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Unformatted text preview: Chapter 8 Chapter Atomic Electron Configurations Atomic and Chemical Periodicity and Chapter goals Chapter • Understanding the role magnetism Understanding plays in determining and revealing atomic structure. atomic • Understand effective nuclear charge Understand and its role in determining atomic properties. properties. • Write the electron configuration of Write neutral atoms and monatomic ions. neutral • Understand the fundamental physical Understand properties of the elements and their periodic trends. periodic Electron Spin and the Fourth Quantum Number and • The fourth quantum number is the spin quantum The number which has the symbol ms. number • The spin quantum number only has two possible The values. values. ms = +1/2 or −1/2 ms = ± 1/2 • This quantum number tells us the spin and This orientation of the magnetic field of the electrons. orientation • Wolfgang Pauli discovered the Exclusion Wolfgang Principle in 1925. Principle No two electrons in an atom can have the same No set of 4 quantum numbers, n, l , ml , and ms Electron Spin • Spin quantum number effects: – Every orbital can hold up to two Every electrons. electrons. • Consequence of the Pauli Exclusion Consequence Principle. Principle. – The two electrons are designated as The having having – one spin up ↑ ms = +1/2 one – and one spin down ↓ ms = −1/2 • Spin describes the direction of the Spin electron’s magnetic field. electron’s Paramagnetism and Diamagnetism Paramagnetism • Unpaired electrons have their spins Unpaired aligned ↑ ↑ or ↓ ↓ (in diff. orbitals) orbitals) – This increases the magnetic field of the This atom. atom. Total spin ≠ 0, because they add up. • Atoms with unpaired electrons are called Atoms paramagnetic . paramagnetic – Paramagnetic atoms are attracted to a Paramagnetic magnet. magnet. Paramagnetism and Diamagnetism Paramagnetism • Paired electrons have their spins Paired unaligned ↑↓. ( in the same orbital) ↑↓ – Paired electrons have no net magnetic Paired field. field. Total spin = 0, because of cancellation, Total ½ −½=0 • Atoms with no unpaired electrons are Atoms called diamagnetic. diamagnetic – Diamagnetic atoms are not attracted to a Diamagnetic magnet. magnet. Atomic Orbitals, Spin, and # of Electrons Atomic • Because two electrons in the same orbital Because must be paired (due to Pauli’s Exclusion Principle), it is possible to calculate the number of orbitals and the number of electrons in each n shell. • The number of orbitals per n level is given The by n2 (see table at end of chapter 7.) by (see • The maximum number of electrons per n The level is 2n2 (two electrons per orbital.) level – The value is 2n2 because of the two paired electrons per orbital. paired #orbitals ml n shell l subshell s 0 1K0 1 s 0 2L0 1 –1,0,1 1p 3 0 3 M0 s 1 –1,0,1 1p 3 2d 5 -2,-1,0,1,2 0 4N0 s 1 –1,0,1 1p 3 2d 5 -2,-1,0,1,2 3 f -3,-2,-1,0,1,2,3 7 Max n2 #e– 1 2 28 4 6 2 6 18 9 10 2 6 16 10 32 14 Atomic Subshell Energies and Electron Assignments Electron • The principle that describes how the The periodic chart is a function of electronic configurations is the Aufbau Principle. configurations • The electron that distinguishes an element The from the previous element enters the lowest energy atomic orbital available. lowest