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GENERAL CHEMISTRY AND INORGANIC CHEMISTRY Dr. BÙI THỊ BỬU HUÊ College of Natural Science 1 Chapter 1. ATOMIC STRUCTURE AND THE PERIODIC TABLE Chapter 2. CHEMICAL BONDS AND MOLECULAR STRUCTURE Chapter 3. CHEMICAL THERMODYNAMICS Chapter 4. CHEMICAL KINETICS Chapter 5. CHEMICAL EQUILIBRIUM Chapter 6. SOLUTIONS Chapter 7. ACIDS AND BASES Chapter 8. CHEMISTRY OF METALS Chapter 9. CHEMISTRY OF NONMETALS Chapter 10. TRANSITION METALS AND COMPLEXES 2 References 1. Brady and Holum, 1996, Chemistry: the Study of Matter and its Changes, 2th Ed., John Wiley & Sons. Inc. New York. 2. Umland, Jean B., 1993, General Chemistry, West publishing company. 3. Zumdahl, Steven S., 1995, Chemical Principal, 2th Ed. DC. Health & company. Toronto. 4. http://www.chemistry.msu.edu/Courses/ 5. http://antoine.frostburg.edu 6. http://chemed.chem.purdue.edu 7. http://www.chem1.com/chemed/genchem.html 8. http://www.cbu.edu/~mcondren/lectures.htm 9. http://ull.chemistry.uakron.edu/GenChem/index.html Chapter 1. ATOMIC STRUCTURE AND THE PERIODIC TABLE  Understand atomic structure of an atom including its mass number, isotopes and orbitals.  Know how to account for the structure of the periodic table of the elements based on the modern theory of atomic structure.  Understand general trends of several important atomic properties. FUNDAMENTAL PARTICLES An atom is composed of three types of subatomic particles: the proton, neutron, and electron Particle Mass (g) Charge Proton 1.6727 x 10 -24 +1 Neutron 1.6750 x 10 -24 0 Electron 9.110 x 10 -28 -1 Atomic Structure Atoms consist of very small, very dense positively charged nuclei surrounded by clouds of electrons at relatively great distances from the nuclei. 6 Nuclide Symbol Mass number = number of protons + number of neutrons = atomic number + neutron number 7 ISOTOPES Isotopes are atoms of the same element with different masses; they are atoms containing the same number of protons but different numbers of neutrons 8 The three isotopes of Hydrogen 1 amu = 1.660 x 10-24 g Particle Mass (g) Charge Proton 1.6727 x 10 -24 +1 Neutron 1.6750 x 10 -24 0 Electron 9.110 x 10 -28 -1 9 THE ATOMIC WEIGHT SCALE AND ATOMIC WEIGHTS  The atomic weight scale is based on the mass of the carbon-12 isotope  One amu is exactly 1/12 of the mass of a carbon-12 atom: 1 g = 6.022 x 1023 amu or 1 amu = 1.660 x 10-24 g 10 THE ATOMIC WEIGHT SCALE AND ATOMIC WEIGHTS Atomic weight = 0.7899 (23.98504 amu) + 0.1000 (24.98584 amu) + 0.1101 (25.98259 amu) = 18.946 amu + 2.4986 amu + 2.8607 amu = 24.30 amu 11 ELECTRONIC STRUCTURES OF ATOMS  Why do different elements have such different chemical and physical properties?  Why does chemical bonding occur at all?  Why does each element form compounds with characteristic formulas?  How can atoms of different elements give off or absorb light only of characteristic colors. 12 13 Electromagnetic Radiation c Where: frequency wavelength c: speed of light c = 2.99 x 108 m/s 14 Electromagnetic Radiation 15 Photons The quantum of electromagnetic energy, generally regarded as a discrete particle having zero mass, no electric charge, and an indefinitely long lifetime. E = hν = hc/λ h = Planck's constant = 6.626 × 10−34 J.s 16 17 18 Electromagnetic Spectrum 19 Dispersion of White Light 20 EMISSION & ABSORPTION SPECTRA 21 ATOMIC SPECTRA 22 Bohr Model In 1913, Niels Bohr (1885–1962):  The electronic energy is quantized: only certain values of electronic energy are possible.  The electrons absorb or emit energy in discrete amounts as they move from one orbit to another. 23 Bohr Model 24 Bohr Model for Hydrogen Atom  Allowed orbits: mvr = nh/2 r = n2a0 a0 = 5.292 x 10 11 m = 0.5292 Å  n = quantum number = 1, 2, 3, 4, 5, 6, etc Potential energy: 25 Ground State The state of least possible energy in a physical system, as of elementary particles. Also called ground level. 26 Excited State Being at an energy level higher than the ground state. 27 Electron Transition in a Hydrogen Atom Lyman series → ultraviolet n>1→ n=1 Balmer series → visible light n>2→n=2 Paschen series → infrared n>3→ n=3 28 29 30 Quantum Mechanics Theory of the structure and behavior of atoms and molecules. 