Periodic Variation in Physical Properties Sizes of atoms and ions Ionization energy Electron affinity Metallic properties
Properties and Electronic Structure Properties depend on • valence electrons • Effective nuclear charge (net charge on an electron) • Protons-electrons attraction (increases) • Electron-electron repulsion (decreases) • Penetration of the orbitals (greater penetration – increases the Zeff) • The shielding effect • Core electrons • Size of the atom
Effective Nuclear Charge Effective Nuclear charge, Zeff: the actual nuclear charge a valence electron experiences. In a many-electron atom, the Zeff depends on two factors: 1. Attraction between electrons and nucleus 2. Repulsion between electrons in orbitals
Effective Nuclear Charge In SWE, electron is treated individually in a field of net nuclear charge that determines the Zeff.. The effective nuclear charge, Zeff, is found: Zeff = Z−σ Z =atomic number σ =screening constant ~ # number of inner electrons, but not equal.
Effective Nuclear Charge (Zeff) • Proton-electron attractionincreases the Zeff. • It depends on: Size of the atom: Coulomb’s Law Penetration of orbitals: (10% probability being next to the nucleus)
Effective Nuclear Charge (Zeff) • Electron-electron repulsion reduces the Zeff • Smaller Zeff on valence electron, easier to remove the electron from the atom
Effective Nuclear Charge on Valence Electrons • Expected Nuclear Charge by the valence electron in • Li atom; Be atom; B atom ….. But: Li. Zeff on the 3rd electron 1.3. Why? Zeff = Z – (shielding by other electrons) Zeff = Z – σ
Effective Nuclear Charge on 2s Electron in Li http://www.wou.edu/las/physci/ch412/Periodic%20trends/periodic_trends.htm
Shielding and (Zeff) in a Group • The electron on the outside energy level has to look through all the other energy levels to see the nucleus: it is shielded from the nucleus by all the inner electrons • Less attraction between the valence electron and the nucleus • Lower effective nuclear charge • If shielding were completely effective, Zeff = 1 • Why isn’t it? 24
Zeff in a Group • Electrons enter new shell (energy level) • Size of atom increases • Zeff decreases: decreased force of attraction (Coulomb’s Law – inversely proportional to r2) • In a Group: Zeff decreases; shielding Increases.
Shielding in a Group • The more positive the nucleus (higher n, higher atomic number), the closer the inner orbitals to the nucleus provides for better shielding of the outer electrons (easier to remove electron in Na than in Li or H) by the core electrons. • In lithium 1s orbital is the same shape as a hydrogen 1s orbital, but it is smaller because the electron is more strongly attracted to the nucleus. • The sodium 1s is even smaller.
Shielding: Electron Repulsion and Orbital Energy Extent of shielding depends on: 1. Core e- provide more shielding than valence electrons • Electrons in the same shell ( same n), extend of shielding depends on l (penetration effects); • s > p > d > f • Correlates with degree of penetration. 3. Electrons in the same sub-shell (same value of n and l) do not effectively shield one another. Electrons are in the degenerate orbitals
Shielding in Period • The electron on the outside energy level has to look through all the other energy levels to see the nucleus • A second electron in the same energy level has almost the same shielding. • The third one has almost the same shielding • Zeff (effective nuclear charge) increases as atomic number increases in the Period. • Why? 29
Shielding Effects: Examples First Ionization energies in kJ/mol • He atom (1s2 + 2 protons): 2372 • He+ ion: (1s1 + 2 protons): 5250 ( it is not 2X 2372, but larger)) • In He, the second electron repels the first, making it easier to remove (one electron shields the other from the full effect of the nucleus); not in He+1 (there is only one electron) • In Li (1s22s1): 520; Li+2 (1s1): 2954 • In Li: we have two1s electrons. They shield very effectively the electron in 2s and smaller radius, stronger attraction
Sizes of Atoms The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.
Atomic Size in a Group (shielding dominates) • Each new member has one more level of inner electrons • Inner electrons shield the outer electrons very effectively • Zeff increases very slightly (more protons) • Atoms get larges, as n-increases • Atoms radii increase in a group
Atomic Size in a Period: Zeff Dominates • Electrons added to the same shell • shielding by inner electrons changes very slightly, if at all. • Outer electrons shield each other poorly • Zeff rises significantly, electrons pulled closer to nucleus • Atomic radius decreases in a period. Na Mg Al C Si P Cl Ar
Overall Rb K Na Ga Li Atomic Radius (nm) Kr Zn V Ar Ne H 10 Atomic Number
Radii in Transition Elements • Size shrinks for the first two to three members because of increased nuclear charge • After: the size remains relatively constant as repulsion of d-electrons (increased radius) counteracts the increase in Zeff. • d-electrons shield very well, but p-orbital penetrates much more than d-orbital: thus Ga (135 pm) is much smaller than Ca (197 pm) • Another anomaly: 13Al (143 pm) versus 31Ga (135 pm). Filling the d-orbitals causes major contraction.
