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Electron Configurations  Chemical Periodicity (Ch 8)

Electron Configurations  Chemical Periodicity (Ch 8). Electron spin & Pauli exclusion principle configurations spectroscopic, orbital box notation Hund’s rule - electron filling rules configurations of ATOMS: the basis for chemical valence

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Electron Configurations  Chemical Periodicity (Ch 8)

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  1. Electron Configurations  Chemical Periodicity (Ch 8) • Electron spin & Pauli exclusion principle • configurations • spectroscopic, orbital box notation • Hund’s rule - electron filling rules • configurations of ATOMS: • the basis for chemical valence • configurations and properties of IONS • periodic trends in : • size • ionization energies • electron affinities Na + Cl  NaCl Mg + O2 MgO Chemical Periodicity

  2. Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS () ORBITALS (m) Each orbital can be assigned up to 2 electrons! WHY ? . . . Because there is a 4th quantum number, the electron spin quantum number, ms. Chemical Periodicity

  3. Electron Spin Quantum Number, ms • It can be proved experimentally that the electron has a spin. This is QUANTIZED. • The two allowed spin directions are defined by the magnetic spin quantum number, ms • ms = +1/2 and -1/2 ONLY. Chemical Periodicity

  4. Electron Spin Quantum Number MAGNETISM is a macroscopic result of quantized electron spin 5_magnet.mov Diamagnetic: NOT attracted to a magnetic field All electrons are paired N2 Paramagnetic: attracted to a magnetic field. Substance has unpaired electrons O2 Chemical Periodicity

  5. Pauli Exclusion Principle • electrons with the same spin keep as far apart as possible • electrons of opposite spin may occupy the same “region of space” (= orbital) • Consequences: • No orbital can have more than 2 electrons • No two electrons in the same atom can have the same set of 4 quantum numbers (n, l, ml, ms) OR • “Each electron has a unique address.” Chemical Periodicity

  6. QUANTUMNUMBERS n (shell) 1, 2, 3, 4, ...  (subshell) 0, 1, 2, ... n - 1 m (orbital) -  ... 0 ... +  ms (electron spin) +1/2, -1/2 Chemical Periodicity

  7. Shells, Subshells, Orbitals n  #orbitals #e- Total PERIOD 1 0 s 1 2 2 1 (H, He) 2 0 s 1 2 1 p 3 6 8 2 (Li…Ne) 3 0 s 1 2 1 p 3 6 3 (Na .. Ar) 2 d 5 10 18 4 0 s 1 2 1 p 3 6 2 d 5 10 3 f 7 14 32 n 0..(n-1) (2 +1) 2*(2 +1) 2n2  = 0 s  = 1 p  = 2 d  = 3 f etc, for n = 5, 6 Chemical Periodicity

  8. Element Mnemonic Competition Hey! Here Lies Ben Brown. Could Not Order Fire. Near Nancy Margaret Alice Sits Peggy Sucking Clorets. Are Kids Capable ? WHAT’s YOURs ?? Chemical Periodicity

  9. Assigning Electrons to Atoms • Electrons are assigned to orbitals successively in order of the energy. • For H atoms, E = - R(1/n2). E depends only on n. • For many-electron atoms, orbital energy depends on both n and . • E(ns) < E(np) < E(nd) ... Chemical Periodicity

  10. (n + )= 5 (n + )= 4 Assigning Electrons to Subshells • In H atom all subshells of same n have same energy. • In many-electron atom: a) subshells increase in energy as value of (n + ) increases. b) for subshells of same (n +), subshell with lower n is lower in energy. 5_manyelE.mov Chemical Periodicity

  11. 2s e- spends more time close to Li3+ nucleus than the 2p e- Therefore 2s is lower in E than 3s Charge felt by 2s e- of Li atom Effective Nuclear Charge • The difference in SUBSHELL energy e.g. 2s and 2p subshells is due to effective nuclear charge, Z*. Chemical Periodicity

  12. Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by an electron. • Z* increases across a period owing to incomplete shielding by inner electrons. • For VALENCE electrons we estimate Z* as: Z* = [ Z - (no. of inner electrons)] • Charge felt by 2s e- in Li Z* = 3 - 2 = 1 • Be Z* = 4 - 2 = 2 • B Z* = 5 - 2 = 3 • and so on! Chemical Periodicity

