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Chemistry-140 Lecture 20 Chapter 8: Electron Configurations & Chemical Periodicity Chapter Highlights electron spin Pauli Exclusion Principle atomic electron configurations atomic properties periodic trends chemical reactions & periodic properties Chemistry-140 Lecture 20 Magnetism & Electron Spin North pole of a magnet producing a magnetic field N N Electron spinning counterclockwise Electron spinning clockwise S South pole of a magnet producing a magnetic field S Chemistry-140 Lecture 20 Magnetism in Chemical Substances o.ox g Paramagnetism a substance is attracted to a magnetic field Sample sealed in a glass tube Electronic balance Diamagnetism: a substance is repelled by a magnetic field Electromagnet to provide magnetic field Chemistry-140 Lecture 20 Quantum Numbers The electron spin magnetic quantum number (ms) has possible values of: 1 1 ms = + or 2 2 It describes the electron character of an electron in an orbital. Electron spins are either paired or unpaired If there is only one electron the value of ms is arbitrary Chemistry-140 Lecture 20 Pauli Exclusion Principle Pauli exclusion principle: no two electrons can have the same set of four quantum numbers ( n l ml ms ). Only two electrons can be in the same orbital (sharing the same values of n l ml ). They must have opposite spins, ms = + 1/2 and - 1/2. Chemistry-140 Lecture 20 Pauli Exclusion Principle Represent the 1s orbital of an H atom as This orbital has quantum numbers n=1 l ml = 0 =0 One electron can occupy this 1s orbital with a value of ms = + 1/2 H: 1s or ms = - 1/2 H: 1s Chemistry-140 Lecture 20 Pauli Exclusion Principle One electron can occupy the 1s orbital of He with a value of ms = + 1/2 He: 1s or ms = - 1/2 He: 1s The second electron must occupy the 1s orbital of He such that the spins are paired He: 1s Chemistry-140 Lecture 20 Electron Configurations in Multi-Electron Atoms Electron configuration: an arrangement of electrons in atomic orbitals The most stable electron configurations are those that minimize energy. Each element has a ground state electron configuration in which all the electrons are in the lowest-energy orbitals. Any other configuration is an excited state Chemistry-140 Lecture 20 Determining Atomic Electron Configurations The lowest energy configuration is most favourable The Pauli exclusion principle must be obeyed. Follow Hund's rule. Electrons in a set of degenerate orbitals (those with equal energy) will go into separate orbitals if that option is available. Chemistry-140 Lecture 20 Orbital Energies in Multi-Electron Atoms Recall for Hydrogen 1 En = -RH 2 n Bohr's equation can be extended to multi-electron atoms since the energy of an electron is proportional to the square of the nuclear charge, Z: Z2 En = -RH 2 n Chemistry-140 Lecture 20 Orbital Energies in Multi-Electron Atoms Nuclear charge "felt" by an electron is called the effective nuclear charge, Z*. Z* = Z - S (S = screening by closer electrons) Due to effective nuclear charge, the energies of orbital subshells vary for a value of n. For a given value of n, Z* decreases as l increases. So the energies of the subshells increase as l increases, ns < np < nd < nf Chemistry-140 Lecture 20 Orbital Energies in Multi-Electron Atoms Chemistry-140 Lecture 20 Orbital Energies for a Multi-Electron Atom Chemistry-140 Lecture 20 Electron Configurations in Multi-Electron Atoms For the simplest case (the Hydrogen atom), one electron occupies the 1s orbital H: 1s This can be represented as: and for Helium: He: 1s H: 1s1 becomes He: 1s2 Chemistry-140 Lecture 20 Valence Electrons Valence electrons: the outer-shell electrons - those electrons present beyond the last full subshell or preceding noble gas. For example: Aluminum has the ground state electron configuration Al: 1s22s22p63s23p1 three valence electrons; 3s23p1 Chemistry-140 Lecture 20 Valence Electrons Core electrons: electrons that are not valence electrons. Those which comprise the preceding noble gas electron configuration. Al: 1s22s22p63s23p1 Al: [Ne]3s23p1 This emphasizes the valence electrons which are used in chemical reactions and bonding. Chemistry-140 