Chapter 7 Periodic Properties of the Elements
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1 Chapter 7 Periodic Properties of the Elements 7.1 Development of the Periodic Table In 1869 Dmitri Medeleev arranged the elements in order of increasing weight and place those with similar properties in the same vertical column. He had to leave some blank spaces and even predicted the properties of elements yet to be discovered. (Table 7.1, page 251) 7.2 Effective Nuclear Charge The outermost electron of Na does not feel the full nuclear charge. The effective nuclear charge is the average shielded charge felt by a particular electron. The effective nuclear charge, Z eff, is found this way: Z eff = Z S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons. In general: - Smaller orbitals (lower n) screen larger ones. - Orbitals with higher l are more effectively screened (by small orbitals) than those with smaller l. - Electrons in the same orbital or subshell do not screen each other effectively 1s e- s screen 2s e- s partially (they spend some time closer to the nucleus). 1s e- s screen 2p e- s better (since they spend more time outside). 2s orbital has lower E than 2p. 7.3 Sizes of Atoms and Ions What is the size of an atom? We do not have a unique and universally accepted answer to this question yet. (Why not?) The atomic radius generally depends on the size of the outermost s orbital. Problem 1 Based on the following data for Cu, estimate the volume occupied by a single atom of the element. ρ = 8.93 g/cm 3 ; MM Cu = g/mol Atom Size The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei (vs. the nonbonding radius). 1
2 Periodic Trends in Atomic Radii 1. Radius increases as we go down a group (why?) Li 1.52 Å 2s Na 1.86 Å 3s K 2.27 Å 4s Rb 2.48 Å 5s 2. Radius decreases as you go from left to right across a row (why?). Na Mg Al 1.86 Å 1.60 Å 1.43 Å (3s orbital, Z eff) Periodic Trends in Ionic Radii 1a. Positive ions are smaller than the neutral atom. 1b. Negative ions are larger than the neutral atom. Na Na + O O Å 1.16 Å 0.73 Å 1.26 Å (Why?) 2. For Ions of the same charge, their size increases as we go down a group. (Why?) Sizes of Ions An isoelectronic series is formed by ions that posses the same number of electrons. Ionic size decreases with an increasing nuclear charge. Problem 2 Which one is larger? Why? a. Ca or C 2+? b. K +, Ca 2+ or Cl -? 7.4 Ionization Energy (IE) IE is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. X (g) X + (g) + e - ΔE = IE Since energy is required for the process, IE (+). If IE is high, then it s difficult to remove an e -. If IE is low, it s not as difficult. In general, larger atoms have lower IE. Successive Ionization Energies It requires more energy to remove each successive electron (we talk about the first, second, third, etc., ionization energies). When all valence electrons have been removed, the ionization energy takes a quantum leap. (See Table 7.2) Periodic Trends in First Ionization Energies As one goes down a column, less energy is required to remove the first electron. (Why?) Generally, as one goes across a row, it gets harder to remove an electron. (Why?) However, there are two apparent discontinuities in this trend. The first occurs between Groups IIA and IIIA. IIA Mg 1s 2 2s 2 2p 6 3s 2 IIIA Al 1s 2 2s 2 2p 6 3s 2 3p 1 The e- is removed from a p orbital, further from the nucleus. There is also a small amount of repulsion by the s electrons. 2
