1. Periodic Properties of the Elements

2. The Periodic Table The modern periodic table was developed in 1872 by Dmitri Mendeleev (1834-1907). A similar table was also developed independently by Julius Meyer (1830-1895). The table groups elements with similar properties (both physical and chemical) in vertical columns. As a result, certain properties recur periodically.

3. The Periodic Table Mendeleev left empty spaces in his table for elements that hadn’t yet been discovered. Based on the principle of recurring properties, he was able to predict the density, atomic mass, melting or boiling points and formulas of compounds for several “missing” elements.

4. The Periodic Table

5. metal/non-metal line

6. The Periodic Table The periodic table is based on observations of chemical and physical behavior of the elements. It was developed before the discovery of subatomic particles or knowledge of the structure of atoms. The basis of the periodic table can be explained by quantum theory and the electronic structure of atoms.

7. Quantum Numbers In addition to n, the principal quantum number, there are three additional quantum numbers which describe the type of orbital ( l ), the spatial orientation of the orbital (m l ), and the spin of the electron (m s ).

8. Quantum Numbers The magnetic quantum number (m l ) specifies the spatial orientation of the orbital. An example is to distinguish between the p x, p y or p z orbitals.

9. Electron Spin Each orbital, regardless of type, can contain zero, one or two electrons. If two electrons occupy the same orbital, they must spin in opposite directions. The spin is quantized, and can be expressed using quantum numbers, or simply specifying the spin as up or down or clockwise and counter-clockwise.

10. The Pauli Exclusion Principle Quantum mechanics dictates that no two electrons in an atom can have the same four quantum numbers. Another way of stating the Pauli Exclusion Principle is that if electrons occupy the same orbital, they must have opposite spins.

11. Multi-electron Atoms Orbitals of any type can be empty, or have 1 or two electrons. Experimental data indicate that if two electrons are in the same orbital, they will spin in opposite directions.

12. Energy Levels In any atom or ion with only 1 electron, the principal quantum number, n, determines the energy of the electron. For n=2, the 2s and 2p orbitals all have the same energy.

13. Energy Levels Likewise, the 3s, 3p and 3d orbitals are all degenerate, with the same energy.

14. Energy Levels In a multi-electron atom, there is interaction between electrons. As a result of this interaction, the various subshells of a principal quantum level will vary in energy.

15. Energy Levels

17. The energy diagram for the first three quantum levels shows the splitting of energies.

18. Energy Levels For a given value of n, the energies of the subshells is as follows: ns<np<nd<nf

19. Energy Levels The subshells have different energies due to the penetrating ability for each type of orbital. Electrons in a 2s orbital can get nearer to the nucleus than those in a 2p orbital.

20. Energy Levels The electrons in the 3s orbital (top diagram) have higher probability to be found near the nucleus, and thus greater penetrating ability than those in 3p or 3d orbitals.

21. Multi-electron Atoms Electron configurations are a way of noting which subshells of an atom contain electrons. Although much of the periodic table was developed before the concept of electron configurations, it turns out that the position of an element on the periodic table is directly related to its electron configuration.

22. Multi-electron Atoms

23. Electron Configurations Write the complete electron configurations for nitrogen and zinc. Write the complete electron configurations for nitrogen and zinc. How many unpaired electrons does each atom have? How many unpaired electrons does each atom have? What is the short hand notation for each element. What is the short hand notation for each element.

24. Hund’s Rule When electrons occupy degenerate orbitals, they occupy separate orbitals with parallel spins. This is the lowest energy, or ground state, configuration.

25. Multi-electron Atoms The electron configurations for Cr and Cu differ from that expected based on their positions in the periodic table.

26. Multi-electron Atoms Electron configurations also get less predictable for the elements near the bottom of the periodic table. With many quantum levels (n) occupied, the energy levels overlap and the lowest energy arrangement becomes more difficult to predict.

27. Periodic Trends Many of the properties of atoms show clear trends in going across a period (from left to right) or down a group. In going across a period, each atom gains a proton in the nucleus as well as a valence electron.

28. Periodic Trends The increase of positive charge in the nucleus isn’t completely cancelled out by the addition of the electron. Electrons added to the valence shell don’t shield each other very much. As a result, in going across a period, the effective nuclear charge (Z eff ) increases.

29. Effective Nuclear Charge The effective nuclear charge (Z eff ) equals the atomic number (Z) minus the shielding factor (σ). Z eff = Z-σ

30. Effective Nuclear Charge Z eff = Z-σ

31. Effective Nuclear Charge Electrons in the valence shell are partially shielded from the nucleus by core electrons.

32. Effective Nuclear Charge Electrons in p or d orbitals don’t get too close to the nucleus, so they are less shielding than electrons in s orbitals. As a result, effective nuclear charge increases across a period.

33. Periodic Trends

34. In going down a group or family, a full quantum level of electrons, along with an equal number of protons, is added. As n increases, the valence electrons are, on average, farther from the nucleus, and experience less nuclear pull due to the shielding by the “core” electrons. As a result, Z eff decreases slightly going down a group.

35. Trends- Atomic Radii Atomic radii are obtained in a variety of ways: 1. For metallic elements, the radius is half the internuclear distance in the crystal, which is obtained from X-ray data. 2. For diatomic molecules, the radius is half the bond length. 3. For other elements, estimates of the radii are made.

36. Trends- Atomic Radii Atomic radii follow trends directly related to the effective nuclear charge. As Z eff increases across a period, the electrons are pulled closer to the nucleus, and atomic radii decrease. As Z eff decreases down a group, the valence electrons experience less nuclear attraction, and the radius increases.

