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Home , Ionization energy

7.16 Ionic Radius

• Part of the distance between the centers of two neighboring in an ionic solid (O2- is used as a standard with radius 140 pm) • Cations are smaller than their parent • Anions are larger then their parent atoms • Ionic sizes of cations as well as anions follow the same trends in the as the sizes of atoms

Fig. 7.34

• Isoelectronic species – atoms and ions with the 7.17 same number of (have the same configuration) • Ionization energy (I) – energy required to remove – size decreases with increasing the of an electron from a gas-phase st - the element (nuclear charge increases) – first ionization energy (I1) - to remove the 1 e + - Example: Compare the sizes of Cl-, Ca2+ and Sc3+. X(g) ® X (g) + e – second ionization energy (I ) - to remove a 2nd e- isoelectronic, of [Ar] 2 X+(g) ® X2+(g) + e- Þ Sc3+ < Ca2+ < Cl- (atomic number ¯) • Ionization energies are positive (endothermic) and Example: Compare the sizes of Ca, Ca2+ and Mg2+. become larger with every subsequent ionization 2+ Ca < Ca (cation is smaller) 0 < I1 < I2 < I3 < I4 ... Mg2+ < Ca2+ (Mg is above Ca) Þ Mg2+ < Ca2+ < Ca – it’s harder to remove an e- from a positive

• First ionization decrease down a and increase from left to right across a (with some exceptions) – down a group electrons are removed from shells that are farther from the nucleus (less tightly bound)

– across a period Zeff increases (valence electrons are more tightly bound to the nucleus) • Considerable jump in the successive ionization energies occurs after removal of all valence electrons – core electrons are much more difficult to remove than valence electrons – explains the charges of stable metal cations Fig. 7.38

1 1 Na ® [Ne]3s • Irregularities in the ionization energy trends I1 = 496 kJ/mol, I2=4562 kJ/mol – decrease in I1 between groups 2(IIA) and 13(IIIA) Stable cation ® Na+ elements Mg ® [Ne]3s2 group IIA ® ns2 group IIIA ® ns2np1 • the np electron is easier to remove than the ns electron – I1 = 738 kJ/mol, I2 = 1450 kJ/mol, I3 = 7734 kJ/mol p-subshells have higher energy and are less tightly bound Stable cation ® Mg2+ – decrease in I1 between groups 15 and 16 elements • Low ionization energy accounts for the metallic group 15 ® ns2np 1np 1np 1 character of elements in the lower left corner of x y z group 16 ® ns2np 2np 1np 1 the table (s, d, f and some of the p block) – easy x y z - • it’s easier to remove the paired electron than the unpaired removal of e provides better conductivity and electrons on the p-orbitals – paired electrons repel each tendency to form cations other stronger than unpaired electrons

7.18 • Electron affinity (A) – energy released when an electron is added to a gas-phase atom st - – first electron affinity (A1) - to add the 1 e X(g) + e- ® X-(g) nd - – second electron affinity (A2) - to add a 2 e X-(g) + e- ® X2-(g) • Electron affinities can be either exothermic or endothermic – by convention assume: (+) for exothermic A and (-) for endothermic A Fig. 7.43

• First electron affinities tend to be larger in the • Irregularities in the electron affinity trends upper right corner of the table similarly to the first – decrease in A between groups 1 and 2 elements ionization energies 1 group 1 ® ns1 group 2 ® ns2 • Successive electron affinities are smaller and • for group 2 the new electron is added to the higher smaller – more endothermic (A1 > A2 > A3 > A4 ...) energy np subshell – it’s harder to add an e- to a negative ion – decrease in A1 between groups 14 and 15 elements 2 1 1 • Considerable drop in the successive electron group 14 ® ns npx npy affinities occurs after achieving a 2 1 1 1 group 15 ® ns npx npy npz configuration – the new electrons are added to a • for group 15 the new electron is added to an already higher principal shell occupied np orbital – pairing of electrons is – explains the charges of the stable anions of groups energetically unfavorable (stronger repulsion) 15, 16 and 17 (N3-, O2-, F-)

2 • s-block elements 7.19 Predicting Periodic Properties – reactive metals – form cations (low I) – reactivity increases down a group (I decreases) • p-block elements – upper right corner ® reactive – form anions (high A) – lower right corner ® less reactive metals • d-block elements – transitional properties between the s- and p-blocks – metals forming cation with variable charges • Diagonal relationships – diagonal (left-to-right) bands of elements with similar properties (similar atomic size, I and A

Fig. 7.51 values) – the form a diagonal band

Assignments:

• Homework: Chpt. 7/1, 3, 5, 7, 9, 11, 15, 21, 23, 27, 31, 33, 39, 41, 43, 45, 47, 51, 57, 61, 65, 67, 69, 73, 75, 97

• Student Companion: 7.2, 7.4, 7.5

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