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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 6 Oct 1997 Chemical Periodicity 1 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. 6 Oct 1997 Chemical Periodicity 2 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. 6 Oct 1997 Chemical Periodicity 3 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 6 Oct 1997 Chemical Periodicity O2 4 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.” 6 Oct 1997 Chemical Periodicity 5 QUANTUM NUMBERS 6 Oct 1997 n (shell) 1, 2, 3, 4, ... (subshell) 0, 1, 2, ... n - 1 m (orbital) - ... 0 ... + ms (electron spin) +1/2, -1/2 Chemical Periodicity 6 Shells, Subshells, Orbitals #orbitals 0 s 1 0 s 1 1 p 3 0 s 1 1 p 3 2 d 5 0 s 1 1 p 3 2 d 5 3 f 7 0..(n-1) (2 +1) n 1 2 3 4 n 6 Oct 1997 #e2 2 6 2 6 10 2 6 10 14 2*(2 +1) Chemical Periodicity Total 2 8 PERIOD 1 (H, He) 2 (Li…Ne) 3 (Na .. Ar) 18 32 2n2 =0 =1 =2 =3 s p d f etc, for n = 5, 6 7 Element Mnemonic Competition Hey! Here Lies Ben Brown. Could Not Order Fire. Near Nancy Margaret Alice Sits Peggy Sucking Clorets. Are Kids Capable ? 6 Oct 1997 Chemical Periodicity 8 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) ... 6 Oct 1997 Chemical Periodicity 9 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. (n + )= 5 (n + )= 4 5_manyelE.mov 6 Oct 1997 Chemical Periodicity 10 Effective Nuclear Charge • The difference in SUBSHELL energy e.g. 2s and 2p subshells is due to effective nuclear charge, Z*. Charge felt by 2s e- of Li atom 6 Oct 1997 2s e- spends more time close to Li3+ nucleus than the 2p eTherefore 2s is lower in E than 3s Chemical Periodicity 11 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 Be B 6 Oct 1997 Z* = 3 - 2 = 1 Z* = 4 - 2 = 2 Z* = 5 - 2 = 3 and so on! Chemical Periodicity 12 Photoelectron Spectroscopy - Measuring IE Photoelectric effect: h + A A+ + eforms basis for DIRECT determination of IE Kinetic energy of electron = h - IE VALENCE therefore: IE = h - KE(e ) ELECTRONS Signal 1s Ne Inner shell or CORE ELECTRONS 2p Ar 1s 309 6 Oct 1997 2s 2s 50 100 Chemical Periodicity 2p 3s 3p 0 IE (MJ/mol) 13 Electron Filling Order 6 Oct 1997 Chemical Periodicity (Figure 8.7) 14 Writing Atomic Electron Configurations Two ways of writing configurations. One is called the spectroscopic notation: SPECTROSCOPIC NOTATION for H, atomic number = 1 1 no. of electrons 1s value of n 6 Oct 1997 value of l Chemical Periodicity 15 Writing Atomic Electron Configurations (2) A second way is called the orbital box notation. ORBITAL BOX NOTATION for He, atomic number = 2 Arrows 2 depict electron spin 1s 1s One electron has n = 1, = 0, ml = 0, ms = + 1/2 Other electron has n = 1, = 0, ml = 0, ms = - 1/2 6 Oct 1997 Chemical Periodicity 16 Electron Configuration tool - see “toolbox”. 6 Oct 1997 Chemical Periodicity 17 Beryllium Lithium Group 2A Z=4 1s22s2 Group 1A Z=3 1s22s1 3p 3p 3s 3s 2p 2p 2s 2s 1s 1s 6 Oct 1997 Chemical Periodicity 18 Boron Carbon Z=5 Z=6 1s2 2s2 2p1 1s2 2s2 2p2 3p 3p 3s 3s 2p 2p 2s 2s Why not ? 1s 6 Oct 1997 1s Chemical Periodicity 19 Carbon Z=6 1s2 2s2 2p2 3p 3s 2p 2s 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. 1s 6 Oct 1997 Chemical Periodicity 20 Nitrogen Oxygen Z=7 Z=8 1s2 2s2 2p3 1s2 2s2 2p4 3p 3p 3s 3s 2p 2p 2s 2s 1s 1s 6 Oct 1997 Chemical Periodicity 21 Fluorine Neon Z=9 Z = 10 1s2 2s2 2p5 1s2 2s2 2p6 3p 3s 3p 3s 2p 2p 2s 1s 6 Oct 1997 2s 1s Chemical Periodicity Note that we have reached the end of the 2nd period, . . . and the 2nd shell is full! 22 GROUPS and PERIODS Sodium Z = 11 1s2 2s2 2p6 3s1 or “neon core” + 3s1 [Ne] 3s1 (uses rare gas notation) Na begins a new period. Li Na K Rb Cs All Group 1A elements: have [core] ns1 configurations. (n = period #) 6 Oct 1997 Chemical Periodicity 23 Periodic Chemical Properties REACTIVITY 5_Li.mov 5_Na.mov 5_K.mov 6 Oct 1997 SIZE IE (Ionization Energy) Li Be Na Mg K Ca Rb Sr Cs Ba Alkalis Alkaline Earths Chemical Periodicity 