Lecture 20 Electron Configurations & the Periodic Table Chemistry-140 Lecture 20 Electron Configurations for Transition Metals Some anomalies occur for the transition elements. V: [Ar]4s23d3 Cr: [Ar]4s13d5 . . . Ni: [Ar]4s23d8 Cu: [Ar]4s13d10 Chemistry-140 Lecture 20 Electron Configurations for Ions Cl:[Ne]3s23p5 Cl-:[Ne]3s23p6 Ca:[Ar]4s2 Ca2+:[Ar] Ni:[Ar]4s23d8 Ni2+:[Ar]3d8 Cu:[Ar]4s13d10 Cu+:[Ar]3d10 cuprous ion Cu2+:[Ar]3d9 cupric ion Cl:[Ar] Chemistry-140 Lecture 21 Chapter 8: Electron Configurations & Chemical Periodicity Chapter Highlights electron spin Pauli Exclusion Principle atomic electron configurations atomic properties periodic trends chemical reactions & periodic properties Chemistry-140 Lecture 21 Atomic Properties & Periodic Trends Members of the same Group have similar chemical properties and the same valence shell electron configurations Example Group 2 Elements: M M M2+ + 1/2 O2 Be: [He]2s2 Mg: [Ne]3s2 Ca: [Ar]4s2 Sr: [Kr]5s2 + 2 e- MO Chemistry-140 Lecture 21 Periodic Trends A Comparison of Mendeleev's eka-Silicon to Germanium Property Prediction for Observed for eka-Silicon (1871) Germanium Atomic weight 72 72.59 Density (g/cm3) 5.5 5.35 Specific Heat (J/g-K) 0.305 0.309 Melting Point (oC) High 947 Colour Dark grey Greyish White Formula of Oxide XO2 GeO2 Density of Oxide (g/cm3) 4.7 4.70 Formula of Chloride XCl4 GeCl4 Boiling Point of Chloride (oC) A little under 100 84 Chemistry-140 Lecture 21 Sizes of Atoms The effective nuclear charge, Z*, increases from left to right in a period and from bottom to top in a group. Z* = Z - S (S = screening by core electrons) The greater Z*, the greater the attractive force between the nucleus and its electrons The greater Z*, the smaller the average distance between the nucleus and its valence electrons (the atomic radius) Chemistry-140 Lecture 21 Atomic Radii C 154 pm Cl 200 pm 176 pm Chemistry-140 Lecture 21 Atomic Radii 1 2 3 4 5 6 7 8 Chemistry-140 Lecture 21 Ionization Energies Ionization energy, IE: The energy required to remove the outermost electron from a gaseous atom or ion. The first ionization energy, I1 is the energy for the removal of an electron from a neutral, gaseous atom: Atom in ground state(g) Atom(g)+ + e- DE = ionization energy, IE Chemistry-140 Lecture 21 Ionization Energies The second ionization energy, I2, is the energy of the removal of a second electron from a gaseous ion with a 1+ charge: M(g)+ M(g)2+ + e- As Z* increases, it requires more energy to remove the outermost electron. Consequently, ionization energy is related to atomic radius, with ionization energy increasing as atomic radius decreases Chemistry-140 Lecture 21 Trends in Ionization Energies Ionization energies can be correlated to electron configuration: Mg(g) Mg: [Ne]3s2 Mg(g)+ + eMg+: [Ne]3s1 IE(1) = 738 kJ/mol Mg(g)+ Mg+: [Ne]3s1 Mg(g)2+ + eMg2+: [Ne] IE(2) = 1451 kJ/mol Mg(g)2+ Mg2+: [Ne] Mg(g)3+ + eMg3+: [He]2s22p5 IE(3) = 7733 kJ/mol Chemistry-140 Lecture 21 Trends in Ionization Energies Chemistry-140 Lecture 21 First Ionization Energies Chemistry-140 Lecture 21 Electron Affinities Electron affinity, EA: the energy change of the reaction of adding an electron to a gaseous atom or ion: Atom in ground state(g) + e- Atom(g)- DE = electron affinity, EA These reactions tend to be exothermic and so the values of EA are generally negative. In general, electron affinity tends to decrease from left to right. Going down a group, there is little change Chemistry-140 Lecture 21 Electron Affinities Chemistry-140 Lecture 21 Ion Size (Cations) Li atom (radius = 152 pm) Li+ cation (radius = 90 pm) Li Li+ 152 pm 1s 2s 90 pm 1s 2s Chemistry-140 Lecture 21 Ion Size (Anions) F atom (radius = 72 pm) F- anion (radius = 119 pm) F- F 72 pm 2s 2p 119 pm 2s 2p Chemistry-140 Lecture 21 Trends in Ion Size Chemistry-140 Lecture 21 Ion Sizes in an Isoelectronic Series Ion Ionic radius (pm) 2- O - F + Na Mg 126 119 116 2+ 86 Number of protons 8 9 11 12 Number of electrons 10 10 10 10 Chemistry-140 Lecture 21 Textbook Questions From Chapter # 8 Electronic configurations: 10, 12, 16, 23, 26, 32 Quantum numbers: 36, 41, 42 Periodic Properties: 45, 46, 52, 54 General & Conceptual 60, 61, 73, 75