1. Chapter 7 Periodic Properties of the Elements

2. Chapter 7 In 2002, there were 115 elements known. The majority of the elements were discovered between 1735 and 1843. How do we organize 115 different elements in a meaningful way that will allow us to make predictions about undiscovered elements? Development of the Periodic Table

3. Chapter 7 Arrange elements to reflect the trends in chemical and physical properties. First attempt (Mendeleev and Meyer) arranged the elements in order of increasing atomic weight. Certain elements were missing from this scheme. Example: In 1871, Mendeleev noted that As properly belonged underneath P and not Si, which left a missing element underneath Si. He predicted a number of properties for this element. In 1886 Ge was discovered. The properties of Ge match Mendeleev’s predictions well. Development of the Periodic Table

4. Mendeleev’s Periodic Table Dmitri Mendeleev

6. Chapter 7 Modern periodic table: arrange elements in order of increasing atomic number. Effective Nuclear Charge Effective nuclear charge is the charge experienced by an electron on a many-electron atom. The effective nuclear charge is not the same as the charge on the nucleus because of the effect of the inner electrons. Development of the Periodic Table

7. Chapter 7 Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nuclear charge. The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons. As the average number of screening electrons (S) increases, the effective nuclear charge (Z eff ) decreases. As the distance from the nucleus increases, S increases and Z eff decreases. Effective Nuclear Charge

8. Chapter 7 The ns orbitals all have the same shape, but have different sizes and different numbers of nodes. Consider: He: 1s 2, Ne: 1s 2 2s 2 2p 6, and Ar: 1s 2 2s 2 2p 6 3s 2 3p 6. The radial electron density is the probability of finding an electron at a given distance. Effective Nuclear Charge

10. Chapter 7 Consider a simple diatomic molecule. The distance between the two nuclei is called the bond distance. If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom. Sizes of Atoms and Ions

11. Chapter 7 As the principal quantum number increases, the size of the orbital increases. Consider the s orbitals. All s orbitals are spherical and increase in size as n increases. The spherical symmetry of the orbitals can be seen in the contour plots. Contour plots are connecting points of equal electron density. Sizes of Atoms and Ions

12. Chapter 7 Periodic Trends in Atomic Radii As a consequence of the ordering in the periodic table, properties of elements vary periodically. Atomic size varies consistently through the periodic table. As we move down a group, the atoms become larger. As we move across a period, atoms become smaller. There are two factors at work: principal quantum number, n, and the effective nuclear charge, Z eff. Sizes of Atoms and Ions

14. Chapter 7 Sizes of Atoms and Ions Periodic Trends in Atomic Radii As the principle quantum number increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence, the atomic radius increases. As we move across the periodic table, the number of core electrons remains constant. However, the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease.

15. Chapter 7 Sizes of Atoms and Ions Trends in the Sizes of Ions Ion size is the distance between ions in an ionic compound. Ion size also depends on nuclear charge, number of electrons, and orbitals that contain the valence electrons. Cations vacate the most spatially extended orbital and are smaller than the parent ion. Anions add electrons to the most spatially extended orbital and are larger than the parent ion.

17. Chapter 7 Sizes of Atoms and Ions Trends in the Sizes of Ions For ions of the same charge, ion size increases down a group. All the members of an isoelectronic series have the same number of electrons. As nuclear charge increases in an isoelectronic series the ions become smaller: O 2- > F - > Na + > Mg 2+ > Al 3+

18. Chapter 7 Ionization Energy The first ionization energy, I 1, is the amount of energy required to remove an electron from a gaseous atom: Na(g)  Na + (g) + e -. The second ionization energy, I 2, is the energy required to remove an electron from a gaseous ion: Na + (g)  Na 2+ (g) + e -. The larger ionization energy, the more difficult it is to remove the electron.

19. Chapter 7 Ionization Energy Variations in Successive Ionization Energies There is a sharp increase in ionization energy when a core electron is removed.

20. Chapter 7 Ionization Energy Periodic Trends in Ionization Energies Ionization energy decreases down a group. This means that the outermost electron is more readily removed as we go down a group. As the atom gets bigger, it becomes easier to remove an electron from the most spatially extended orbital. Ionization energy generally increases across a period. As we move across a period, Z eff increases. Therefore, it becomes more difficult to remove an electron. Two exceptions: removing the first p electron and removing the fourth p electron.

21. Chapter 7 Ionization Energy Periodic Trends in Ionization Energies The s electrons are more effective at shielding than p electrons. Therefore, forming the s 2 p 0 becomes more favorable. When a second electron is placed in a p orbital, the electron-electron repulsion increases. When this electron is removed, the resulting s 2 p 3 is more stable than the starting s 2 p 4 configuration. Therefore, there is a decrease in ionization energy.

