- Chapter 7 - Periodic Properties of the Elements Summary 7.1 Development of the periodic table 7.2 Effective nuclear charge 7.3 Size of atoms and ions 7.4 Ionization energy 7.5 Electron affinities 7.6 Metals, Nonmetals, and Metalloids 7.7 Group trends for the active metals 7.8 Group trends for selected Nonmetals 1
Development of Periodic Table The periodic table is the most significant tool that chemists use for organizing elements and recalling chemical facts. Elements in the same column contain the same number of outer-shell electrons or valence electrons and have consequently similar chemical properties. Development of Periodic Table Mendeleev and Meyer arranged the elements in order of increasing atomic weight. Certain elements were missing from their scheme. 2
Development of Periodic Table in 1871 Mendeleev noted that As properly belonged underneath P and not Si, which left a missing element underneath Si. He predicted a number of properties for this missing element (which he called eka-silicon or Germanium) with chemical properties similar to those of silicon. 3
Periodic Trends In this chapter, we will rationalize observed trends in: Sizes of atoms and ions. Ionization energy. Electron affinity. Effective Nuclear Charge In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors. The electron is attracted to the nucleus, but repelled by core electrons that shield or screen it from the full nuclear charge. This shielding is called the screening effect. 4
Effective Nuclear Charge 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. Na (Z=11) The effective charge Z eff = Z S = 11 + -10 + =1 + 5
Sizes of Atoms Consider a collection of argon atoms in the gas phase. During collisions electron clouds cannot penetrate each other to a significant extent. The apparent radius is determined by half of the closest distances separating the nuclei during such collisions. This radius is the nonbonding radius. Sizes of atoms in molecules The bonding atomic radius is defined as onehalf of the distance between covalently bonded nuclei which is shorter than nonbonding radius. 6
Example The distance separating the iodine (I 2 ) nuclei is 2.66 A, thus the radius is 1.33 A. Sizes of Atoms Knowing the atomic radii allows the estimation of the bond lengths between different elements in molecules. In the compound CCl 4 the measured length of C-Cl bond is 1.77 A which is very close to the sum of (0.77A + 0.99 A ) for C and Cl respectively 7
Sizes of Atoms Bonding atomic radius tends to decrease from left to right across a row due to increasing Z EFF which draws the electrons closer to the nucleus causing the atom to decrease in size. Bonding atomic radius tends to increase from top to bottom of a column due to increasing value of n 8
Sizes of Ions The radii of ions are based on the distances between ion nuclei in an ionic compound. Ionic size depends upon: Nuclear charge. Number of electrons. Orbitals in which electrons reside. Sizes of Cations Cations are smaller than their parent atoms. The outermost electron is removed and repulsions are reduced. 9
Sizes of Anions Anions are larger than their parent atoms. Electrons are added and repulsions are increased. Ions increase in size as you go down a column. Due to increasing value of n. Sizes of Ions 10
Sizes of Ions In an isoelectronic series, ions have the same number of electrons. In an isoelectronic series, ionic size decreases with an increasing nuclear charge. Ionization Energy Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. First ionization energy is the energy required to remove the first electron. Second ionization energy is the energy required to remove the second electron, etc. 11
Ionization Energy More energy is needed to remove each successive electron since the effective nuclear charges increase (more attractions). 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. 12
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, Zeff increases. Trends in First Ionization Energies Within each group the ionization energy generally decreases with increasing atomic number. The alkali metals have the lowest first ionization energies 13
Irregularities in First Ionization Energies There are two apparent irregularities: The first occurs between Groups IIIA and IIA. Electrons are more easily removed from p-orbitals than a s-orbital. The second occurs between Groups VA and VIA. Electron removed comes from doubly occupied orbital. Repulsion from other electron in the same orbital helps in its removal 14
Electron Affinity Energy change accompanying addition of an electron to a gaseous atom: Cl (g) + e- Cl (g) Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. The greater the attraction between atom and added electron the more negative is the electron affinity. A positive value means that an electron cannot be attached to an atom. 15
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. Added electron must go in p- orbital, not s- orbital (highly energetic subshell which is empty in the neutral atom. 16
Trends in Electron Affinity The second occurs between Groups IVA and VA. Group VA has no empty orbitals. Extra electron must go into occupied orbital, creating repulsion. Properties of Metal, Nonmetals, and Metalloids 17
Metals versus Nonmetals Metals tend to form cations. Nonmetals tend to form anions. Metals versus Nonmetals 18
Metals Tend to be malleable and good conductors of heat and electricity. Compounds formed between metals and nonmetals tend to be ionic. Metal oxides tend to be basic. Nonmetals brittle to hard substances that are poor conductors of heat and electricity. Tend to gain electrons in reactions with metals to acquire noble gas configuration. 19
Nonmetals Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic. Metalloids Have some characteristics of metals, some of nonmetals. For instance, silicon looks shiny, but is brittle and fairly poor conductor. 20
Group Trends Soft, metallic solids. Alkali Metals 21
Alkali Metals Found only as compounds in nature. Have low densities and melting points. Also have low ionization energies. Alkali Metals Their reactions with water are famously exothermic. Used in distillation stills. Story of the guy burning his hand in Notts. 22
Alkaline Earth Metals Have higher densities and melting points than alkali metals. Have low ionization energies, but not as low as alkali metals. Alkaline Earth Metals Beryllium does not react with water but others react readily with water. Reactivity tends to increase as you go down a group. 23
Group 6A Oxygen, sulfur, and selenium are nonmetals. Tellurium is a metalloid. The radioactive polonium is a metal. Oxygen Two allotropes: O 2 O 3, ozone Three anions: O 2, oxide O 2 2, peroxide O 2, superoxide Tends to take electrons from other elements (oxidation) 24
Sulfur Weaker oxidizing agent than oxygen. Most stable allotrope is S 8, a ringed molecule. 30 allotropes Group VIIA: Halogens Name comes from the Greek halos and gennao: salt formers 25
Group VIIA: Halogens Large, negative electron affinities Therefore, tend to oxidize other elements easily React directly with metals to form metal halides Chlorine added to water supplies to serve as disinfectant 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 2000. 26