Penetrating and Shielding Penetrating • the radial distribution function the shows that the 2s orbital shows penetrates more deeply into the 1s orbital than does the 2p • the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more the repulsive force, they are more shielded from the attractive force of the nucleus of • the deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively are • the result is that the electrons in the the 2s sublevel are lower in (more the more negative) energy than the negative energy electrons in the 2p electrons Atomic Subshell Energies and Electron Assignments Assignments The Aufbau Principle describes the electron describes filling order in atoms. This filling is product of the effective is nuclear charge, Z*. For the same n, Z* is higher For for s orbital: s > p > d > f for Then, e− in s is the most Then, attracted by nucleus and has attracted the lowest energy Atomic Subshell Energies and Electron Assignments Assignments One mnemonic One to remember the to correct filling correct order for electrons in atoms is the increasing (n + ) value Atomic Subshell Energies and Electron Assignments Assignments or we can use this periodic chart or Atomic Electron Configurations Atomic • Now we will use the Aufbau Principle to Now determine the electronic configurations of the elements on the periodic chart. the • 1st row elements 1s ↑ 1 H 2 He ↑↓ Configuration 1 1s 1s 2 Atomic Electron Configurations Atomic Hund’s rule tells us that the electrons will fill the p and d orbitals by placing electrons in each orbital singly and with same spin until half-filled. That is the rule of maximum spin. Then the electrons will pair to finish the p orbitals. 1s 3 Li 2s 2p ↑↓ ↑ ↑↓ ↑ ↓ 4 Be Configuration 1s 2 2s1 1s 2 2s 2 5B ↑↓ ↑ ↓ ↑ 1s 2 2s 2 2p1 6C ↑↓ ↑ ↓ ↑ ↑ 1s 2 2s 2 2p 2 7N ↑↓ ↑ ↓ ↑ ↑ ↑ 1s 2 2s 2 2p 3 8O ↑↓ ↑ ↓ ↑↓ ↑ ↑ 1s 2 2s 2 2p 4 9F ↑↓ ↑ ↓ ↑↓ ↑↓ ↑ 1s 2 2s 2 2p 5 10 Ne ↑↓ ↑ ↓ ↑↓ ↑↓ ↑↓ 1s 2 2s 2 2p 6 Electrons in orbitals of or same kind, such as p or d orbitals, in the same shell (n), have the same energy; the are said to be degenerate. Atomic Electron Configurations Atomic 3rd row elements… 3s 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar [ Ne ] [ Ne ] [ Ne] [ Ne] [ Ne ] [ Ne] [ Ne ] [ Ne ] 3p Configuration ↑ ↑↓ ↑↓ ↑ ↑↓ ↑↑ ↑↓ ↑↑↑ ↑↓ ↑↓ ↑ ↑ ↑↓ ↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑↓ [ Ne] 3s1 [ Ne] 3s2 [ Ne] 3s2 3p1 [ Ne] 3s2 3p2 [ Ne] 3s2 3p3 [ Ne] 3s2 3p4 [ Ne] 3s2 3p5 [ Ne] 3s2 3p6 Atomic Electron Configurations Atomic 4th row elements… 3d 19 K [ Ar ] 4s ↑ 4p Configuration [ Ar ] 4s1 Atomic Electron Configurations Atomic 4th row elements… 3d 4s 19 K [ Ar ] ↑ 20 ↑↓ Ca [ Ar ] 4p Configuration [ Ar ] 4s1 [ Ar ] 4s2 Atomic Electron Configurations Atomic 4th row elements… The five d orbitals are degenerate 3d 4s 19 K [ Ar ] ↑ 20 ↑↓ 21 Ca [ Ar ] S c [ Ar ] ↑ ↑↓ 4p Configuration [ Ar ] 4s1 [ Ar ] 4s2 [ Ar ] 4s2 3d1 Atomic Electron Configurations Atomic 4th row elements… 3d 4s 19 K [ Ar ] ↑ 20 ↑↓ 21 Ca [ Ar ] Sc [ Ar ] ↑ Ti [ Ar ] ↑ ↑ 22 ↑↓ ↑↓ 4p Configuration [ Ar ] 4s1 [ Ar ] 4s2 [ Ar ] 4s2 3d1 [ Ar ] 4s2 3d 2 Atomic Electron Configurations Atomic 4th row elements… The five d orbitals are degenerate 3d 4s 19 K [ Ar ] ↑ 20 ↑↓ 21 Ca [ Ar ] Sc [ Ar ] ↑ Ti [ Ar ] ↑ ↑ ↑↓ 22 ↑↓ 23 ↑↓ V [ Ar ] ↑ ↑ ↑ 4p Configuration [ Ar ] 4s1 2 [ Ar ] 4s [ Ar ] 4s2 3d1 [ Ar ] 4s2 3d 2 [ Ar ] 4s2 3d 3 Atomic Electron Configurations Atomic 4th row elements… 3d 4s 19 K [ Ar ] ↑ 20 ↑↓ 21 Ca [ Ar ] Sc [ Ar ] ↑ Ti [ Ar ] ↑ ↑ ↑↓ 22 ↑↓ 23 ↑↓ V [ Ar ] ↑ ↑ ↑ 4p Configuration [ Ar ] 4s1 2 [ Ar ] 4s [ Ar ] 4s2 3d1 [ Ar ] 4s2 3d 2 [ Ar ] 4s2 3d 3 Atomic Electron Configurations Atomic 4th row elements… The [Ar] 4s1 3d5 configuration of Cr is more stable than [Ar] 4s2 3d4 (expected) 3d 4s 19 K [ Ar ] ↑ 20 ↑↓ 21 Ca [ Ar ] Sc [ Ar ] ↑ Ti [ Ar ] ↑ ↑ ↑↓ 22 ↑↓ 23 ↑↓ V [ Ar ] ↑ ↑ ↑ 24 Cr [ Ar ] ↑ ↑ ↑ ↑ ↑ ↑ 4p Configuration [ Ar] 4s1 [ Ar] 4s 2 [ Ar] 4s 2 3d1 [ Ar] 4s 2 3d 2 [ Ar] 4s 2 3d 3 [ Ar] 4s1 3d 5 There is an extra measure of stability associated with half - filled and completely filled orbitals. Atomic Electron Configurations Atomic 4th row elements… The [Ar] 4s1 3d10 configuration of Cu is more stable than [Ar] 4s2 3d9 (expected) 3d 25 Mn [ Ar ] ↑ ↑ ↑ ↑ ↑ 26 27 28 29 Fe [ Ar ] ↑↓ ↑ ↑ ↑ ↑ Co [ Ar ] ↑↓ ↑↓ ↑ ↑ ↑ Ni [ Ar ] ↑↓ ↑↓ ↑↓ ↑ ↑ 4s 4p ↑↓ ↑↓ ↑↓ ↑↓ Cu [ Ar ] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Another exception like Cr and for essentially the same reason. Configuration [ Ar ] 4s2 3d5 [ Ar ] 4s2 3d 6 2 7 [ Ar ] 4s 3d [ Ar ] 4s2 3d8 [ Ar ] 4s1 3d10 Atomic Electron Configurations Atomic 4th row elements… 3d 25 Mn [ Ar ] ↑ ↑ ↑ ↑ ↑ 26 27 28 29 30 Fe [ Ar ] ↑ ↑ ↑ ↑ ↑ ↓ Co [ Ar ] ↑ ↑ ↑ ↑ ↑ ↓↓ Ni [ Ar ] ↑ ↑ ↑ ↑ ↑ ↓↓↓ 4s ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ Cu [ A r ] ↑ ↑ ↑ ↑ ↑ ↑ ↓↓↓↓↓ Zn [ Ar ] ↑ ↑ ↑ ↑ ↑ ↑ ↓↓↓↓↓ ↓ 4p Configuration [ Ar ] 4s2 3d5 [ Ar ] 4s2 3d 6 [ Ar ] 4s2 3d 7 [ Ar ] 4s2 3d8 [ Ar ] 4s1 3d10 [ Ar ] 4s2 3d10 Atomic Electron Configurations Atomic 4th row elements… (remember Hund’s rule): ↑ ↑ __ is better (lower energy) than ↑↓ __ __ 4p 4p 3d 4s 4p 31 Ga [ Ar ] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ 32 33 34 35 36 Ge [ Ar ] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ As [ Ar ] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Se [ Ar ] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Br [ Ar ] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Kr [ Ar ] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Configuration [ Ar ] 4s2 3d10 4p1 [ Ar ] 4s2 3d10 4p2 ↑↑ [ Ar ] 4s2 3d10 4p3 ↑↑↑ [ Ar ] 4s2 3d10 4p4 ↑↓ ↑ ↑ 2 10 5 ↑↓ ↑↓ ↑ [ Ar ] 4s 3d 4p 2 10 6 ↑↓ ↑↓ ↑↓ [ Ar ] 4s 3d 4p Atomic Electron Configurations Atomic Lanthanides (4f) Ba [Xe] 6s2 56 La [Xe] 5d1 6s2 57 Ce [Xe] 4f1 5d1 6s2 58 Pr [Xe] 4f3 6s2 59 Praseodymium Yb [Xe] 4f14 6s2 70 Ytterbium Lu [Xe] 4f14 5d1 6s2 Lutetium Periodic Table Periodic s, p, d, and f-block in the Periodic Table s, P 1A 1 1 H 2A 23 4 3 4 5 6 7 Li Be 11 12 Na Mg 19 20 K Ca 37 38 Rb Sr 55 56 Cs Ba 87 88 Fr Ra (P–1)d 3B 4B 5B 6B 7B 8B 8B 8B 21 22 23 24 25 26 27 28 Sc Ti V Cr Mn Fe Co Ni 39 40 41 42 43 44 45 46 Y Zr Nb Mo Tc Ru Rh Pd 57 72 73 74 75 76 77 78 La Hf Ta W Re Os Ir Pt 89 104 105 106 107 108 109 Ac Rf Db Sg Bh Hs Mt (P)s (P–2)f 58 59 Ce Pr 90 91 Th Pa 1B 2B 29 30 Cu Zn 47 48 Ag Cd 79 80 Au Hg 60 61 62 63 64 65 66 Nd Pm Sm Eu Gd Tb Dy 92 93 94 95 96 97 98 U Np Pu Am Cm Bk Cf 3A 5 B 13 Al 31 Ga Ga 49 In 81 Tl 4A 5A 6A 7A 6789 CNOF 14 15 16 17 Si P S Cl 32 33 34 35 Ge As Se Br 50 51 52 53 Sn Sb Te I 82 83 84 85 Pb Bi Po At (P)p 67 68 69 70 71 Ho Er Tm Yb Lu 99 100 101 102 103 Es Fm Md No Lr 8A 2 He 10 Ne 18 Ar 36 Kr 54 Xe 86 Rn Valence Electrons