Ionization Energy Ionization Energy

Periodic Trends in Ionization Energies

• The first ionization energy, I1, is the amount of energy • Ionization energy decreases down a group. required to remove an electron from a gaseous atom: • This means that the outermost electron is more readily Na(g) → Na+(g) + e-. removed as we go down a group. • The second ionization energy, I , is the energy required • As the atom gets bigger, it becomes easier to remove an 2 electron from the most spatially extended orbital. to remove an electron from a gaseous ion: • Ionization energy generally increases across a period. + 2+ - Na (g) → Na (g) + e . • As we move across a period, Zeff increases. Therefore, it • The larger ionization energy, the more difficult it is to becomes more difficult to remove an electron. remove the electron. • Two exceptions: removing the first p electron and removing the fourth p electron.

Ionization Energy Ionization Energy Variations in Successive Ionization Energies Periodic Trends in Ionization Energies • There is a sharp increase in ionization energy when a •The s electrons are more effective at shielding than p core electron is removed. electrons. Therefore, forming the s2p0 becomes more favorable.

• When a second electron is placed in a p orbital occupied by a single electron, the electron-electron repulsion increases. When this electron is removed, the resulting s2p3 is more stable than the starting s2p4 configuration. Therefore, there is a decrease in ionization energy. Prentice Hall © 2003 Chapter 7 Prentice Hall © 2003 Chapter 7

1 Ionization Energy

Electron Configuration of Ions • Cations: electrons removed from orbital with highest principle quantum number, n, first: Li (1s2 2s1) ⇒ Li+ (1s2) Fe ([Ar]3d6 4s2) ⇒ Fe3+ ([Ar]3d5) • Anions: electrons added to the orbital with highest n: F (1s2 2s2 2p5) ⇒ F− (1s2 2s2 2p6) • Lets try a problem.

Electron Affinities

• Electron affinity is the opposite of ionization energy. • Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion: Cl(g) + e- → Cl-(g) • Electron affinity can either be exothermic (as the above example) or endothermic: Ar(g) + e- → Ar-(g)

2 Electron Affinities Metals, Nonmetals, and Metalloids

• Look at electron configurations to determine whether electron affinity is positive or negative. • Should we try another problem?

Electron Affinities Metals, Nonmetals, and Metalloids Figure 7.11 Metals • Metallic character refers to the properties of metals (shiny or lustrous, malleable and ductile, oxides form basic ionic solids, and tend to form cations in aqueous solution). • Metallic character increases down a group. • Metallic character decreases across a period. • Metals have low ionization energies. • Most neutral metals are oxidized rather than reduced.

3 Metals, Nonmetals, and Metals, Nonmetals, and Metalloids Metalloids Metals Metals • When metals are oxidized they tend to form • Most metal oxides are basic: characteristics cations. Metal oxide + water → metal hydroxide • All group 1A metals form M+ ions. Na2O(s) + H2O(l) → 2NaOH(aq) Nonmetals 2+ • All group 2A metals form M ions. • Nonmetals are more diverse in their behavior than metals. • Most transition metals have variable charges. • When nonmetals react with metals, nonmetals tend to gain electrons: metal + nonmetal → salt

2Al(s) + 3Br2(l) → 2AlBr3(s)

Metals, Nonmetals, and Metals, Nonmetals, and Metalloids Metalloids Metals Nonmetals • Most nonmetal oxides are acidic: nonmetal oxide + water → acid

P4O10(s) + H2O(l) → 4H3PO4(aq) Metalloids • Metalloids have properties that are intermediate between metals and nonmetals. • Example: Si has a metallic luster but it is brittle. • Metalloids have found fame in the semiconductor Figure 7.14 industry. Prentice Hall © 2003 Chapter 7 Prentice Hall © 2003 Chapter 7