The Periodic Table

1. The Periodic Table • Jedediah Mephistophles Soltmann

2. Dmitri Mendeleev • Studied the properties of elements and organized the elements by similar properties (families) and by increasing atomic mass. • He left blanks for elements he knew had to exist, such as:

3. Ekaaluminum (gallium) • In 1871 Mendeleev predicted the existence of yet undiscovered element he named eka-aluminum (because of its proximity to aluminum in the periodic table). The table below compares the qualities of the element predicted by Mendeleev with actual characteristics of Gallium (discovered in 1875).

4. Ekasilicon (Germanium) • Germanium was isolated in 1882, and provided the best confirmation of the theory up to that time, due to its contrasting more clearly with its neighboring elements than the two previously confirmed predictions of Mendeleev do with theirs.

5. Effective Nuclear Charge • Protons in the nucleus attract the electrons • Electrons repel each other. • So inner electrons push the outer electrons (shielding), negating much of the pull of the nucleus. Thus higher energy levels means less lower effective nuclear charge. • Zeff = Z - S

6. Calculating Zeff • Na11: 1s2 2s2 2p6 3s1Zeff = 11 – 10 = 1 • Mg12: 1s2 2s2 2p6 3s2Zeff = 12 – 10 = 2 • Cl17: 1s2 2s2 2p63s2 3p5Zeff = 17 – 10 = 7 • As you can see, the outer electrons of chlorine are pulled more by the nucleus than those of the sodium or magnesium.

7. Isoelectronic Atoms/Ions • Iso = same • electronic = from electrons • Isoelectronic particles are those with the same # of electrons in the same configuration.

8. Size of Atoms

9. Atomic Size on the Periodic Table • As we compare elements in a period, the Zeff increases which means that the valence electrons are being pulled harder by the nucleus. So, from left to right, the atomic size decreases.

10. Atomic Size on the Periodic Table • As we compare elements in a family, the main difference is the number of shells. From top to bottom, the number of shells increases, so the atomic size increases.

11. Do Now • What is the effective nuclear charge of: • An electron in the 3rd energy level of Mo? • An electron in the 2nd energy level of S? • An electron in the 4th energy level of Br? • List these elements in size order: P, S, As, Se • List these particles in size order: S, S2-, O

12. Do Now Answers • What is the effective nuclear charge of: • An electron in the 3rd energy level of Mo? • Zeff=42-10=32 • An electron in the 2nd energy level of S? • Zeff = 16 - 2 = 14

13. Do Now Answers • An electron in the 4th energy level of Br? • Zeff=35-28=7 • List these elements in size order: P, S, As, Se • S, P, Se, As • List these particles in size order: S, S2-, O • O, S, S2-

14. Bond Length • When a bond forms, two atoms are held next to each other by electrical attractions. So the distance from nucleus to nucleus is called the bond length. • Bond length is thus the sum of atomic radii. • For example a C-H bond has a length of 1.14A, because C has a radius of .77A and H has a radius of .37A. .37A + .77A = 1.14A.

15. Chart of Atomic Radii

16. What is the bond length of: • C-S? • S-H? • N-Cl? • Na-Cl?

17. What is the bond length of: • C-S = 1.79A • S-H = 1.39A • N-Cl = 1.74A • Na-Cl = 2.79A

18. Why is the bond length of NaCl 2.79A? • NaCl is an ionic compound and thus depends on the radii of the ions, not the atoms! • Na+ has a radius of .98A and Cl- has a radius of 1.81A. Thus the sum is 2.79A!

19. Ionic Radii

20. Ionization Energy • Ionization energy is the minimum energy required to remove an electron from the ground state of an isolated gaseous atom, or ion. • Na(g) --> Na+ (g) + e- IE = 496 kJ/mol • Na+ (g) --> Na2+ (g) + e- IE = 4560 kJ/mol • Why does the first electron come from sodium so much easier than the 2nd?

21. Because... • Na11: 1s2 2s2 2p6 3s1 • The first electron comes from the 3rd energy level, but the next electron must come from a lower energy level, closer to the nucleus, with a higher Zeff. Thus it takes a lot more energy to get 2 electrons than 1 from a sodium atom.

22. So think about this... • An element in the 3rd period requires 787 kJ/mol to remove its first electron. • It requires 1575 kJ/mol to remove the 2nd electron. • It requires 3220 kJ/mol to remove the 3rd electron. • It requires 4350 kJ/mol to remove the 4th electron. • It requires 16,100kJ/mol to remove the 5th electron. • What element is this?

23. Chart of Successive Ionizations The Answer is Silicon

24. Ionization and the Periodic Table • It is easier to remove a valence electron from a bigger element than a smaller one. Why? • A valence electron in a smaller atom is closer to the nucleus, and thus held more tightly by electrical attraction.

25. Ionization across a Period • We now know that the size of the atoms decreases as we compare the elements going from left to right across a period. This means that more energy is required to remove electrons from elements on the right (nonmetals) and less for elements on the left (metals). • Ionization energy increases from left to right. • Could this be why metals give off electrons easily?

26. Ionization within a Family • We also know that each successive member of a family is larger because of additional energy levels. This means that elements near the top of the periodic table require more energy to remove an electron than elements near the top. • Ionization energy decreases from top to bottom.

27. Ionization as a Periodic Function

28. Electron Affinity • Instead of taking electrons, we could also add electrons. One such property of atoms is called Electron Affinity. • Electron affinity is electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions. • Like Ionization energy, there are successive electron affinities.

29. However... • Electron affinity is not a clear periodic property like ionization energy. The reason is that energy shifts based on whether subshells or orbitals are partially filled or completely filled. This makes it hard to come up with a good rule. • Still, it makes sense that a smaller atom can attract electrons better than a larger atom. So more energy is released when a smaller atom captures an electron than a larger atom.

30. Electron Affinity and the Periodic Table • If smaller atoms release more energy, than electron affinity should increase from left to right across a period. • Likewise, electron affinity should decrease from top to bottom.

31. Electron Affinity and the Periodic Table

32. Metals What defines a metal? We’ve used words like: luster, ductility, malleability, & conductivity. Why do metals behave this way?

33. Metallic Behavior Metals tend to be larger atoms. Since it is easier to remove an electron from a larger atom, it should make sense then that metals tend to form cations. Conversely we can say that the larger an atom is (or the lower its first ionization is) the more metallic the atom is. So if we compared O, S, and Se (all nonmetals) we could say that selenium, being the largest atom, is the most metallic - even though it is a nonmetal.

34. Nonmetallic Behavior Nonmetals tend to be smaller atoms. Since it is easier to add an electron to a smaller atom, it should make sense then that nonmetals tend to form anions. Conversely we can say that the smaller an atom is (or the higher its first ionization is) the more nonmetallic the atom is. So if we compared Li, Na, and K (all metals) we could say that lithium, being the smallest atom, is the most nonmetallic - even though it is a metal.