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Unit 5.2 Periodic Trends Teacher: Dr. Van Der Sluys Objectives • Trends in – Electron configuration – Atomic Radii – Electronegativity – Ionization Energy – Electron Affinity Periodic Table e- configuration from the periodic table 1 IA 18 VIIIA 1 H 1s1 2 IIA 13 IIIA 2 Li 2s1 Na 3s1 Be 2s2 B 2p1 Mg 3s2 3 IIIB 4 IVB 5 VB K 4s1 Rb 5s1 Ca 4s2 Sr 5s2 Sc 3d1 V Cr Mn Fe Co 3d3 4s13d5 3d5 3d6 3d7 Nb Mo Tc Ru Rh 4d3 5s14d5 4d5 4d6 4d7 6 Cs 6s1 7 Fr 7s1 Ba 6s2 Ra 7s2 Ti 3d2 Zr 4d2 Hf 5d2 3 4 5 Y 4d1 La 5d1 Ac Rf 6d1 6d2 6 VIB 7 VIIB 8 Ta W Re Os 5d3 6s15d5 5d5 5d6 Db Sg Bh 6d3 7s16d5 6d5 9 VIIIB Ir 5d7 Hs Mt 6 6d 6d7 10 11 IB 12 IIB Ni Cu Zn 3d8 4s13d10 3d10 Ni Ag Cd 4d8 5s14d10 4d10 Ni Au Hg 5d8 6s15d10 5d10 Al 3p1 Ga 4p1 In 5p1 Tl 6p1 14 IVA 15 VA 16 VIA 17 VIIA He 1s2 F •B C N O Ne •2p2 1 2p3 2p4 2p5 2p6 2p Cl Ar Si P S 3p2 3p3 3p4 3p5 3p6 Ge As Se Be Kr 4p2 4p3 4p4 4p5 4p6 Sn Sb Te I Xe 5p2 5p3 5p4 5p5 5p6 Pb Bi Po At Rn 6p2 6p3 6p4 6p5 6p6 1 Periodic Table and Subshells Metals, Nonmetals and Metalloids • Metals are on the bottom left side of the table and represent the vast majority of the total number of elements. • Nonmetals are on the top right side of the table. • Metalloids are along the stair case, except aluminum, which is a metal. Periodic Table as a Map West (South) Mid-plains East (North) METALS Alkali Alkaline Transition These elements tend to give up e - and form CATIONS METALLOID NON-METALS Noble gas Halogens Calcogens These elements tend to accept e - and form ANIONS These elements will give up e- or accept e- • The periodic table can be classified by the behavior of their electrons 1 IA 1 18 VIIIA 2 IIA 13 IIIA 14 IVA 15 VA 16 VIA 17 VIIA 2 3 3 IIIB 4 IVB 5 VB 6 VIB 7 VIIB 8 9 VIIIB 10 11 IB 12 IIB 4 5 6 7 2 Atomic Radius • The atomic radius of an element is half of the distance between the centers of two atoms of that element that are just touching each other. Generally, the atomic radius decreases across a period from left to right and increases down a given group. The atoms with the largest atomic radii are located in Group I and at the bottom of groups. Electronegativity • Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy (IE) and electron affinity. • Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus. In a group, the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly electronegative element is fluorine . Ionization Energy • The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be. The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy. Ionization energies increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet. 3 Electron Affinity • Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about the electron affinities of certain groups in the periodic table. The Group IIA elements, the alkaline earths, have low electron affinity values. These elements are relatively stable because they have filled s subshells. Group VIIA elements, the halogens, have high electron affinities because the addition of an electron to an atom results in a completely filled shell. Group VIII elements, noble gases, have electron affinities near zero, since each atom possesses a stable octet and will not accept an electron readily. Elements of other groups have low electron affinities. 1. Electronegativity Trend Metals < 1.6 Nonmetals > 2.0 http://library.kcc.hawaii.edu/external/chemistry/electronegativity.html 2. Trend in Atomic Radius 4 3. Trend in Ionization Potential Ionization potential: The energy required to remove the valence (outer most) electron from an atomic species. Largest toward NE corner of PT since these atoms hold on to their valence e- the tightest. 4. Trend in Electron Affinity Electron Affinity: The energy release when an electron is added to an atom. Most favorable toward NE corner of PT since these atoms have a great affinity for e-. Trends in Groups: Shells • The trends down a group can best be explained by the increasing number of shells of electrons. • Each element in a group has the same number of valence electrons, but they are always in shells that have increasing n values and are further from the nucleus. 5 Core Charge • The trends in periods from left to right can best be explained by – the increasing core charge • Core charge = # protons - inner shell electrons Core Charge Group 1 2 H 1 - 0 = +1 Li 3 - 2 = +1 Na 11 - 10 = +1 K 19 - 18 = +1 Rb 37 - 36 = +1 Cs 55 - 54 = +1 He 2 - 0 = +2 Be 4 - 2 = +2 Mg 12 Ca 20 Sr 38 Ba 56 13 B 5 - 2 = +3 Al 13 Ga 31 In 49 Tl 81 14 C 5 - 2 = +4 Si 14 Ge 32 Sn 50 Pb 82 15 N 7 - 2 = +5 P 15 As 33 Sb 51 Bi 83 16 O 8 - 2 = +6 S 16 Se 34 Te 52 Po 84 17 F 9 - 2 = +7 Cl 17 Br 35 I 53 At 85 18 Ne 10 - 2 = +8 Ar 18 Kr 36 Xe 54 Rn 86 Core charge = atomic number - # core electrons Summary of Trend 1. Electronegativity Electronegativity:: Largest toward NE corner of PT 3. Ionization energy: Largest toward NE corner of PT 2. Atomic Radius: Largest toward SW corner of PT 6