Periodic Properties. of the Elements. 2009, Prentice-Hall, Inc. Periodic Properties of the Elements. 2009, Prentice-Hall, Inc.

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1 Chemistry, The Central Science, 11th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 7 John D. Bookstaver St. Charles Community College Cottleville, MO Chapter 7 11, 19, 21, 25, 39, 45, 47, 49, 61, 65, 69, 71 Development of Table in the same group generally have similar chemical properties. Physical properties are not necessarily similar, however. 1

2 During the nineteenth century, chemists began to categorize the elements according to similarities in their physical and chemical properties. The end result of these studies was our modern periodic table. Dmitri Mendeleev In 1869 he published a table of the elements organized by increasing atomic mass Development of Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped. 2

3 Development of Table Mendeleev, for instance, predicted the discovery of germanium (which he called ekasilicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon. Law Henry Moseley In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) elements*. He rearranged the elements in order of increasing atomic number. * There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus

4 X-Ray Spectra Moseley 1913 X-ray emission is explained in terms of transitions in which e - drop into orbits close to the atomic nucleus. Correlated frequencies to nuclear charges. = A (Z b) 2 Used to predict new elements (43, 61, 75) later discovered. Alkali Metals Alkaline Earths The table Transition Metals Halogens Main Group Noble Gases Main Group Lanthanides and Actinides Trends In this chapter, we will rationalize observed trends in Sizes of atoms and ions. Ionization energy. Electron affinity. 4

5 Effective Nuclear Charge In a many-electron atom, (1) electrons are both attracted to the nucleus and (2)repelled by other electrons. The nuclear charge that an electron experiences depends on both factors. Screening and Penetration Z eff = Z S Z 2 E n = - R eff H n 2 Atomic Radius increase top to bottom down a group each additional electron shell shields the outer electrons from the nuclear charge Z eff = Z - S where Z eff => effective nuclear charge Z => nuclear charge, atomic number S => shielding constant 5

6 Atomic Radius increases from upper right corner to the lower left corner What Is the Size of an Atom? The bonding atomic radius is defined as one-half distance between covalently bonded nuclei. Atomic Radius 6

7 Atomic Radius Ionic Radius same trends as for atomic radius positive ions smaller than atom negative ions larger than atom Cationic Radii 7

8 Anionic Radii Sizes of Ions Ionic size depends upon: The nuclear charge. The number of electrons. The orbitals in which electrons reside. Sizes of Ions Cations are smaller than their parent atoms. The outermost electron is removed and repulsions between electrons are reduced. 8

9 Sizes of Ions Anions are larger than their parent atoms. Electrons are added and repulsions between electrons are increased. Ions increase in size as you go down a column. This is due to increasing value of n. Sizes of Ions Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge. 9

10 Ionization Energy The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. The first ionization energy is that energy required to remove first electron. The second ionization energy is that energy required to remove second electron, etc. 9-4 Ionization Energy Mg(g) Mg + (g) + e - Mg + (g) Mg 2+ (g) + e - I 1 = 738 kj I 2 = 1451 kj Z 2 I = R eff H n 2 Ionization Energy It requires more energy to remove each successive electron. When all valence electrons have been removed, the ionization energy takes a quantum leap. 10

11 Trends in First Ionization Energies As one goes down a column, less energy is required to remove the first electron. For atoms in the same group, Z eff is essentially the same, but the valence electrons are farther from the nucleus. Trends in First Ionization Energies Generally, as one goes across a row, it gets harder to remove an electron. As you go from left to right, Z eff increases. Trends in First Ionization Energies However, there are two apparent discontinuities in this trend. 11

12 Trends in First Ionization Energies The first occurs between Groups IIA and IIIA. In this case the electron is removed from a p-orbital rather than an s-orbital. The electron removed is farther from nucleus. There is also a small amount of repulsion by the s electrons. Trends in First Ionization Energies The second occurs between Groups VA and VIA. The electron removed comes from doubly occupied orbital. Repulsion from the other electron in the orbital aids in its removal. Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e Cl 12

13 Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. Trends in Electron Affinity There are again, however, two discontinuities in this trend. Trends in Electron Affinity The first occurs between Groups IA and IIA. The added electron must go in a p-orbital, not an s-orbital. The electron is farther from nucleus and feels repulsion from the s-electrons. 13

14 Trends in Electron Affinity The second occurs between Groups IVA and VA. Group VA has no empty orbitals. The extra electron must go into an already occupied orbital, creating repulsion. of Metal, Nonmetals, and Metalloids Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties. 14

15 Metals versus Nonmetals Metals tend to form cations. Nonmetals tend to form anions. Metals They tend to be lustrous, malleable, ductile, and good conductors of heat and electricity. Compounds formed between metals and nonmetals tend to be ionic. Metal oxides tend to be basic. Metals 15

16 Nonmetals These are dull, brittle substances that are poor conductors of heat and electricity. They tend to gain electrons in reactions with metals to acquire a noble gas configuration. Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic. Nonmetals Metalloids These have some characteristics of metals and some of nonmetals. For instance, silicon looks shiny, but is brittle and fairly poor conductor. 16

17 Group Trends Alkali metals are soft, metallic solids. The name comes from the Arabic word for ashes. Alkali Metals Alkali Metals 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. 17

18 Alkali Metals Their reactions with water are famously exothermic. Alkali Metals 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 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. 18

19 Alkaline Earth Metals 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. 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 2 O 3, ozone There can be three anions: O 2, oxide O 2 2, peroxide O 1 2, superoxide It tends to take electrons from other elements (oxidation). 19

20 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. Group VIIA: Halogens 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 20

21 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. Group VIIIA: Noble Gases Xe forms three compounds: XeF 2 XeF 4 (at right) XeF 6 Kr forms only one stable compound: KrF 2 The unstable HArF was synthesized in