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Chapter 8 Periodicity 2010

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Chapter 8 Periodicity 2010 EM radiation Announcements is a wave and a units called particle --Exam 3 Oct 3 are Quant having Includes chapters 7/8/9/10 a The excluded items include the following. These won’t appear on a quiz on in any Exam. absorbe emitted amplitude energy frequency wavelength 1. Classical distinction between energy and matter (p. 217) d 2. Numerical problems involving the Rydberg equation (equations 7.3 and 7.4) 3. Spectral analysis in the laboratory (pp. 226-227) involve energy changes in related by related by 4. Numerical problems involving the Heisenberg uncertainty principle (p. 231) 5. Trends among the transition elements (p. 261) atomsatoms 6. Trends in electron affinity (pp. 265-266) c = !v 7. Pseudo-noble gas configuration (p. 269) electron E = hv 8. Lattice energy (pp. 283-285) 9. IR spectroscopy (p. 292) s molecules 10. Numerical problems involving electronegativity (p. 296) described by 11. Electronegativity and oxidation number (p. 297) 12. Section 11.3: MO theory and electron delocalization 13. All sections in chapter 12 except 12.3 (types of intermolecular forces). e- filling e- configuration comprising Wave functions gives determined by --Quiz 6 Next Week Tuesday Chapter 6 having quantum uunmbers Aufbau Rules Quantum described by Core Numbers Electrons Wave Function e- filling spdf electronic configuration comprising Orbital size (Orbital) Valence Principal define & energy determined by Electrons n = 1,2,3,.. s described by Aufbau basis for defines Rules Periodic Quantum which involve Table Angular Numbers momentum, l Orbital which summarizes defines Oribital Pauli Hund’s shape which are Energy Exclusion Rule Periodic Properties Magnetic Orbital defines ml orientation Electron Spin, ms defines spin Learning Objectives • Understand the meaning of periodic law • Write electron configurations for elements and ions. • Understand effective nuclear charge and how it can be used used to explain atomic radii, ionization energy. • Describe how the radii of ions relate to those of atoms. • Explain the concept of an isoelectronic series in ions • Predict ionization energy, electron affinity of elements. • Explain the observed changes in values of the successive ionization energies for a given atom. • Explain the general variations in first ionization energies among the elements and relate these variations to variations Electronic Configuration and Periodicity in atomic radii. • Explain the variations in electron affinities among the elements. • Trends and Variation of Chemical Properties of the Main Chapter 8 Group Elements Each electron in an atom can described by 4 quantum A shell describes the n quantum number (n = 1 = K; n = 2 = numbers n, l, ml, ms. You can think of the quantum L and so on). A subshell refers to the specific n,l quantum numbers as specifying the “house” or “wavefunction” or number as in a 2s subshell or 2p subshell or 3s subshell or “orbital” where the electron is in space. 3d subshell. Each atomic orbital can hold two electrons of opposite spin! angular Name Symbol Permitted Values Property principal momentum magnetic principal n positive integers (1,2,3...) orbital energy (size) angular l orbital shape (0, 1, 2, and angular momentum integers from 0 to n-1 3 correspond to s, p, d, and momentum f orbitals, respectively.) magnetic integers from -l to 0 to +l orbital orientation in magnetic ml space - spin +1/2 or -1/2 direction of e spin spin ms The shape of an orbital is given by the “l” quantum The order of filling of the orbitals can be number (l = {0,1,2 up to n-1}. The number of orbitals remembered using a mnemonic device. and its orientation in space is given by the angular Memorize this to help you! momentum quantum number ml {-l,..0..+l}. s-orbital p-orbital d-orbital The shapes are radial boundry plots that incorporate 95% probability of finding an electron. f-orbital How many 2p orbitals are there in an atom? n=2 If l = 1, then ml = -1, 0, or +1 2p 3 orbitals l = 1 If two electrons can be placed in one orbital how many electrons can go exist in the 3d subshell? n=3 If l = 2, then ml = -2, -1, 0, +1, or +2 3d 5 orbitals which can each holding 2 electrons for a total of 10 e- in the subshell l = 2 Determining Sublevel Names and Orbital Quantum Numbers The 3 p orbital Boundary Surface Plots (Shapes) Give the orbital name, possible magnetic quantum numbers, and number of orbitals for each sublevel with the following quantum numbers: (a) n = 3, l = 2 (c) n = 5, l = 1 (b) n = 2, l = 0 (d) n = 4, l = 3 SOLUTION: 2 px 2 py 2 pz n l Orbital name possible ml values # of orbitals (a) 3 2 3d -2, -1, 0, 1, 2 5 (b) 2 0 2s 0 1 Three p-orbitals (c) 5 1 5p -1, 0, 1 3 superimposed (d) 4 3 4f -3, -2, -1, 0, 1, 2, 3 7 The 5-d Boundry Surface Plots (Shapes) The seven f-orbitals (radial plots) dyz dxz dxy 2 dz Five d-orbitals superimposed 2 2 dx - dy Electronic configuration of the elements: The lowest energy (ground state) four quantum numbers describe an electron electronic configuration of all in a ground state atom. elements are constructed by filling lowest energy orbitals sequentially in what is called the “Aufbau Process”. 