3 The second occurs between Groups VA and VIA. VA P 1s 2 2s 2 2p 6 3s 2 3p 3 VIA S 1s 2 2s 2 2p 6 3s 2 3p 4 The electron comes from a double occupied orbital so there is more repulsion; plus, half-filled p orbitals are more stable. Problem 3 Explain the following anomaly: N IE = 1402 kj; O IE = 1314 kj This is opposite from prediction. Why? Electron Configurations of Ions Main-group ions have noble gas e- configurations. Cl - : 17p +, 18e - 1s 2 2s 2 2p 6 3s 2 3p 6 Al 3+ : 13p +,10e - 1s 2 2s 2 2p 6 Transition metal ions lose e - s from the outermost s orbital first. Fe: [Ar]4s 2 3d 6 Fe 2+ : [Ar]3d 6 Fe 3+ : [Ar]3d 5 Problem 4 Write the electron configuration for: a. Ca 2+ b. Cr Electron Affinities Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e Cl ΔE = EA Most EA values are negative (release energy). Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. Exceptions to the EA Trend 1. For noble gases EA is positive (unfavorable) since energy needs to be added to insert an electron in a higher E orbital. 2. Groups IA and IIA. The added e - must go in a p-orbital, not an s-orbital. The e - is farther from nucleus and feels repulsion from the s e - s. 3. Groups IVA and VA. Group VA has no empty orbitals. The extra electron must go into an already occupied orbital, creating repulsion. Problem 5 1. Which element should have the more negative EA, Al or Si? 2. Which atom gains an e - more readily, Al or Si? 3. Which is easier to remove an e - from: Na or P? Li or Rb? 4. Which is smallest, Se or Cl? 5. Rank in order of increasing size: O 2-, Mg 2+, Na +. 3
4 7.6 Metals, Nonmetals, and Metaloids Metals: Shiny, malleable, ductile Good conductors of heat and electricity High melting points (most of them) Low IE s (easy to form +ions, lose e - s, get oxidized) Metal + Nonmetal ionic compounds. Metal Oxides are basic: Metal oxide + water metal hydroxide Na 2O (s) + H 2O (l) 2NaOH (aq) CaO (s) + H 2O (l) Ca(OH) 2(aq) Metal oxide + acid salt + water NiO (s) + 2HCl (aq) NiCl 2(aq) + H 2O (l) Nonmetals Not shiny Poor conductors of heat and electricity Low melting points (in general) Tend to gain e - s (favorable EA form ions, reduced) Compounds with only nonmetals are molecular Nonmetal oxides are acidic: Nonmetal oxide + water acid CO 2(g) + H 2O (l) H 2CO 3(aq) P 4O 10(s) + 6 H 2O (l) 4H 3PO 4(aq) Nonmetal oxide + Base salt + water CO 2(g) + 2NaOH (aq) Na 2CO 3(aq) + H 2O (l) Metalloids Display intermediate properties Shiny, but brittle Semiconductors 7.7 Trends for Group 1A and Group 2A Metals Alkali metals (IA) are soft, metallic solids. The name comes from the Arabic word for ashes. They are found only in compounds in nature, not in their elemental forms. They have low densities and melting points. They also have low ionization energies. Their reactions with water are famously exothermic. Alkali metals (except Li) react with oxygen to form peroxides. K, Rb, and Cs also form superoxides: K + O 2 KO 2 They produce bright colors when placed in a flame. Alkaline Earth Metals (IIA) Alkaline earth metals have higher densities and melting points than alkali metals. Their ionization energies are low, but not as low as those of alkali metals. 4
5 Beryllium does not react with water and magnesium reacts only with steam, but the others react readily with water. Reactivity tends to increase as you go down the group. 7.8 Trend for Selected Nonmetals Group 6A Oxygen, sulfur, and selenium are nonmetals. Tellurium is a metalloid. The radioactive polonium is a metal. Oxygen There are two allotropes of oxygen: O 3, ozone There can be three anions:, oxide 2, peroxide 1, superoxide It tends to take electrons from other elements (oxidation). Sulfur Sulfur is a weaker oxidizer than oxygen. The most stable allotrope is S 8, a ringed molecule. Group VIIA: Halogens The halogens are prototypical nonmetals. The name comes from the Greek words halos and gennao: salt formers. They have large, negative electron affinities. Therefore, they tend to oxidize other elements easily. They react directly with metals to form metal halides. Chlorine is added to water supplies to serve as a disinfectant. Group VIIIA: Noble Gases The noble gases have astronomical ionization energies. Their electron affinities are positive. Therefore, they are relatively unreactive. They are found as monatomic gases. Xe forms three compounds: XeF 2 XeF 4 XeF 6 Kr forms only one stable compound: KrF 2 The unstable HArF was synthesized in
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