  13. VALENCE ELECTRONS Ne Inner shell or CORE ELECTRONS 1s 2s 2p Signal Ar 1s 2p 3s 3p 2s 50 309 100 0 IE (MJ/mol) Photoelectron Spectroscopy - Measuring IE Photoelectric effect: h + A  A+ + e- forms basis for DIRECT determination of IE Kinetic energy of electron = h - IE therefore: IE = h - KE(e-) Chemical Periodicity

  14. Electron Filling Order(Figure 8.7) Chemical Periodicity

  15. Writing Atomic Electron Configurations Two ways of writing configurations. One is called the spectroscopic notation: Chemical Periodicity

  16. Writing Atomic Electron Configurations (2) A second way is called the orbital box notation. One electron has n = 1,  = 0,ml = 0,ms = + 1/2 Other electron has n = 1,  = 0, ml = 0,ms = - 1/2 Chemical Periodicity

  17. Electron Configuration tool - see “toolbox”. Chemical Periodicity

  18. BerylliumGroup 2AZ = 41s22s2 LithiumGroup 1AZ = 31s22s1 Chemical Periodicity

  19. CarbonZ = 61s2 2s2 2p2 Why not ?  BoronZ = 51s2 2s2 2p1 Chemical Periodicity

  20. CarbonZ = 61s2 2s2 2p2 The configuration of C is an example of HUND’S RULE: the lowest energy arrangement of electrons in a subshell is that with the MAXIMUM no. of unpaired electrons Electrons in a set of orbitals having the same energy, are placed singly as long as possible. Chemical Periodicity

  21. OxygenZ = 81s2 2s2 2p4 NitrogenZ = 71s2 2s2 2p3 Chemical Periodicity

  22. NeonZ = 101s2 2s2 2p6 FluorineZ = 9 1s2 2s2 2p5 Note that we have reached the end of the 2nd period, . . . and the 2nd shell is full! Chemical Periodicity

  23. Sodium Z = 11 1s2 2s2 2p6 3s1 GROUPS and PERIODS or “neon core” + 3s1 [Ne] 3s1 (uses rare gas notation) Na begins a new period. All Group 1A elements: Li Na K Rb Cs have [core] ns1 configurations. (n = period #) Chemical Periodicity

  24. REACTIVITY SIZE IE (Ionization Energy) Be Mg Ca Sr Ba Alkaline Earths Periodic Chemical Properties Li Na K Rb Cs Alkalis 5_Li.mov 5_Na.mov 5_K.mov Chemical Periodicity

  25. Metals (ns2) - easily oxidized to M2+ - less reactive than alkalis of same period reactivity: Be < Mg < Ca < Sr < Ba WHY? - Alkaline Earths • Size INCREASES as  group • VALENCE e- are farther from nucleus • same Z* - Valence e- less tightly held • Therefore valence e- are easier to remove Typical reactions / compounds Oxides: M +1/2O2 (g)  MO (s) CaO (lime) - #5 Ind. Chem Halides: M + X2 (g)  MX Carbonates: CaCO3 (limestone)  CaO + CO2 Sulfates: CaSO4.2H2O (gypsum)  CaSO4. 0.5H2O (plaster-of-paris) + 3/2H2O RECALL: Solubility rules and PRECIPITATION REACTIONS Chemical Periodicity

  26. p block s block d block f block Relationship of Electron Configuration and Regions of the Periodic Table Chemical Periodicity

  27. Transition Metals Table 8.4 • Transition metals (e.g. Sc .. Zn in the 4th period) have the configuration [argon] nsx (n - 1)dy • also called “d-block” elements. 3d orbitals used for Sc - Zn Chromium Iron Copper Chemical Periodicity

  28. Ion Configurations To form cations from elements : remove 1 e- (or more) from subshell of highest n [or highest (n + )]. P [Ne] 3s2 3p3 - 3e-  P3+ [Ne] 3s2 3p0 Chemical Periodicity

  29. Ion Configurations (2) Transition metals ions: remove ns electrons and then (n - 1)d electrons. Fe [Ar] 4s2 3d6 loses 2 electrons  Fe2+ [Ar] 4s0 3d6 E4s ~ E3d - exact energy of orbitals depend on whole configuration Chemical Periodicity