Valence electrons in shell with highest n, i.e., the outermost electrons i.e., electrons, those beyond the core electrons electrons 1s2 2s2 2p6 3s1 2p 3s 1s2 2s2 2p6 3s2 3p2 2p 3s 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 2p 4s 1s2 2s2 2s 1s2 2s2 2p6 3s2 3p6 4s2 3d7 1s 2p 3s They determine the chemical properties of an They element. For the representative elements, they are the ns and np electrons; for transition and for elements they are the ns and (n−1)d electrons. ns 1A P 1A 1 1H 1s1 23 2s1 3 3s1 4 5 6 7 Li 11 Na 19 K 37 Rb 55 Cs 87 Fr 4s1 5s1 6s1 7s1 # of valence electrons = 1 2A 4 Be 12 Mg 20 Ca 38 Sr 56 Ba Ba 88 Ra 2s2 3s2 4s 2 5s2 6s2 7s2 # of valence electrons = 2 3A 5 B 13 Al 31 Ga Ga 49 In 81 Tl 2s2 2p1 3s2 3p1 4s 4p 2 1 5s2 5p1 6s2 6p1 # of valence electrons = 3 7A 9 F 17 Cl 35 Br Br 53 I 85 At # of valence electrons = 7 2s2 2p5 3s2 3p5 4s2 4p5 5s2 5p5 6s2 6p5 For the representative elements, the # of valence electrons = # of group The element X has the valence shell electron configuration, ns2 np4. electron X belongs to what group? belongs 1A 8A chalcogens 2 1 H 2A 34 Li Be 11 12 Na Mg 3B 4B 5B 6B 7B 8B 8B 8B 19 20 21 22 23 24 25 26 27 28 K Ca Sc Ti V Cr Mn Fe Co Ni 37 38 39 40 41 42 43 44 45 46 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd 55 56 57 72 73 74 75 76 77 78 Cs Ba La Hf Ta W Re Os Ir Pt 87 88 89 104 105 106 107 108 109 Fr Ra Ac Unq Unp Unh Uns Uno Une 3A 5 B 13 1B 2B Al 29 30 31 Cu Zn Ga 47 48 49 Ag Cd In 79 80 81 Au Hg Tl 4A 5A 6A 7A 6789 CNOF 14 15 16 17 Si P S Cl 32 33 34 35 Ge As Se Br 50 51 52 53 Sn Sb Te I 82 83 84 85 Pb Bi Po At He 10 Ne 18 Ar 36 Kr 54 Xe 86 Rn Energy (Orbital) Diagram Energy 4p 3d 4s E 3s 3p 2p 2s Be 1s 1s2 2s2 Orbital Box Diagrams Be 1s 1s 2s 2p 3s Orbital Box Diagrams N 1s 1s 2s 2p Formation of Cations Formation electrons lost from subshell with highest electrons n and l first (from valence electrons) examples K K+ 1s2 2s2 2p6 3s2 3p6 4s1 [Ar] 4s1 [Ar] 4s 1s2 2s2 2p6 3s2 3p6 [Ar] Ca Ca Ca2+ Al Al Al3+ In In In3+ In 1s2 2s2 2p6 3s2 3p6 4s2 [Ar] 4s2 [Ar] 4s 1s2 2s2 2p6 3s2 3p6 [Ar] 1s2 2s2 2p6 3s2 3p1 1s 3s [Ne] [Ne] [Kr] 4d10 5s2 5p1 5s [Kr] 4d10 Transition Metal Cations Transition In the process of ionization transition metals In the ns electrons are lost before the (n-1)d Fe: [Ar] 3d6 4s2 → Fe2+: [Ar] 3d6 Fe: Fe2+: [Ar] 3d6 → Fe3+: [Ar] 3d5 Cu: [Ar] 3d10 4s1 → Cu+: [Ar] 3d10 Cu+: [Ar] 3d10 → Cu2+: [Ar] 3d9 Fe, Fe2+, Fe3+, Cu, and Cu2+ are paramagnetic Two problems of ions, charge, and electron configuration Two An anion has a 3− charge and electron configuration 1s2 2s2 2p6 3s2 3p6. What is the symbol of the ion? The neutral atom has gained 3e- to form the ion, then The the neutral atom had 15 e-. In the neutral atom the # ethe = # p+ = Atomic number, that is 15. The element is, p+ then, phosphorus (phosphorus). Symbol of ion is P3−. A cation has a 2+ charge and its electron charge configuration is [Ar] 3d7. What is the symbol of the ion? configuration Here, the neutral atom has lost 2e-. It is a transition metal, due to the 3d electrons. Remember they firstly metal, lose e-s in 4s orbital. Symbol of ion is Co2+. Neutral atom has 18 + 7 + 2 = 27 e- = 27 p+ = atomic # [Ar] 3d7 lost 3d Atomic Properties and Periodic Trends Trends Periodic Properties of Periodic the Elements the 1. Atomic Radii 2. Ionization