Periodic Relationships Among the Elements Chapter 8 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

The periodic table places the most reactive metals in Group 1A and the most reactive nonmetals in Group 7A. When these elements are mixed, we predict a vigorous reaction to ensue, as evidenced by the formation of sodium chloride. M any of the chemical properties of the elements can be understood in terms of their electron configurations. Because electrons fill atomic orbitals in a fairly regular fashion, it is not surprising that elements with similar electron configurations, such as sodium and potassium, behave similarly in many respects. Chemists in the nineteenth century recognized periodic trends in the physical and chemical properties of the elements, long before quantum theory came onto the scene. Although these chemists were not aware of the existence of electrons and protons, their efforts to systematize the chemistry of the elements were remarkably successful. Their main sources of information were the atomic masses of the elements and other known physical and chemical properties. 8.2 Periodic Classifi cation of the Elements Figure 8.2 shows the periodic table together with the outermost ground-state electron confi gurations of the elements. According to the type of subshell being filled, the elements can be divided into categories—the representative elements, the noble gases, the transition elements (or transition metals), the lanthanides, and the actinides. The representative elements (also called main group elements ) are the elements in Groups 1A through 7A, all of which have incompletely fi lled s or p subshells of the highest principal quantum number. With the exception of helium, the noble gases (the Group 8A elements) all have a completely fi lled p subshell. (The electron confi gurations are 1 s 2 for helium and ns 2 np 6 for the other noble gases, where n is the principal quantum number for the outermost shell.) The transition metals are the elements in Groups 1B and 3B through 8B, which have incompletely fi lled d subshells, or readily produce cations with incompletely 2

3 The Modern Periodic Table Period Group Alkali Metal Noble Gas Halogen Alkali Earth Metal

4 When the Elements Were Discovered

5 ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 ns 2 np 6 d1d1 d5d5 d 10 4f 5f Ground State Electron Configurations of the Elements

6 Classification of the Elements

7 Electron Configurations of Cations and Anions Na [Ne]3s 1 Na + [Ne] Ca [Ar]4s 2 Ca 2+ [Ar] Al [Ne]3s 2 3p 1 Al 3+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. H 1s 1 H - 1s 2 or [He] F 1s 2 2s 2 2p 5 F - 1s 2 2s 2 2p 6 or [Ne] O 1s 2 2s 2 2p 4 O 2- 1s 2 2s 2 2p 6 or [Ne] N 1s 2 2s 2 2p 3 N 3- 1s 2 2s 2 2p 6 or [Ne] Atoms gain electrons so that anion has a noble-gas outer electron configuration. Of Representative Elements

EXAMPLE 8.1 An atom of a certain element has 15 electrons. Without consulting a periodic table, answer the following questions: (a) What is the ground-state electron confi guration of the element? (b) How should the element be classifi ed? (c) Is the element diamagnetic or paramagnetic? Strategy (a) We refer to the building-up principle discussed in Section 7.9 and start writing the electron confi guration with principal quantum number n 5 1 and continuing upward until all the electrons are accounted for. (b) What are the electron confi guration characteristics of representative elements? transition elements? noble gases? (c) Examine the pairing scheme of the electrons in the outermost shell. What determines whether an element is diamagnetic or paramagnetic? Solution (a) We know that for n 5 1 we have a 1 s orbital (2 electrons); for n 5 2 we have a 2 s orbital (2 electrons) and three 2 p orbitals (6 electrons); for n 5 3 we have a 3 s orbital (2 electrons). The number of electrons left is and these three electrons are placed in the 3 p orbitals. The electron confi guration is 1 s 2 2 s 2 2 p 6 3 s 2 3 p 3. (b) Because the 3 p subshell is not completely fi lled, this is a representative element. Based on the information given, we cannot say whether it is a metal, a nonmetal, or a metalloid. (c) According to Hund’s rule, the three electrons in the 3 p orbitals have parallel spins (three unpaired electrons). Therefore, the element is paramagnetic. 8

