Periodic Relationships Among the Elements Chapter 8 Periodic Relationships Among the Elements

Electron Configurations of the Elements Valence Electrons • the outer electrons of an atom, the ones involved in Bonding • for the representative elements (main group elements), the number of valence electrons is equal to the A Group Number

Which of the following electron configurations do represent similar chemical properties of their atoms? 1s22s22p63s2 (ii) 1s22s22p3 (iii) 1s22s22p63s23p64s23d104p5 (iv) 1s22s1 (v) 1s22s22p6 (vi) 1s22s22p63s23p3 (ii), (vi)

Ground State Electron Configurations of the Elements ns2np6 Ground State Electron Configurations of the Elements ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 4f 5f

Electron Configurations of Cations and Anions Of Representative Elements Atoms lose electrons so that cation has a noble-gas outer electron configuration. Na [Ne]3s1 Na+ [Ne] Ca [Ar]4s2 Ca2+ [Ar] Remember: electrons are first removed from orbitals with the highest principal quantum number. Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Cations and Anions Of Representative Elements +1 +2 +3 -3 -2 -1

Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne Isoelectronic: have the same number of electrons What neutral atom is isoelectronic with H- ? H-: 1s2 same electron configuration as He

Which of the following pairs are not isoelectronic with Ar? K+ and Cl- (d) Ca2+ and S2- Al3+and N3- (e) P3- and Sc3+ Rh3+ and Ir3+ (a) K+ [Ar] Cl- [Ar] (b) Al3+ [Ne] N3- [Ne] NOT (c) Rh3+ [Rn]5s14d5 Ir3+ [Rn] 5s24d4 NOT (d) Ca2+ [Ar] S2- [Ar] (e) P3- [Ar] Sc3+ [Ar]

Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5

A metal ion with a net +3 charge has five electrons in the 3d subshell A metal ion with a net +3 charge has five electrons in the 3d subshell. What is this metal? (a) Cr (b) Mn (c) Fe (d) Co (e) Ni This species has +3 charges, which indicates that it has three more protons than the electrons. According to the question that it has five electrons in the 3d subshell, and thus the total electrons in valence shells for its atomic type will be 8. Note that the transition metals (with d electrons) losing its s electrons prior to its d electrons and thus the valence electron configuration will be 4s23d6, which is Fe.

Periodic Variation in Physical Properties

Atomic Radius

Sizes of atoms: Trends • size increases while going down a Group • Why? • Because orbital size increases with increasing n number, the size of atom increases. The highest energy electrons can be farther away from the nucleus. • size decreases going across Period • effective nuclear charge increases. The larger the effective nuclear charge, the stronger hold of the nucleus on electrons. The electron clouds shrink. • Remember n does NOT increases--e- cannot be farther from nucleus.

Ionic Radius Cation is always smaller than atom from which it is formed. This is because the nuclear charge remains the same but the reduced electron repulsion resulting from removal of electrons make the electron clouds shrink. Anion is always larger than atom from which it is formed. This is because the nuclear charge remains the same but electron repulsion resulting from the additional electron enlarges the electron clouds.

From top to bottom, both the atomic and ionic radius increase within group. Across a period the anions are usually larger than cations. For ions derived from different groups, size comparison is meaningful only if the ions are isoelectronic. E.g. Na+ (Z=11) is smaller than F- (Z=9). The larger effective charge results in a smaller radius. Radius of tripositive ions < dipositive ions<unipositive ions Al3+ <Mg2+ <Na+ Radius of uninegative ions < dinegative ions < trinegative ions N3- > O2->F-

Which of the following is a correct order of atomic radii? Ar < P < Na < Ca < Cs (b) Cs < Rb < K < Na < Li (c) Na < Mg < Al < Si < P (d) Rb < Mg < C < F < He (e) Li < H < Al < K < Ar (a)

Which of the following is a correct order of ionic radii? Na+ < Mg2+ < Al3+ < O2- < F- < N3- Al3+ < Mg2+ < Na+ < F- < O2- < N3- (c) F- < O2- < N3- < Na+ < Mg2+ < Al3+ (d) Al3+ < Mg2+ < Na+ < N3- < O2- < F- (e) Na+ < Mg2+ < Al3+ < O2- < N3- < F- (b)

Ionization energy (I.E.) is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. I1 + X (g) X+(g) + e- I1 first ionization energy, remove 1st electron I2 + X+(g) X2+(g) + e- I2 second ionization energy, remove 2nd electron I3 + X2+(g) X3+(g) + e- I3 third ionization energy, remove 3rd electron I1 < I2 < I3

General Trend in First Ionization Energies Increasing First Ionization Energy Increasing First Ionization Energy

First IE: Trends down a Group • remember that Cs is more reactive than Na • Easier to remove an e- from Cs than from Na • As size increases, I.E. decreases Trends across a Period • remember that Cl gains rather than loses an e- • Easier to remove an e- from Al than from Cl • As size decreases, I.E. increases

Exceptions: Ionization Energy Occur between Group 2A and 3A Group 3A: ns2np1. one single electron in p orbital. • removing the first p electron is less than expected •Why? This p electron is shielded by inner ns2 electrons. Less energy is needed to remove a single p electron than to remove a pair of s electron of the same n level. 2. Occur Group 5A and 6A 5A: ns2np3 6A: ns2np4 • removing the fourth p electron is less than expected •Why? 2nd electron in a p orbital increases the electron-electron repulsion, which makes it easier to ionize an atom of Group 6A.

Cs < Na < Al < S < Cl Arrange the following in order of the increasing first ionization energy: Na, Cl, Al, S, Cs Hint: Ionization energy increases across a row of the periodic table and decreases down a column or group. Cs < Na < Al < S < Cl

Electron affinity (EA) is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. X (g) + e- X-(g) F (g) + e- F-(g) DH = -328 kJ/mol EA = +328 kJ/mol O (g) + e- O-(g) DH = -141 kJ/mol EA = +141 kJ/mol

Variation of Electron Affinity With Atomic Number (H – Ba) Overall trend: increases from left to right across the period. The values varied little in the group. The EA of metals are generally lower than those of nonmetals.

Properties of Oxides Across a Period basic acidic Across a period, oxides change from basic to amphoteric to acidic. Going down a group, the oxides become more basic.