The Periodic Table

1. The Periodic Table

2. The Development of the Periodic Table • In the time of Mendeleev, we had not yet discovered the electron • Mendeleev was convinced that the properties of elements had a periodicity if we classified the elements in order of increasing atomic weight • He had so much confidence in the periodicity of properties that he suggested that some elements remained to be discovered, and he left "holes" in his table to accommodate them once they were discovered • He also predicted the properties of these remaining elements very well • One problem was the fact that Ar (element 18) had a greater mass than Na (element 19) • After the discovery of the nucleus by Rutherford, it was found that the number of protons in the nuclei was the important criterion (not the mass)

3. The Periodic Classification of Elements • Each column (group) of the modern periodic table shares the same electronic configuration for their valence electrons • There are a few exceptions for the transition metals, the lanthanides, and the actinides

4. The Periodic Classification of Elements • The valence electrons are the peripheral electrons (the electrons with the highest principal quantum number) of an atom • Valence electrons are those involved in the formation of chemical bonds • The fact that each member of a group shares the same number of valence electrons explains the similarities in their reactivity • N.B. The properties of the elements in IA, IIA, and VIIA groups are very similar through the group • In the IIIA, IVA, VA, et VIA groups, the properties of the elements sometimes change gradually when descending the group

5. The Electronic Configurations of Cations and Anions becomes • In an ionic compound, the cation of a representative element is produced by releasing electrons so that the cation reaches the electron configuration of a noble gas: • In an ionic compound, the anion of a representative element is produced by accepting electrons so that the anion reaches the electron configuration of a noble gas: becomes becomes becomes becomes or [Ne] becomes or [Ne] becomes or [Ne]

6. The Electronic Configurations of Cations and Anions • These ions of representative elements and noble gases are isoelectronic (have the same number of electrons and therefore share the same electronic configuration) • For transition metals, we can often find more than one type of cations and these cations are often not isoelectronic with a noble gas • ex.; neither Fe2+ nor Fe3+ are isoelectronic with a noble gas, but each can be found in nature • Even if the ns orbital is filled before the (n-1)d orbital, when a transition metal forms a cation, it is the ns orbital that empties first • The ns and (n-1)d orbitals are very close in energy, and the electron-electron and electron-nucleus change when going from atom to cation, thus causing a change in the energetic order between ns and (n-1)d • ex.; Mn2+ has the configuration [Ar]3d5 and Zn2+ has the configuration [Ar]3d10

7. The Effective Nuclear Charge • The effective nuclear charge, Zeff, is the positive charge felt by a valence electron • Zeff Z (where Z is the nuclear charge) since the other electrons screen the nucleus from the valence electrons • Zeff = Z -  where  is the screening constant • To give an example of the importance of this screening effect, consider the two ionization energies of He • The first electron is extracted with 2373 kJ of energy • It takes 5251 kJ to remove the second electron • Without the screening effect, the two ionization energies would be the same

8. The Effective Nuclear Charge • All of the electrons contribute to the screening effects, but in the simplest model,  is the number of electrons in the layers with principal quantum numbers lower than that of the valence electrons • ex.; Mg: 1s22s22p63s2 Z = 12,  = 10, Zeff = +2 • ex.; Mg2+: 1s22s22p6 Z = 12,  = 2, Zeff = +10 • N.B. These effective nuclear charges are what are felt by the valence electrons (n = 3 for the atom, et n = 2 for the cation, in these examples) • For a neutral atom of an element, the effective nuclear charge is the number of the element’s group (according to this simple model)

9. The Atomic Radius • The radius is a poorly defined property because the electron density, in principle, extends to infinity • The size of an atom can be estimated from experimental data • in a metal, the atomic radius is more or less half the distance between two adjacent atoms • for an element that exists as a diatomic molecule, the atomic radius is more or less half the length of the bond

10. The Atomic Radius • Going from left to right, the atomic radii of elements decrease • This decrease in the atomic radii is due to the increasing effective nuclear charge as we go from left to right • valence electrons are held closer to the nucleus by the increasing effective nuclear charge

