization energy  ã atomic and ionic size ã electronegativity& electron affinity  ã reactions of the main group elements ã  writing molecular & net ionic equations UNIT 4:Chemical Bonding ã analysis of ionic & covalent bonding ã  writing Lewis structures, resonance structures formal charges ã shapes of molecules bond angles ã bond polarity ã dipole moments ã hybridization theory orbital diagrams UNIT 5:Intermolecular Forces ã intermolecular forces intramolecularbonds ã dispersion forces, dipole-dipole forces ã hydrogen bonding ã relationship of melting & boiling point & solubility to intermolecular forces ã classification of substances UNIT 6:Liquids & Solutions ã properties of solids ã phase changes & phase diagrams ã physical properties of solutions ã concentration units ã colligativeproperties  ã give a brief overview of Thomson’s experiments which led to the discovery of the electron. (2.2) ã give a brief overview of Rutherford’s experiments which led to the discovery of the nucleus. (2.2) ã describe the mass & charge of each of the 3 fundamental subatomic particles: proton, electron & neutron. (2.2) ã explain the terms: atomic number, mass number, molar mass, atomic mass unit, isotope and atomic mass. (2.3, 3.2) ã describe the main features of the Rutherford model of the hydrogen atom. (class notes) ã define wavelength and frequency, and solve problems using the relationship . (7.1) ã identify the regions of the electromagnetic spectrum based on their frequencies and wavelengths. (7.1) ã explain the difference between continuous and line spectra. (7.3) ã describe the experimental setup to obtain the atomic spectrum of hydrogen. (class notes, emission experiment) ã give a brief outline of Planck’s quantum theory and solve problems based on Planck’s equation . (7.1) ã describe the Bohr model of the hydrogen atom. (7.3) ã use the Bohr equation to calculate the energy of the electron in a given Bohr orbit. (7.3, emission experiment) calculate ∆E for the transition of an electron from one Bohr energy level to another. Also calculate and associated  with these transitions. (7.3, emission experiment) ã recognize how calculations of E, and for energy transitions in the hydrogen atom enabled Bohr to explain the observed line spectrum. Also explain the importance of the Bohr model in the eventual development of atomic theory. (7.3, emission experiment) ã explain why, based on the Bohr model, hydrogen gives a line, not a continuous spectrum. (class notes, emission experiment) ã explain the shortcomings of the Rutherford model of the hydrogen atom. (class notes) ã explain the fundamental differences between the Bohr and Rutherford models of the hydrogen atom. (class notes) ã recognize the difference between emission and absorption spectra. (class notes) Unit II: AtomicTheory  (Chang, Ch. 2, 7 & 8)Objectives:  ã recognize the difference between the ground state & excited states in the hydrogen atom spectrum. (class notes, emission experiment) ã calculate the ionization energy of the hydrogen atom (7.3, emission experiment) ã recognize that the Bohr model was only successful for one electron species, and explain the general limitations of the Bohr model. (7.3) ã outline the general features of the quantum mechanical model of the hydrogen atom based on the idea of the  wavelike properties of matter. (7.5) ã outline the contribution of each of the following scientists to the development of the quantum mechanical model of the atom: Bohr (7.5), Heisenberg (7.5), Schrödinger (7.5), De Broglie (7.4), Pauli(7.8). ã list the quantum numbers: n , l  , m l  obtained by solving the Schrödinger wave equation for the electron in the hydrogen atom. Predict the allowable values of these quantum numbers. (7.6) ã explain the meaning of “wave function” and “orbital”. (7.5) ã relate the shapes of orbitalsto the quantum number l  . (7.6, 7.7) ã sketch the shapes of s and  p orbitals. (7.7) ã recognize the shape s of d  orbitals. (7.7) ã explain the meaning of electron spin and its relationship to the quantum number m s (7.6) ã apply the Pauli exclusion principle and Hund’s rule to write the electron configuration and orbital filling (box) diagrams for multielectronatoms. (7.8) ã explain and apply the concept of “effective nuclear charge”. (7.9 & 8.3) ã identify the s ,  p , d  , and  f  blocks in the periodic table. (7.9) ã relate the position of an element in the periodic table to its electron configuration. (7.9) ã know the anomalous electron configurations for first and second row transition metals. (7.9) ã explain and predict the trends in atomic radii. (8.3) ã predict whether an atom is paraor diamagnetic based on its electron configuration. (7.8) ã give the allowable sets of the four quantum numbers (n, l, m l   , m s ) for each electron in a given atom. (7.8) Continuation…

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