Energy 3. Electron Affinity 4. Ionic Radii Atomic Properties and Periodic Trends Atomic • Establish a classification scheme of the elements Establish based on their electron configurations. based • Noble Gases – All of them have completely filled electron All shells. They are not very reactive. shells. • Since they have similar electronic structures, Since their chemical reactions are similar. their – He 1s2 – Ne [He] 2s2 2p6 – Ar [Ne] 3s2 3p6 – Kr [Ar] 4s2 4p6 Kr [Ar] – Xe [Kr] 5s2 5p6 – Rn [Xe] 6s2 6p6 Atomic Properties and Periodic Trends Atomic Representative Elements are Representative are the elements in A groups on periodic chart. These elements will have These their “last” electron in an their outer s or p orbital. These elements have fairly regular variations in their regular properties. properties. Metallic character, for expl, Metallic increases from right to left and top to bottom. Atomic Properties and Periodic Trends Atomic • • • • d-Transition Elements Elements on periodic Elements chart in B groups. chart Sometimes called Sometimes transition metals. transition Each metal has d Each electrons. electrons. nsx (n-1)dy configurations These elements make the These transition from metals to nonmetals. nonmetals. Exhibit smaller variations Exhibit from row-to-row than the representative elements. representative Atomic Properties and Periodic Trends Atomic • • • • f - transition metals Sometimes called inner Sometimes transition metals. transition Electrons are being Electrons added to f orbitals. Electrons are being Electrons added two shells below the valence shell! the Consequently, very Consequently, slight variations of properties from one element to another. element Atomic Properties and Periodic Trends Atomic Outermost electrons (valence electrons) have the greatest Influence on the chemical have properties of elements. Atomic Properties and Periodic Trends Atomic Atomic radii describe the relative sizes of atoms. relative Atomic radii increase within a column going from the top to column the bottom of the periodic table. the The outermost electrons are The assigned to orbitals with increasingly higher values of n. The underlying electrons require some space, so the electrons of the outer shells must be further from the nucleus. nucleus. Atomic Properties and Periodic Trends Atomic Atomic radii decrease within a row going from within Left to right on the periodic table. This last fact seems This contrary to intuition. contrary How does nature make How the elements smaller the even though the electron even number is increasing? number Atomic Radii Atomic • The reason the atomic radii decrease across a The period is due to shielding or screening effect. shielding screening – Effective nuclear charge, Zeff, experienced by experienced an electron is less than the actual nuclear charge, Z. charge, – The inner electrons block the nuclear charge’s The effect on the outer electrons. effect • Moving across a period, each element has an Moving increased nuclear charge and the electrons are going into the same shell (2s and 2p or 3s and 3p, etc.). etc.). – Consequently, the outer electrons feel a Consequently, stronger effective nuclear charge. stronger – For Li, Zeff ~ +1 — For B, Zeff ~ +3 – For Be, Zeff ~ +2 For Atomic Radii Atomic • Example: Arrange these elements based