Atomic Radius: Examples • Using only the periodic table, rank each set of main-group elements in order of decreasing atomic size: • A) Ca, Mg, Sr • B) K, Ga, Ca • C) Br, Rb, Kr • D) Sr, Ca, Rb
Order the following according to increasing atomic radius. Ge Si Se Cl • Ge < Si < Se < Cl • Se < Si < Ge < Cl 3. Si < Cl < Ge < Se 4. Cl < Si < Se < Ge 5. Si < Ge < Se < Cl
Ionic Radius Ionic radius: the radius of a cation or an anion. Determines the physical and chemical properties of an ionic compound such as • Crystal structure • Melting point • solubility
Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed. 8.3
Sizes of Ions • Ionic size depends upon: • Nuclear charge. • Number of electrons. • Orbitals in which electrons reside.
Sizes of Ions • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge. • All have the configuration 1s12s22p6 (10 electrons)
Ionic Radius: Examples • In each of the following pairs, indicate which one of the two species is larger: • A) N3- or F- • B) Mg+2 or Ca+2 • C) Fe+2 or Fe+3 • Explain your choice.
Order the following according to increasing atomic/ionic radius. N3- Li+ C O2- • C < Li+ < O2- < N3- • N3- < O2- < C < Li+ 3. Li+ < C < N3- < O2- 4. Li+ < C < N3- < O2- 5. Li+ < C < O2- < N3-
Ionization Energy Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • First ionization energy is that energy required to remove first electron. • Second ionization energy is that energy required to remove second electron, etc. • D:\Chapter_07\Present\eMedia_Library\IonizationEnergies\IonizationEnergies.html
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.
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, Zeff is essentially the same, but the valence electrons are farther from the nucleus.
Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell
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, shielding is almost the same.
Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.
Trends in First Ionization Energies • The first occurs between Groups IIA and IIIA. • Electron removed from p-orbital rather than s-orbital • Electron farther from nucleus • Small amount of repulsion by selectrons.
Trends in First Ionization Energies • The second occurs between Groups VA and VIA. • Electron removed comes from doubly occupied orbital. • Repulsion from other electron in orbital helps in its removal.
Trends in Successive IE • The effective nuclear charge increases as you remove electrons. Thus IE3>IE2>IE1 • Big jump after all outer electrons removed. • It takes much more energy to remove a core electron than a valence electron because there is less shielding, smaller size (energy shell removed) • greater effective nuclear charge.
Successive Ionization Energies (kJ/mol) Na Mg Al Si P S Cl Ar IE1 496 738 578 787 1012 1000 1251 1520 IE2 4562 1451 1817 1577 1903 2251 2297 2665 IE3 6912 7733 2745 3231 2912 3361 3822 3931 IE4 9543 10540 11575 4356 4956 4564 5158 5770 IE5 13353 13630 14830 16091 6273 7013 6540 7238 IE6 16610 17995 18376 19784 222338495 9458 8781 IE7 20114 21703 23293 23783 25397 27106 11020 11995
Explain the trend in IE • For Mg • IE1 = 735 kJ/mole • IE2 = 1445 kJ/mole • IE3 = 7730 kJ/mole • For Al • IE1 = 580 kJ/mole • IE2 = 1815 kJ/mole • IE3 = 2740 kJ/mole • IE4 = 11,600 kJ/mole
Which will have the highest ionization energy? • C • N • O • Al • Si
Which will be the largest? I= ionization energy 1. I1 of Na 2. I2 of Na 3. I1 of Mg 4. I2 of Mg 5. I3 of Mg #5
Electron Affinity D:\Chapter_07\Present\eMedia_Library\PeriodicTrendElctrnAffnity\PeriodicTrendElctrnAffnity.html
Electron Affinity, kJ/mol • Electron affinity: the energy change that occurs when an electron is accepted by an atom in gaseous state. X(g) + e-→ X-(g) • A large negative value indicates a strong attraction between the atom and the added electron Cl(g) + e- → Cl-(g) ΔE = -349 kJ/mol • A positive value indicates the addition of electron is unfavorable Ne(g) + e- → Ne-ΔE = 29 kJ/mol
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. • Added electron must go in p-orbital, not s-orbital. • Electron is farther from nucleus and feels repulsion from s-electrons.