Periodic Relationships Among the Elements Chapter 8

8.1 When the Elements Were Discovered

8.2 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

8.2 Classification of the Elements

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 8.2

Cations and Anions Of Representative Elements 8.2

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 8.2

Electron Configurations of Cations of Transition Metals 8.2 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

Effective nuclear charge (Z eff ) is the “positive charge” felt by an electron. Na Mg Al Si Z eff Core Z Radius Z eff = Z -  0 <  < Z (  = shielding constant) Z eff  Z – number of inner or core electrons 8.3

Effective Nuclear Charge (Z eff ) 8.3 increasing Z eff

8.3

Atomic Radii 8.3

Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed. 8.3

Chemistry in Action: The 3 rd Liquid Element? Liquid? 113 elements, 2 are liquids at 25 0 C – Br 2 and Hg 223 Fr, t 1/2 = 21 minutes

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 (g) X 3 + (g) + e - I 1 first ionization energy I 2 second ionization energy I 3 third ionization energy 8.4 I 1 < I 2 < I 3

8.4

Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell 8.4

General Trend in First Ionization Energies 8.4 Increasing First Ionization Energy

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) 8.5 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

8.5

8.6

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 8.6

Group 2A Elements (ns 2, n  2) M M e - Be (s) + 2H 2 O (l) No Reaction Increasing reactivity 8.6 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

Group 3A Elements (ns 2 np 1, n  2) 8.6 4Al (s) + 3O 2(g) 2Al 2 O 3(s) 2Al (s) + 6H + (aq) 2Al 3+ (aq) + 3H 2(g)

Group 4A Elements (ns 2 np 2, n  2) 8.6 Sn (s) + 2H + (aq) Sn 2+ (aq) + H 2 (g) Pb (s) + 2H + (aq) Pb 2+ (aq) + H 2 (g)

Group 5A Elements (ns 2 np 3, n  2) 8.6 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)

Group 6A Elements (ns 2 np 4, n  2) 8.6 SO 3(g) + H 2 O (l) H 2 SO 4(aq)

Group 7A Elements (ns 2 np 5, n  2) X + 1e - X - 1 X 2(g) + H 2(g) 2HX (g) Increasing reactivity 8.6

Group 8A Elements (ns 2 np 6, n  2) 8.6 Completely filled ns and np subshells. Highest ionization energy of all elements. No tendency to accept extra electrons.

Properties of Oxides Across a Period basicacidic 8.6

Chemistry in Action: Discovery of the Noble Gases Sir William Ramsay