Cations and Anions Of Representative Elements

10 Na + : [Ne]Al 3+ : [Ne] F - : 1s 2 2s 2 2p 6 or [Ne] O 2- : 1s 2 2s 2 2p 6 or [Ne]N 3- : 1s 2 2s 2 2p 6 or [Ne] Na +, Al 3+, F -, O 2-, and N 3- are all isoelectronic with Ne What neutral atom is isoelectronic with H - ? H - : 1s 2 same electron configuration as He Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration

11 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]4s 2 3d 6 Fe 2+ : [Ar]4s 0 3d 6 or [Ar]3d 6 Fe 3+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Mn: [Ar]4s 2 3d 5 Mn 2+ : [Ar]4s 0 3d 5 or [Ar]3d 5

12 Effective nuclear charge (Z eff ) is the “positive charge” felt by an electron. Na Mg Al Si Z eff Core Z Radius (pm) Z eff = Z -  0 <  < Z (  = shielding constant) Z eff  Z – number of inner or core electrons

13 Effective Nuclear Charge (Z eff ) increasing Z eff

14 Atomic Radii metallic radius covalent radius

15

16 Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.

17 Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. I 1 + X (g) X + (g) + e - I 2 + X + (g) X 2 + (g) + e - I 3 + X 2+ (g) X 3 + (g) + e - I 1 first ionization energy I 2 second ionization energy I 3 third ionization energy I 1 < I 2 < I 3

18 General Trends in First Ionization Energies Increasing First Ionization Energy

19 Electron affinity 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 - X - (g) O (g) + e - O - (g)  H = -328 kJ/mol EA = +328 kJ/mol  H = -141 kJ/mol EA = +141 kJ/mol

20 Group 1A Elements (ns 1, n  2) M M e - 2M (s) + 2H 2 O (l) 2MOH (aq) + H 2(g) 4M (s) + O 2(g) 2M 2 O (s) Increasing reactivity

21 Group 1A Elements (ns 1, n  2)

22 Group 2A Elements (ns 2, n  2) M M e - Be (s) + 2H 2 O (l) No Reaction Increasing reactivity Mg (s) + 2H 2 O (g) Mg(OH) 2(aq) + H 2(g) M (s) + 2H 2 O (l) M(OH) 2(aq) + H 2(g) M = Ca, Sr, or Ba

23 Group 2A Elements (ns 2, n  2)

24 Group 3A Elements (ns 2 np 1, n  2) 4Al (s) + 3O 2(g) 2Al 2 O 3(s) 2Al (s) + 6H + (aq) 2Al 3+ (aq) + 3H 2(g)

25 Group 3A Elements (ns 2 np 1, n  2)

26 Group 4A Elements (ns 2 np 2, n  2) Sn (s) + 2H + (aq) Sn 2+ (aq) + H 2 (g) Pb (s) + 2H + (aq) Pb 2+ (aq) + H 2 (g)

27 Group 4A Elements (ns 2 np 2, n  2)

28 Group 5A Elements (ns 2 np 3, n  2) N 2 O 5(s) + H 2 O (l) 2HNO 3(aq) P 4 O 10(s) + 6H 2 O (l) 4H 3 PO 4(aq)

29 Group 5A Elements (ns 2 np 3, n  2)

30 Group 6A Elements (ns 2 np 4, n  2) SO 3(g) + H 2 O (l) H 2 SO 4(aq)

31 Group 6A Elements (ns 2 np 4, n  2)

32 Group 7A Elements (ns 2 np 5, n  2) X + 1e - X - 1 X 2(g) + H 2(g) 2HX (g) Increasing reactivity

33 Group 7A Elements (ns 2 np 5, n  2)

34 Group 8A Elements (ns 2 np 6, n  2) Completely filled ns and np subshells. Highest ionization energy of all elements. No tendency to accept extra electrons.

35 Compounds of the Noble Gases A number of xenon compounds XeF 4, XeO 3, XeO 4, XeOF 4 exist. A few krypton compounds (KrF 2, for example) have been prepared.

36 The metals in these two groups have similar outer electron configurations, with one electron in the outermost s orbital. Chemical properties are quite different due to difference in the ionization energy. Comparison of Group 1A and 1B Lower I 1, more reactive

37 Properties of Oxides Across a Period basicacidic