11. The Atomic Radius • Descending a group, the atomic radius increases • When descending a group, the effective nuclear charge seen by the valence electrons remains more or less the same • However, the valence electrons occupy a higher principal quantum level, and thus the average distance from the nucleus increases

12. The Ionic Radius • The ionic radius is the radius of a cation or an anion • When an atom becomes an anion, its size increases because of the greater repulsion between the electrons, thus causing the electrons to occupy a larger volume • When an atom becomes a cation, the opposite occurs, and the radius decreases • In addition, in many cases, the valence shell with quantum level n is emptied and the valence shell then becomes the one with principal quantum number (n-1)

13. The Atomic Radius • Even if N3-, O2-, F-, Na+, Ca2+, et Al3+ are isoelectronic, their ionic radii are very different due to their effective nuclear charges • Their configuration is 1s22s22p6 and =2, and therefore Zeff = +5, +6, +7, +9, +10, and +11, respectively, going from N3- to Al3+ • A greater effective nuclear charge more strongly attracts the electrons towards the nucleus and the ionic radius gets smaller

14. The Atomic Radius • N.B. Atomic and ionic radii may differ substantially

15. Atomic and Ionic Radii • Example: Arrange the following atoms in order of decreasing radii : C, Li, Be. • Solution: In general, the radius decreases from left to right, thus Li > Be > C • Example: Decide which ion is the smallest in each of the following pairs : (a) K+, Li+ ; (b) Au+, Au3+ ; (c) P3-, N3- • Solution: (a) Li+, because the filled quantum level is n=1 rather than n=3 for K+ (b) Au3+ because there will be less repulsion between electrons (c) N3- because the filled quantum level is n=2 rather than n=3 for P3-

16. Ionization Energy • The ionization energy is the minimum energy required to extract an electron from an isolated atom in its ground state energy + X(g)  X+(g) + e- • The higher the ionization energy, the more difficult it is to remove an electron from the atom • There are also second and third, …., ionization energies energy + X+(g)  X2+(g) + e- energy + X2+(g)  X3+(g) + e- …. • It becomes more difficult to extract the second, third , ... . , electron since the repulsion between electrons decreases after each ionization and it is very difficult to separate an electron from a cation due to the favorable electrostatic interaction between these two opposite charges

17. Ionization Energy • The ionization energy decreases when descending a group • The effective nuclear charge does not change, but the valence electron that is removed is further from the core, and therefore more weakly retained and easily removed • The ionization energy increases, in general, when going from left to right on the periodic table • From left to right, the effective nuclear charge increases and the valence electron that is removed is more strongly held by the nucleus

18. Ionization Energy • There are two exceptions when going from left to right • A group IIIA (ns2np1) element is easier to ionize than a group IIA (ns2) element • For a group IIIA element, an electron is removed from the np orbital which is higher in energy than the ns orbital of a group IIA element • The increase in effective charge once again becomes the dominant factor when we reach group IVA

19. Ionization Energy • A group VIA (ns2np4) element is easier to ionize than a group VA (ns2np3) element • For a group VIA element, there are two electrons in one of the np orbitals, and these two electrons repel each other, and one of the two is more easily removed than the np electrons of an element in Group VA where each np electron has an orbital to itself • The increase in effective charge once again becomes the dominant factor when we reach group VIIA

20. The Unique Characteristics of the Second Period • The elements of the second period are Li to F • The properties of the first member of each representative group (the elements of the second period) are often very different from other group members, i.e., the first member is “unique” • eg.; The chemistry of Si resembles the chemistry of Ge more then it does the chemistry of C • eg.; The ability to form good double bonds is limited almost exclusively to the atoms of the second period • This effect is mainly due to the relatively small atomic radius of the first group member

21. The Diagonal Relationship • The diagonal relationship refers to similarities between two immediately adjacent elements of different groups and different periods on the periodic table (ie., diagonally) • The effect is most important for Li/Mg, Be/Al, and B/Si • The effect is largely due to the fact that the charge/radius ratio of their cations are very similar