31 32 33 34 35 36 The Schrödinger Equation  Has been solved exactly only for oneelectron species such as the hydrogen atom and the ions He+ and Li2+ .  Simplifying assumptions are necessary to solve the equation for more complex atoms and molecules. 37 Basic Ideas of Quantum Mechanics 1. Atoms and molecules can exist only in certain energy states. In each energy state, the atom or molecule has a definite energy. When an atom or molecule changes its energy state, it must emit or absorb just enough energy to bring it to the new energy state (the quantum condition). Atoms and molecules possess various forms of energy. Let us focus our attention on their electronic energies. 38 2. When atoms or molecules emit or absorb radiation (light), they change their energies. The energy change in the atom or molecule is related to the frequency or wavelength of the light emitted or absorbed by the equations: ΔE = hν or ΔE = hc/λ The energy lost (or gained) by an atom as it goes from higher to lower (or lower to higher) energy states is equal to the energy of the photon emitted (or absorbed) during the transition. 39 3. The allowed energy states of atoms and molecules can be described by sets of numbers called quantum numbers: n, l, m 40 41 42 43 Atomic Orbitals  An atomic orbital is a region of space around the nucleus in which the probability of finding an electron is high.  Determined by a set of quantum numbers: n, l, m.  4 types: s, p, d, f. 44 Atomic Orbitals, s-type 45 Atomic Orbitals, p-type 46 Atomic Orbitals, d-type 47 48 49 Electronic Configurations • The shorthand representation of the occupancy of the energy levels (shells and subshells) of an atom by electrons. 50 51 52 53 54 55 56 57 58 Hund's Rules 59 Electronic Configuration H atom (1 electron): 1s1 He atom (2 electrons): 1s2 Li atom (3 electrons): 1s2, 2s1 Cl atom (17 electrons): 1s2, 2s2, 2p6, 3s2, 3p5 60 Electronic Configuration As atom (33 electons): 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p3 or [Ar] 4s2, 3d10, 4p3 61 62 63 Electronic Configuration Negative ions: add electron(s), 1 electron for each negative charge S-2 ion: (16 + 2) electrons: 1s2, 2s2, 2p6, 3s2, 3p6 64 Electronic Configuration Positive ions remove electron(s), 1 electron for each positive charge Mg+2 ion: (12-2) electrons: 1s2, 2s2, 2p6 65 How many valence electrons are in Cl: [Ne]3s2 3p5? 2, 5, 7 66 For Cl to achieve a noble gas configuration, it is more likely that: electrons would be added electrons would be removed 67 68 Regions by Electron Type 69 70 71 Trends in the Periodic Table • • • • Atomic radius Ionic radius Ionization energy Electron affinity 72 Atomic Radius decrease left to right across a period Zeff = Z - S where: Zeff = effective nuclear charge Z = nuclear charge, atomic number S = shielding constant 73 Atomic Radius  Increase top to bottom down a group  Increases from upper right corner to the lower left corner 74 Atomic Radius 75 Atomic Radius vs. Atomic Number 76 Ionic Radii 77 Ionic Radii • Same trends as for atomic radius 78 Comparison of Atomic and Ionic Radii  Positive ions smaller than atom  Negative ions larger than atom 79 Ionic Radii Isoelectronic Series • series of negative ions, noble gas atom, and positive ions with the same electronic confiuration • size decreases as “positive charge” of the nucleus increases 80 Ionization Energy • energy necessary to remove an electron to • • • form a positive ion low value for metals, electrons easily removed high value for non-metals, electrons difficult to remove increases from lower left corner of periodic table to the upper right corner 81 Ionization Energy first ionization energy Energy to remove first electron from an atom. second ionization energy Energy to remove second electron from a +1 ion, etc. 82 Ionization Energy vs. Atomic Number 83 Electron Affinity • Energy released when an electron is added to an atom • Same trends as ionization energy, increases from lower left corner to the upper right corner • Metals have low “EA” • Nonmetals have high “EA” 84 Magnetism • Result of the spin of electrons • Diamagnetism: no unpaired electrons • Paramagnetism: one or more unpaired electrons 85 86 87