37. Trends- Atomic Radii Atomic size roughly halves across a period, and doubles going down a group.

38. Electron Configurations of Ions The atoms of the main group elements (groups IA-VIIA) will form ions by losing or gaining electrons. The resulting ion will have the same electron configuration as a noble gas (group VIIIA). These configurations are usually very stable.

39. Electron Configurations of Ions Atoms or ions with the same electron configuration (or number of electrons) are called isoelectronic. Atoms or ions with the same electron configuration (or number of electrons) are called isoelectronic. For example, Na +, Mg 2+, Ne, F -, and O 2- are isoelectronic. The size will decrease with increasing positive charge. O 2- > F - >Ne> Na + > Mg 2+

40. Electron Configurations of Ions When atoms lose electrons, the electrons are always removed from the highest quantum level first. For the first row of transition metals, this means that the electrons in 4s subshell are lost before the 3d subshell. Fe: [Ar]4s 2 3d 6 Fe 2+ : [Ar] 3d 6 or [Ar]4s 0 3d 6

41. Common Ionic Charges The charges of ions of elements in groups 1A-7A (the main groups) are usually predictable. Group 1A metals form +1 ions, group 2A metals form +2 ions, etc. The non-metals of group 5A have a -3 charge, those of group 6A have a -2 charge, and the halogens form ions with a -1 charge.

42. Typical Ionic Charges

43. Trends – Ionic Size Cations are always smaller than the neutral atom. The loss of one or more electrons significantly increases Z eff, resulting in the valence electrons being pulled closer to the nucleus.

44. Ionic Size - Cations Within a group, assuming the same ionic charge, the size of the ion increases going down the group, due to more core electrons shielding the nucleus as n increases.

46. Trends – Ionic Size Across period, the cations get more positive, and as a result, considerably smaller.

47. Trends – Ionic Size Anions are always larger in size than the neutral atom. The addition of one or more electrons results in greater electron- electron repulsion, which causes the valence electrons to “spread out” a bit.

48. Size of Anions

49. Anions are always larger than the neutral atom.

50. Size of Anions Within a group, assuming the same ionic charge, the size of the ion increases going down the group, due to more core electrons shielding the nucleus as n increases.

51. Trends – Ionic Size

52. Trends – Ionization Energy Ionization energy is the energy required to remove an electron from a mole of gaseous atoms or ions. X(g) + energy  X + (g) + e - Elements can lose more than one electron, so there are 1 st, 2 nd, 3 rd, etc., ionization energies.

53. Ionization Energy It always requires energy to remove an electron from a neutral atom. As more electrons are removed and the ion becomes positively charged, it requires increasingly greater energy to remove electrons.

54. Trends – Ionization Energy Ionization energy is a measure of how tightly the electrons in the highest occupied orbitals are held by the nucleus. As a result, it is directly related to the effective nuclear charge. Ionization energy increases going across a period, and decreases going down a group.

55. Trends – Ionization Energy

57. Ionization Energy

58. Trends – Electron Affinity Electron Affinity is the energy change when an electron is added to a mole of gaseous atoms. Electron Affinity is the energy change when an electron is added to a mole of gaseous atoms. X(g) + e -  X - (g) ∆E = electron affinity A negative value for the electron affinity indicates that the process releases energy, and that the anion is easily formed.

59. Trends – Electron Affinity There is less of a predictable trend in electron affinities. In going across a period (ignoring the noble gases), the electron affinity should become more negative. Although this is observed, there are many inconsistencies.

60. Trends – Electron Affinity

62. Trends- Electron Affinity In going down a group, the electron affinity should become less negative. Although this trend is observed, there is only a slight change in electron affinities within a group. There may also be inconsistencies in the general trend.

63. Trends – Electron Affinity

64. Trends- Electron Affinity The electron affinity of fluorine is less negative than expected. This may be due to additional electron-electron repulsion when an electron is added to such a small atom.

65. Metallic Character

66. Across a period, metallic behavior decreases. Non-metals are often crumbly solids, liquids or gases at room temperature.

67. Metallic Character Metallic behavior increases going down a group.

68. Group IA – the Alkali Metals In discussing the chemistry, preparation and properties of the group IA elements, it is important to remember that hydrogen is not a group IA metal. It’s properties and reactivity would place it within group 7A (diatomic non- metals), rather than group IA.

69. Group 1A Metals The group 1A metals are soft shiny metals with fairly low densities (Li, Na and K are less dense than water) and low melting points. Sodium melts at 98 o C, and cesium melts at 29 o C. The softness, low density and low melting points are the result of weaker metallic bonding due to only one valence electron in this group.

70. Group 1A Metals - Production Due to the high reactivity with oxygen and water, all of the metals are found in nature in ionic form (M 1+ ). The pure metal must be produced in an oxygen and water-free environment. Typically, an electrical current is passed through the melted chloride salt. The metal and the chlorine gas are collected separately.

71. Reactivity Trends The chemical behavior of the group IA metals illustrates periodic trends. As the valence electron occupies a higher quantum level, it experiences less nuclear attraction, and is more easily removed.

72. Group 1A Metals + Water The reaction with water forms hydrogen gas and the aqueous metal hydroxide. The reaction is so vigorous, that the hydrogen may ignite. 2 M(s) + 2 H 2 O(l)  H 2 (g) + 2 MOH(aq)

73. Metallic Character The group IA metals react with water to produce hydrogen and the metal hydroxide. Metallic behavior increases going down a group.