Periodic Table

1. Periodic Table • It is a systematic catalog of the elements. • Elements are arranged in order of atomic number.

2. Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

3. Periodic Table • The rows on the periodic chart are periods. • Columns are groups. • Elements in the same group have similar chemical properties.

4. Groups These five groups are known by their names.

5. Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H). Gases – Group 18 and F, O , N , H, Cl Liquid – Br Rest are solids

6. Periodic Table Metalloids border the stair-step line (with the exception of Al, Po, and At).

7. Periodic Table Metals are on the left side of the chart. Properties – luster and high electrical and heat conductivity malleable All except Hg are solids at room temperature

8. Development of Periodic Table • Elements in the same group generally have similar chemical properties. • Physical properties are not necessarily similar, however.

9. Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

10. Development of Periodic Table Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon meaning under silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

11. Development of Periodic Table • 1913 – Henry Mosley developed concept of atomic number • Bombarded different elements with high-energy electrons, found each element produced X-rays of a unique frequency and that the frequency generally increased with the atomic mass • He arranged the X-ray frequencies in order by assigning them a whole number and called it the atomic number • He correctly identifies the atomic number as the number of protons in a nucleus

12. Development of Periodic Table • Concept of atomic number clarified some problems in the periodic table • Ex) atomic weight of Ar (18) is greater than that of K (19) yet the chemical properties of Ar are much more like those of Ne and Kr than those of Na and Rb • When elements are arranged by atomic number, rather than increasing atomic weight, Ar and K appear in the correct places • This made is possible to discover previously unknown elements

13. Periodic Trends • In this chapter, we will rationalize observed trends in • Sizes of atoms and ions. • Ionization energy. • Electron affinity.

14. Effective Nuclear Charge • Coulombs Law states that the strength of the interaction between two electrical charges depends on the magnitude of the charges and the distances between them .

15. Effective Nuclear Charge • The attractive force between an electron and the nucleus depends on the nuclear charge and on the average distance between the nucleus and the electron • The force increases as the nuclear charge increases and decreases as the electron moves further away

16. 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.

17. Effective Nuclear Charge Can treat each electron as though it were moving in the net electric field created by the nucleus and the electron density of the other electrons We view this net electric field as if it results from a single positive charge in the nuclues

18. Effective Nuclear Charge In many electron atoms the inner electrons partially screen outer electrons from the attraction of the nucleus

19. Effective Nuclear Charge The effective nuclear charge, Zeff, is found this way: Zeff = Z − S where Z is the atomic number and S (positive number) is a screening constant, usually close to the number of inner electrons.

20. Effective Nuclear Charge

21. Effective Nuclear Charge • For many-electron atom the energies of orbitals with the same n value increases with increasing l value • In C 1s22s22p6, the energy of the 2p orbital (l=1) is higher than that of 2s (l=0) even though both orbitals are in the n=2 shell • Difference in energies is due to the radial probability functions for the orbitals • The greater attraction between 2s electron and the nucleus leads to a lower energy for the 2s orbital than for the 2 p orbital

22. Effective Nuclear Charge • Trends in valence electron Zeff values: The effective nuclear charge increases from left to right across any period of the periodic table • Number of core electrons stays the same across the period, protons increase • Down a column the effective nuclear charge experienced by the valence electrons changes far less than it does across a period

23. Effective Nuclear Charge • n is larger than the value of n for the electron of interest contributes 0 • n is equal to the value of n for the electron of interest contribute 0.35 • n is 1 less than the n value for the electron of interest contribute .85 • n is even smaller contribute 1.00 • Ex) F • 1s22s22p5

24. Effective Nuclear Charge – Slater • Ex) F • 1s22s22p5 • Electron of interest n = 2 • n = 2 (7 electrons will look at the effect of 6 on 1 ) • S = (6*.35) + (2*.85) = 3.8 • Zeff = Z- S = 9 – 3.8 = 5.2+

25. What Is the Size of an Atom? The shortest distance separating the two nuclei during collisions is twice the radii of the atoms – non bonding atomic radius or van der Waals radius

26. What Is the Size of an Atom? The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. – shorter than the nonbonding atomic radius The attractive force between two adjacent atoms in the molecule leading to a chemical bond causes this

27. What Is the Size of an Atom? Example: I2 molecule, the distance separating the nuclei is observed to be 2.66Å, which means the bonding atomic radius of an iodine atom is (2.66Å)/2 = 1.33Å

28. What Is the Size of an Atom? Knowing atomic radii allows us to estimate bond lengths Ex) Cl-Cl bond length in Cl2= 1.99Å so bonding atomic radius is 0.99Ǻ In CCl4 bond length 1.77Ǻ bonding atomic radius Close to the sum (0.77+0.99)

29. Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row (due to increasing Zeff). …increase from top to bottom of a column (due to increasing value of n).

30. Sizes of Ions • Ionic size depends upon: • The nuclear charge. • The number of electrons. • The orbitals in which electrons reside.

31. Sizes of Ions • Cations are smaller than their parent atoms. • The outermost electron is removed and repulsions between electrons are reduced.

32. Sizes of Ions • Anions are larger than their parent atoms. • Electrons are added and repulsions between electrons are increased.

33. Sizes of Ions • Ions carrying the same charge ionic radius will increase in size as you go down a column. • This is due to increasing value of n.

34. Sizes of Ions • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge.

35. Ionization Energy • The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • The first ionization energy is that energy required to remove first electron. • The second ionization energy is that energy required to remove second electron, etc.

36. 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. • Supports the idea that only the outermost electrons are involved in the sharing and transfer of electrons that give rise to chemical bonding and reactions

37. 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.

38. 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.

39. Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

40. Trends in First Ionization Energies • The first occurs between Groups IIA and IIIA. • In this case the electron is removed from a p-orbital rather than an s-orbital. • The electron removed is farther from nucleus. • There is also a small amount of repulsion by the s electrons.

41. Trends in First Ionization Energies • The second occurs between Groups VA and VIA. • The electron removed comes from doubly occupied orbital. • Repulsion from the other electron in the orbital aids in its removal.

42. Electron Configurations of Ions • When electrons are removed from an atom to form a cation, they are always removed first from the occupied orbitals having the largest n value • Electrons added to form an anion are added to the empty or partially filled orbital having the lowest value of n

43. Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Measures the attraction or affinity of the atom fro the added electron – negative sign would indicate that energy is released during the process Cl + e− Cl−

44. Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row.

45. Trends in Electron Affinity There are again, however, two discontinuities in this trend.

46. Trends in Electron Affinity • The first occurs between Groups IA and IIA. • The added electron must go in a p-orbital, not an s-orbital. • The electron is farther from nucleus and feels repulsion from the s-electrons.

47. Trends in Electron Affinity • The second occurs between Groups IVA and VA. • Group VA has no empty orbitals. • The extra electron must go into an already occupied orbital, creating repulsion.