24 Alkaline Earths Metals (ns2) - easily oxidized to M2+ - less reactive than alkalis of same period reactivity: Be < Mg < Ca < Sr < Ba WHY? - • 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 6 Oct 1997 Chemical Periodicity 25 Relationship of Electron Configuration and Regions of the Periodic Table s block d block p block f block 6 Oct 1997 Chemical Periodicity 26 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 6 Oct 1997 Iron Chemical Periodicity Copper 27 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 3p 3p 3s 3s 2p 2p 6 Oct 1997 2s 2s 1s 1s Chemical Periodicity 28 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 Fe2+ Fe 4s 4s 3d E4s ~ E3d - exact energy of orbitals depend on whole configuration 6 Oct 1997 3d Fe3+ 4s Chemical Periodicity 3d 29 Ion Configurations (3) How do we know the configurations of ions? From the magnetic properties of ions. Ions (or atoms) with UNPAIRED ELECTRONS are: PARAMAGNETIC. Ions (or atoms) without unpaired electrons are: DIAMAGNETIC. 6 Oct 1997 Chemical Periodicity 30 General Periodic Trends • Atomic and ionic radii : SIZE • Ionization energy : E(A+) - E(A) • Electron affinity : E(A-) - E(A) Higher Z*. Electrons held more tightly. Larger orbitals. Electrons held less tightly. 6 Oct 1997 Chemical Periodicity 31 Atomic Size INCREASES down a Group • Size goes UP on going down a GROUP • Because electrons are added further from the nucleus, there is less attraction. 6 Oct 1997 Chemical Periodicity 32 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. 6 Oct 1997 Chemical Periodicity 33 Atomic Radii 6 Oct 1997 Chemical Periodicity 34 Trends in Atomic Size (Figure 8.10) Radius (pm) 250 K 1st transition series 3rd period 200 Na 2nd period Li 150 Kr 100 Ar Ne 50 He 0 0 5 10 15 20 25 30 35 40 Atomic Number 6 Oct 1997 Chemical Periodicity 35 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! 6 Oct 1997 Chemical Periodicity 36 Ion Sizes - CATIONS Does the size go up or down when an atom loses an electron to form a cation? + Li, 152 pm 3 e-, 3 p Forming a cation Li+, 60 pm 2 e-, 3 p • CATIONS are SMALLER than the parent atoms. • The electron/proton attraction goes UP so size DECREASES. 6 Oct 1997 Chemical Periodicity 37 Ion Sizes - ANIONS Does the size go up or down when gaining an electron to form an anion? F, 64 pm 9 e-, 9 p Forming an anion F-, 136 pm 10 e-, 9 p • ANIONS are LARGER than the parent atoms. • electron/proton attraction goes DOWN so size INCREASES. 6 Oct 1997 Chemical Periodicity 38 Trends in Ion Sizes CATIONS ANIONS (59 pm) (207 pm) Trends in relative ion sizes are the same as atom sizes. 6 Oct 1997 Chemical Periodicity 39 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 6 Oct 1997 Chemical Periodicity 40 Ionization Energy (IE) Mg (g) atom [Ne]2s 735 kJ Mg+ (g) + e- [Ne]2s1 Mg+ (g) + 1451 kJ Mg2+ (g) + e- [Ne]2s0 Mg (g) + Mg3+ Mg2+ (g) + 7733 kJ Mg3+ (g) + e- [He]2s22p5 Mg2+ • Energy ‘cost’ is very high to remove an INNER SHELL e- (shell of n < nVALENCE). • This is why oxidation. no. = Group no. 6 Oct 1997 Chemical Periodicity Mg+ Mg 41 Trends in First Ionization Energy 1st Ionization energy (kJ/mol) 2500 He Ne 2000 Ar 1500 Kr 1000 500 0 1 H 6 Oct 1997 3 Li 5 7 9 11 Na 13 15 17 19 21 K Chemical Periodicity 23 25 27 29 31 33 35 Atomic Number 42 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 6 Oct 1997 Chemical Periodicity 43 2nd IE / 1st IE 2nd IE: A+ A++ + e- Li Na K 6 Oct 1997 Chemical Periodicity 44 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 6 Oct 1997 Chemical Periodicity 45 Trends in Electron Affinity (Table 8.5, Figure 8.14) Atom EA (kJ) • Affinity for electron increases B C across a period (EA becomes more negative). N O -27 -122 0 -141 F -328 F Cl Br I -328 -349 -325 -295 • Affinity decreases down a group (EA becomes less negative). 6 Oct 1997 Chemical Periodicity 46 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 6 Oct 1997 Chemical Periodicity 47