  30. Ion Configurations (3) From the magnetic properties of ions. Ions (or atoms) with UNPAIRED ELECTRONS are: PARAMAGNETIC. Ions (or atoms) without unpaired electrons are: DIAMAGNETIC. How do we know the configurations of ions? Chemical Periodicity

  31. General Periodic Trends • Atomic and ionic radii : SIZE • Ionization energy : E(A+) - E(A) • Electron affinity : E(A-) - E(A) Chemical Periodicity

  32. Atomic Size INCREASESdown a Group • Size goes UP on going down a GROUP • Because electrons are added further from the nucleus, there is less attraction. Chemical Periodicity

  33. Atomic Size DECREASES across a period Size goes DOWN on going across a PERIOD. Size decreases due to increase in Z*. Each added electron feels a greater and greater +ve charge. Chemical Periodicity

  34. Atomic Radii Chemical Periodicity

  35. Trends in Atomic Size (Figure 8.10) Chemical Periodicity

  36. Sizes of Transition Elements(Figure 8.11) • 3d subshell is inside the 4s subshell. • 4s electrons feel a more or less constant Z*. • Sizes stay about the same and chemistries are similar! Chemical Periodicity

  37. + Li, 152 pm 3 e-, 3 p Li+, 60 pm 2 e-, 3 p Ion Sizes - CATIONS Does the size go up or down when an atom loses an electron to form a cation? Forming a cation • CATIONS are SMALLER than the parent atoms. • The electron/proton attraction goes UP so size DECREASES. Chemical Periodicity

  38. - F, 64 pm 9 e-, 9 p F-, 136 pm 10 e-, 9 p Ion Sizes - ANIONS Does the size go up or down when gaining an electron to form an anion? Forming an anion • ANIONS are LARGER than the parent atoms. • electron/proton attraction goes DOWN so size INCREASES. Chemical Periodicity

  39. CATIONS ANIONS (59 pm) (207 pm) Trends in relative ion sizes are the same as atom sizes. Trends in Ion Sizes Chemical Periodicity

  40. Oxidation-Reduction Reactions • Why do metals lose electrons in their reactions? • Why does Mg form Mg2+ ions and not Mg3+? • Why do nonmetals take on electrons? - related to IE and EA Chemical Periodicity

  41. Mg3+ Mg2+ Mg+ Ionization Energy (IE) Mg (g) + 735 kJ  Mg+ (g) + e- [Ne]2s1 Mg (g) atom [Ne]2s Mg+ (g) + 1451 kJ  Mg2+ (g) + e- [Ne]2s0 Mg2+ (g) + 7733 kJ  Mg3+ (g) + e- [He]2s22p5 • Energy ‘cost’ is very high to remove an INNER SHELL e- (shell of n < nVALENCE). • This is why oxidation. no. = Group no. Mg Chemical Periodicity

  42. Trends in First Ionization Energy Chemical Periodicity

  43. Trends in Ionization Energy (2) • IE increases across a period because Z* increases. • Metals lose electrons more easily than nonmetals. • Metals are good reducing agents. • Nonmetals lose electrons with difficulty. • IE decreases down a group • Because size increases, reducing ability generally increases down the periodic table. • E.g. reactions of Li, Na, K Chemical Periodicity

  44. 2nd IE / 1st IE 2nd IE: A+ A++ + e- Li Na K Chemical Periodicity

  45. Electron Affinity (EA) • A few elements GAIN electrons to form anions. • Electron affinity is the energy released when an atom gains an electron. A(g) + e- A-(g) E.A. = DE = E(A-) - E(A) • If E(A-) < E(A) then the anion is more stable and there is an exothermic reaction Chemical Periodicity

  46. F -328 Cl -349 Br -325 I -295 • Affinity decreases down a • group • (EA becomes less negative). Trends in Electron Affinity(Table 8.5, Figure 8.14) Atom EA (kJ) B -27 C -122 N 0 O -141 F -328 • Affinity for electron increases across a period (EA becomes more negative). Chemical Periodicity

  47. SUMMARY • Electron spin: diamagnetism vs. paramagnetism • Pauli exclusion principle - allowable quantum numbers • configurations • spectroscopic notation • orbital box notation • Hund’s rule - electron filling rules • configurations of ATOMS: the basis for chemical valence • period 2 ; groups • transition metals • configurations and properties of IONS • periodic trends in : • size • ionization energies • electron affinities Chemical Periodicity

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