on Example: their increasing atomic radii. increasing – Se, S, O, Te O < S < Se < Te Se In the same group atomic size increases In as n (and Z) increases as ─ Br, Ca, Ge, F Br, F < Br < Ge < Ca Br same group same period same Ionization Energy Ionization • First ionization energy (IE1) – The minimum amount of energy required to The remove the most loosely bound electron from an isolated gaseous atom to form a 1+ ion. an • Symbolically: Atom(g) + energy → ion+(g) + e­ Endothermic Mg(g) + 738kJ/mol → Mg+ + e­ IE1= 738kJ/mol IE1= Ionization Energy Ionization • Second ionization energy (IE2) – The amount of energy required to remove the The second electron from a gaseous 1+ ion. second • Symbolically: – ion+ + energy → ion2+ + eion Mg+ + 1451 kJ/mol → Mg2+ + eIE2= 1451 kJ/mol Atoms can have 3rd (IE3), 4th (IE4), etc. ionization energies. The values are consecutively getting larger. Ionization Energy Ionization Periodic trends for Ionization Periodic Energy: Energy 1) IE2 > IE1 It always takes more energy It to remove a second electron from an ion than from a neutral atom. 2) IE1 generally increases 2) moving from IA elements to VIIIA elements. Important exceptions at Be & Important B, N & O, etc. due to s and p B, and half-filled subshells. 3) IE1 generally decreases 3) moving down a family. IE1 for Li > IE1 for Na, etc IE for First Ionization Energies of Some Elements of 2500 2000 Ionization Energy (kJ/mol) 1500 1000 500 0 12 34 567 8 9 10 11 12 13 14 15 16 17 18 19 20 Atomic Number Ionization Energy Ionization • Example: Arrange these elements based on Example: their (increasing) first ionization energies. their – Sr, Be, Ca, Mg Sr < Ca < Mg < Be Sr – Al, Cl, Na, P Na < Al < P < Cl Na – O, Ga, Sr, Se Sr < Ga < Se < O Sr Ionization Energy Ionization • The reason Na forms Na+ and not Na2+ is that the energy difference between IE1 and that IE2 is so large. IE – Requires more than 9 times more energy Requires to remove the second electron than the first one. first • The same trend is persistent throughout The the series. the – Thus Mg forms Mg2+ and not Mg3+. – Al forms Al3+ and not Al4+. and H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K 1312 1312 2371 520 900 800 1086 1402 1314 1681 2080 496 738 577 786 1012 1000 1255 1520 419 Ionization Energies (kJ/mole) Ionization 5247 7297 1757 2430 2352 2857 3391 3375 3963 4565 1450 1816 1577 1896 2260 2297 2665 3069 11810 14840 3659 4619 4577 5301 6045 6276 6912 7732 2744 3229 2910 3380 3850 3947 4600 21000 25020 6221 7473 7468 8418 9376 9540 10550 11580 4356 4954 4565 5146 5770 5879 32810 37800 47300 9443 53250 64340 10980 13320 71300 84850 11020 15160 17860 92000 12190 15230 13360 16610 20110 25490 13620 18000 21700 25660 15030 18370 23290 27460 16080 19790 23780 29250 6272 21270 25410 29840 6996 8490 28080 31720 6544 9330 11020 33600 7240 8810 11970 13840 7971 9619 11380 14950 Electron Affinity (EA) Electron • Electron affinity is the amount of energy Electron absorbed or emitted when an electron is absorbed added to an isolated gaseous atom to form an ion with a 1- charge. an • Sign conventions for electron affinity. – If EA > 0 energy is absorbed (difficult) – If EA < 0 energy is released (easy) • Electron affinity is a measure of an atom’s Electron ability to form negative ions. ability • Symbolically: atom(g) + e- → ion-(g) EA (kJ/mol) Electron Affinity Electron • General periodic trend for electron affinity is – the values become more negative from left the become to right across a period on the periodic chart (affinity for electron increases). – the values become more negative from the bottom to top up a group on the periodic chart. chart. −Noble gases have EA > 0 (full electron confg) • An element with a high ionization energy generally has a high affinity for an electron, i.e., EA is largely negative. That is the case for halogens (F, Cl, Br, I), O, and S. Electron Affinity Electron F (Z= 9) and Cl (Z = 17) have the most negative EA Noble gases, He (2), Ne (10), and Ar (18), EA > 0; also Be, Mg, N Electron Affinity Electron Two examples of electron affinity values: Mg(g) + e- + 231 kJ/mol → Mg-(g) EA = 231kJ/mol Br(g) + e- → Br-(g) + 323 kJ/mol EA = -323 kJ/mol Br has a larger affinity for e− than Mg. The greater the affinity an atom has for an e− , the more negative EA is, the smaller it is. Ionic Radii Ionic Cations (positive ions) are always smaller than their respective neutral atoms. When one or more electrons are removed, the attractive force of the protons is now exerted on less electrons. Element Atomic Atomic Radius (Å) Radius Ion Ionic Ionic Radius (Å) Radius Na 11 p+, 11e- Mg 12p+, 12 e- Al 13 p+, 13e- 1.86 1.60 1.43 Na+ Mg2+ 11 p+, 10e- 12 p+, 10 e1.16 0.85 Al3+ 13 p+, 10e0.68 Ionic Radii Ionic Anions (negative ions) are always larger than their neutral atoms. F 1s2 2s2 2p5 + e− → F− 1s2 2s2 2p6 same Z nine electrons Element Atomic Radius(Å) Ion Ionic Radius(Å) ten electrons N 7 p+, 7e- O F 0.75 0.73 0.72 N37 p+, 10ep+, 10e 1.71 O2F− 8 p+, 10e- 9 p+, 10ep+, 10e p+, 10e 1.26 1.19 Ionic Radii Ionic Cation (positive ions) radii decrease from left to right across a period. Increasing nuclear charge attracts the electrons and decreases the radius. Rb+ and Sr2+ are isoelectronic, same # of e-s Ion Sr2+ In3+ Z = 37 p+ Ionic Radii(Å) Rb+ Z = 38 p+ Z = 49 p+ 1.66 1.32 0.94 Ionic Radii Ionic Anion (negative ions) radii decrease from left to right across a period. Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius. For these isoelectronic anions… 10 e− and 7 p+ 8 p+ 9 p+ Ion Ion N3- O2- F− Ionic Radii(Å) 1.71 1.26 1.19 Ionic Radii Ionic Example: Arrange these ions in order of decreasing radius. Ga3+, K+, Ca2+ K+ > Ca2+ > Ga3+ Cl−, Se2−, Br−, S2− Se2− > Br− isoelectronic > S2− > Cl− isoelectronic, same # of electrons Se2−(34 p+) > Br−(35 p+); they have 36 e− each. S2−(16 p+) > Cl−(17 p+); they have 18 e− each. Br− > S2− because Br− is in the 4th period, S2− is in the 3rd. Ionic Radii of isoelectronic species Ionic Isoelectronic species have the same number of electrons. Here are some examples with the number of (protons) and + or − charges N3−(Z=7) > O2−(Z=8) > F−(Z=9) > Ne(Z=10) neutral > Na+(Z=11) > Mg2+(Z=12) > Al3+ (Z=13) all have 10e− The nuclear charge (+) increases from left to right, so does attraction force to electrons: r decreases. S2−(Z=16) > Cl− (Z=17) > Ar0 (Z=18) > K+ (Z=19) > Ca2+ (Z=20) > Sc3+ (Z=21) all of them have 18e− ...
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This note was uploaded on 11/07/2011 for the course CHM 2045 taught by Professor Geiger during the Fall '08 term at University of Central Florida.

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