Sheets Page 1 Lecture 7 Lecture 7: Periodic properties Read: BLB 2.5; 7.1–7.6 HW: BLB 2:37; 7:11,23,25,27,31,45a–c,47a,e,f,53,61,94 Sup 7:1–12 Know: • screening effects • periodic properties atomic and ion sizes isoelectronic series electron configurations of ions ionization energies electron affinities Exam 1: Monday, Feb 9 @ 6:30!!! start preparing now!! only non-text programmable calculators are allowed—no PDAs, blackberries, cell phones, etc. will be permitted. Bring: pencils, student ID and a calculator—Absolutely NO text-programmable calculators or wireless devices (will be checked) Form a study group, use the CRC, take advantage of SI (info on web), use the online resources, and work those problems—practice, practice, practice Bonus deadline for Skill check tests 3 & 4 is tomorrow Thursday, 1/29 @ midnight; Bonus deadline for Skill check test 5 is Thursday, 2/5 & Skill check test 6 is SUNDAY, 2/8 Sheets’s office hours: Mondays 12:30-2; Tuesdays 10:30-12 in 324 Chem (or 326 Chem).Sheets Page 2 Lecture 7 Key to understanding periodic trends As n ↑, atomic orbitals become & less stable As Z (# protons) ↑, any given orbital becomes & more stable; Zeff ↑ Trade-off: Z ••• e– attraction e– ••• e– repulsion recall Coulomb’s Law: E !Q1Q2d column = “group” or “family”; row = “period”Sheets Page 3 Lecture 7 Electron configuration & the periodic table • electron configuration determines 1. periodic table organization (Chap 6) 2. properties of the elements atomic size ionization energy electron affinities reactivity • properties of elements determined by (n) & () of orbitals and atomic z (# ) (nuclear charge) • valence shell configuration is key to elementʼs properties ⇒ valence e– determine the chemistry of an element! • core electrons: smaller radius, lower E, mostly not participating in chemistry of reactionsSheets Page 4 Lecture 7 Atomic size Screening, again (see p 3 ff of Lecture 5 notes) • Zeff = Z – S • weak screening by e– with same n S is average number of e– between nucleus & e– of interest; often (but not always) use number of core e–Sheets Page 5 Lecture 7 Ion sizes What is happening when we make a cation or anion?? • cations are than parent atoms; why?? • anions are than parent atoms; why?? • atom sizes going down family; ion sizes alsoSheets Page 6 Lecture 7 Isoelectronic series • isoelectronic: same number of e–, same e– configuration • example 10 e– each: [He]2s22p6 Z 8 9 10 11 12 13 O2– F– Ne Na+ Mg2+ Al3+ Zeff size For this example, Zeff = Z – 2 from [He]Sheets Page 7 Lecture 7 Makinʼ ions • to form a cation: remove e– first from orbitals with highest n; these are the valence e– • in transition metals (TM): s electrons are part of valence e– Ag [Kr] 5s14d10 Ag+ [Kr] 4d10 • when forming TM ions: remove e– first then maybe remove e–, if needed ions with different charges may be formed Fe [Ar] 4s23d6 Fe2+ [Ar] 3d6 Fe3+ [Ar] 3d5 • to form an anion: add e– to empty or partially filled orbital with available nSheets Page 8 Lecture 7 Ionization energy (IE) • energy to remove 1 e– in the gas phase • I1: first ionization energy (IE) M(g) → M+(g) + e– ΔE = I1 • I2: second ionization energy M+(g) → M2+(g) + e– ΔE = I2 • In > 0; therefore energy is endothermic ann = ∞Sheets Page 9 Lecture 7 Ionization energy (IE) • I1 (kJ/mol) Na Mg Al Si P S Cl Ar 495 738 578 786 1012 1000 1251 1521 • e– closer to nucleus, more difficult to remove • as atomic size , IE ↑ • therefore, IE left to right • exceptions: extra energy to remove e– from filled subshells (Mg, Ar) or from half-filled subshell (P)Sheets Page 10 Lecture 7 Electron affinity (EA) • energy to add 1 e– in the gas phase M(g) + e– → M–(g) ΔE = EA • EA: either endothermic (ΔE +, energetically unfavorable) or exothermic (ΔE –, energetically favorable) • EA of cation is always exothermic (ΔE –); because opposite process of IE M+(g) + e– → M(g) EA(M+) = – I1(M) • complex property due to trade-off between Z+ ••• e– attraction e– ••• e– repulsionSheets Page 11 Lecture 7 Electron affinity trends • most negative (energetically favorable) for halogens (group 7A) why?? • group 2A (Be, Mg, Ca) do not want to fill a new subshell have positive values (energetically unfavorable) for EA (unstable negative ions) • group 1A (Li, Na, K) negative ions are not stable but have ns2 configuration • noble gases (group 8A) have positive values for EA why??Sheets Page 12 Lecture 7 Review of periodic trends • atomic size • ionization energy • electron affinity: both + and –; group 7A most negative • metallic character: metals lose e– (oxidation) nonmetals gain e– (reduction)Sheets Page 13 Lecture 7 The modern periodic table • elements listed in order of increasing atomic number (Z) • elements in same column (group or family) have similar chemical properties • metals toward left side of PT reactivity ↑ going down group • most metallic character, lower left of PT • nonmetals toward right side of PT reactivity ↑ upward • most nonmetallic character, upper right of PT • noble gases at far right of PT: quite inertSheets Page 14 Lecture 7 Metals • chemistry of metals: form group 1A (alkali): 1+ cations on atom, valence e–: ns1 most active (Cs, Fr) group 2A (alkaline earth): 2+ cations on atom, valence e–: ns2 less reactive than group 1A transition: 1+, 2+, 3+ cations • reactivity ↑ as IE ↓; easier to lose that e– as n ↑Sheets Page 15 Lecture 7 Nonmetals • group 8A (noble gases) on atom, valence e–: ns2np6 almost completely inert (except XeFn) all gases • chemistry of nonmetals: form group 7A (halogens): 1– anions on atom, valence e–: ns2np5 most reactive nonmetals, particularly F2 reactions dominated by: X2 + 2e– ← 2X– • reactivity of halogens ↑ as EA ↑ group 6A (oxygen family): 2– anions on atom, valence e–: ns2np4Sheets Page 16 Lecture 7 Nonmetals
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