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

2 When the Elements Were Discovered

3 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

4 Classification of the 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 configuration of the element? (b) How should the element be classified? (c) Is the element diamagnetic or paramagnetic? Page 333

6 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

Cations and Anions Of Representative Elements

8 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 Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration

9 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

10 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

11 Effective Nuclear Charge (Z eff ) increasing Z eff

12 Atomic Radii metallic radius covalent radius

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14 Trends in Atomic Radii

Example 8.2 Referring to a periodic table, arrange the following atoms in order of increasing atomic radius: P, Si, N. Page 338

16 Comparison of Atomic Radii with Ionic Radii

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

18 The Radii (in pm) of Ions of Familiar Elements

Example 8.3 For each of the following pairs, indicate which one of the two species is larger: (a)N 3− or F − (b)Mg 2+ or Ca 2+ (c)Fe 2+ or Fe 3+ Page 340

20 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

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22 Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell Variation of the First Ionization Energy with Atomic Number

23 General Trends in First Ionization Energies Increasing First Ionization Energy

Exceptions 1.Between Group 2A and 3A (ex: Be and B) Group 3A is lower b/c of the single electron in the outermost p subshell which is shielded 2.Between Group 5A and 6A (ex: N and O) Group 6A doubles up one electron, the proximity of two electrons in the same orbital results in greater electrostatic repulsion and lower energy 24

Example 8.4 (a)Which atom should have a smaller first ionization energy: oxygen or sulfur? (b)Which atom should have a higher second ionization energy: lithium or beryllium? Page 346

26 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 - F - (g) O (g) + e - O - (g)  H = -328 kJ/mol EA = +328 kJ/mol  H = -141 kJ/mol EA = +141 kJ/mol

27 Variation of Electron Affinity With Atomic Number (H – Ba)

Example 8.5 Why are the electron affinities of the alkaline earth metals, shown in Table 8.3, either negative or small positive values? Page 349

29 Diagonal Relationships on the Periodic Table

30 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 Low IE = very reactive Never found in the pure state

31 Group 1A Elements (ns 1, n  2)

32 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

33 Group 2A Elements (ns 2, n  2)

34 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)

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

36 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)

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

38 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)

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

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

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

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

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

44 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.

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46 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.

47 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

48 Properties of Oxides Across a Period basicacidic