4s Name Symbol Permitted Values Property 3p 1. Lower energy (n-quantum number) principal n positive integers (1,2,3, orbital energy (size) orbitals fill first. 3s !) 2. Hund’s Rule-degenerate orbitals fill angular l 2p momentum integers from 0 to n-1 orbital shape one at a time before electrons are paired. 3. Pauli Exclusion Principle: No two orbital orientation in 2s magnetic ml integers from -l to 0 to electrons can have same 4-quantum +l space numbers) - spin ms +1/2 or -1/2 direction of e spin Electrons fill the lowest energy orbitals first, 2 at a time! 1s The order of filling of the orbitals can be Chemists use spdf notation and orbital box remembered using a mnemonic device. diagrams to denote or show the “ground state Memorize this to help you! electronic configuration” of elements. spdf orbital box Spin quantum Element Notation diagram number. An arrow denotes 1 an electron H 1s with “spin up” (+1/2) or “spin- He 1s2 down” (-1/2). n principal quantum # l quantum number # of electrons in orbital The Pauli Exclusion principle states: “No two Electronic configuration using Aufbau Process electrons can have the same 4-quantum numbers”. Atomic Orbital Box Full-electronic Condensed-electronic The spin numbers can not be the same (spin up Number/Element Diagram configuration configuration and spin down allowed only). 1 1 (n, l, ml and ms) H 1s 1s Example: He 1s2 1s2 Atomic Orbital Box Full-electronic Condensed-electronic Number/Element Diagram configuration configuration written with noble gas configuration 2 1 1 Li 1s22s1 [He]2s1 Li 1s 2s [He]2s Be 1s22s2 [He]2s2 Atomic Orbital Box Full-electronic Condensed-electronic Number/Element Diagram configuration configuration B 1s22s22p1 [He]2s22p1 2 2 2 2 2 C 1s 2s 2p [He]2s 2p 1s22s22p3 [He]2s22p3 1s22s22p4 [He]2s22p4 1s22s22p5 [He]2s22p5 1s22s22p6 [He]2s22p6 Anamolies occur when filling the d-orbitals Odd-filling behavior here! 4th and 9th position. Unpaired electrons in orbitals gives rise to Unpaired electrons in orbitals gives rise to paramagnetism and is attracted to a magnetic field. paramagnetism and is attracted to a magnetic field. Diamagnetic species contain all paired electrons and Diamagnetic species contain all paired electrons and is “repelled” by the magnetic field. is “repelled” by the magnetic field. Magnetic field off Magnetic field on Magnetic field on • Diamagnetic atoms or ions: Diamagnetic – All e- are paired. all electrons paired – Weakly repelled in a magnetic field. 2p Paramagentic Diamagentic • Paramagnetic atoms or ions: Paramagnetic Paramagnetic Diamagnetic – Unpaired e- exist in an orbital unpaired electrons – Attracted to an external magnetic field. 2p Metals and non-metals form ions with electronic Metals and non-metal ions tend to form electronic configurations closest to their nearest noble gas states closest to their nearest noble gas configuration. configuration. 1A 2A 3A 4A 5A 6A 7A 8A Metals loose electrons (oxidized) to become Isoelectronic species are two different elements cations. Non-metals gain electrons to become that have the same electronic configuration--but not anions. The electronic configuration of each the same nuclear configuration. reflects this change in the number of electrons. oxidation Na: [1s22s22p63s1] =====> Na+: [1s22s22p6] = [Ne] Na [Ne]3s1 Na+ [Ne] Metals lose electrons so oxidation 2 2 6 2 p1 3+ 2 2 6 Ca [Ar]4s2 Ca2+ [Ar] that cation has a noble-gas Al: [1s 2s 2p 3s 3 ] =====> Al : [1s 2s 2p ] = [Ne] 2 1 3+ outer electron configuration. reduced Al [Ne]3s 3p Al [Ne] N: [1s22s22p3] =====> N3-: [1s22s22p6] = [Ne] reduced O: [1s22s22p4] =====> O2-: [1s22s22p6] = [Ne] H 1s1 H- 1s2 or [He] reduced Non-metals gain F: [1s22s22p5] =====> F-: [1s22s22p6] = [Ne] F 1s22s22p5 F- 1s22s22p6 or [Ne] electrons so that anion has a noble-gas outer O 1s22s22p4 O2- 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all said to be “isoelectronic electron configuration. with Ne” as they have the same electronic N 1s22s22p3 3- 2 2 6 N 1s 2s 2p or [Ne] configuration....all subshells are filled. When a cation is formed from an atom of a Use condensed electron configurations to write the reaction transition metal, electrons are removed first from for the formation of each transition metal ion, and predict the ns orbital, then from the (n-1)d orbital.
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