24. Chapter 7 Ionization Energy Electron Configuration of Ions Cations: electrons removed from orbital with highest principle quantum number, n, first: Li (1s 2 2s 1 )  Li + (1s 2 ) Fe ([Ar]3d 6 4s 2 )  Fe 3+ ([Ar]3d 5 ) Anions: electrons added to the orbital with highest n: F (1s 2 2s 2 2p 5 )  F  (1s 2 2s 2 2p 6 )

25. Chapter 7 Electron Affinities Electron affinity is the opposite of ionization energy. Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion: Cl(g) + e -  Cl - (g) Electron affinity can either be exothermic (as the above example) or endothermic: Ar(g) + e -  Ar - (g)

26. Chapter 7 Electron Affinities Look at electron configurations to determine whether electron affinity is positive or negative. The extra electron in Ar needs to be placed in the 4s orbital which is significantly higher in energy than the 3p orbital.

27. Electron Affinities

28. Metals, Nonmetals, and Metalloids

29. Summation of Periodic Trends

30. Chapter 7 Metals, Nonmetals, and Metalloids Metals Metallic character refers to the properties of metals (shiny or lustrous, malleable and ductile, oxides form basic ionic solids, and tend to form cations in aqueous solution). Metallic character increases down a group. Metallic character decreases across a period. Metals have low ionization energies. Most neutral metals are oxidized rather than reduced.

31. Chapter 7 Metals, Nonmetals, and Metalloids Metals When metals are oxidized they tend to form characteristics cations. All group 1A metals form M + ions. All group 2A metals form M 2+ ions. Most transition metals have variable charges.

32. Prentice Hall © 2003Chapter 7 Metals, Nonmetals, and Metalloids Metals

33. Chapter 7 Metals, Nonmetals, and Metalloids Metals Most metal oxides are basic: Metal oxide + water  metal hydroxide Na 2 O(s) + H 2 O(l)  2NaOH(aq) Nonmetals Nonmetals are more diverse in their behavior than metals. When nonmetals react with metals, nonmetals tend to gain electrons: metal + nonmetal  salt 2Al(s) + 3Br 2 (l)  2AlBr 3 (s)

34. Chapter 7 Metals, Nonmetals, and Metalloids Nonmetals Most nonmetal oxides are acidic: nonmetal oxide + water  acid P 4 O 10 (s) + H 2 O(l)  4H 3 PO 4 (aq) Metalloids Metalloids have properties that are intermediate between metals and nonmetals. Example: Si has a metallic luster but it is brittle. Metalloids have found fame in the semiconductor industry.

35. Chapter 7 Group Trends for the Active Metals Group 1A: The Alkali Metals Alkali metals are all soft. Chemistry dominated by the loss of their single s electron: M  M + + e - Reactivity increases as we move down the group. Alkali metals react with water to form MOH and hydrogen gas: 2M(s) + 2H 2 O(l)  2MOH(aq) + H 2 (g)

36. Chapter 7 Group Trends for the Active Metals Group 1A: The Alkali Metals Alkali metal produce different oxides when reacting with O 2 : 4Li(s) + O 2 (g)  2Li 2 O(s)(oxide) 2Na(s) + O 2 (g)  Na 2 O 2 (s)(peroxide) K(s) + O 2 (g)  KO 2 (s)(superoxide) Alkali metals emit characteristic colors when placed in a high temperature flame. The s electron is excited by the flame and emits energy when it returns to the ground state.

37. Prentice Hall © 2003Chapter 7 Group Trends for the Active Metals Group 1A: The Alkali Metals

38. Prentice Hall © 2003Chapter 7 Group Trends for the Active Metals Group 2A: The Alkaline Earth Metals

39. Chapter 7 Group Trends for the Active Metals Group 2A: The Alkaline Earth Metals Alkaline earth metals are harder and more dense than the alkali metals. The chemistry is dominated by the loss of two s electrons: M  M 2+ + 2e -. Mg(s) + Cl 2 (g)  MgCl 2 (s) 2Mg(s) + O 2 (g)  2MgO(s) Be does not react with water. Mg will only react with steam. Ca onwards: Ca(s) + 2H 2 O(l)  Ca(OH) 2 (aq) + H 2 (g)

40. Chapter 7 Group Trends for Selected Nonmetals Hydrogen Hydrogen is a unique element. Most often occurs as a colorless diatomic gas, H 2. It can either gain another electron to form the hydride ion, H , or lose its electron to become H + : 2Na(s) + H 2 (g)  2NaH(s) 2H 2 (g) + O 2 (g)  2H 2 O(g) H + is a proton. The aqueous chemistry of hydrogen is dominated by H + (aq).

41. Prentice Hall © 2003Chapter 7 Group Trends for Selected Nonmetals Group 6A: The Oxygen Group

42. Chapter 7 Group Trends for Selected Nonmetals Group 7A: The Halogens The chemistry of the halogens is dominated by gaining an electron to form an anion: X 2 + 2e -  2X -. Fluorine is one of the most reactive substances known: 2F 2 (g) + 2H 2 O(l)  4HF(aq) + O 2 (g)  H = -758.7 kJ. All halogens consists of diatomic molecules, X 2.

43. Chapter 7 Group Trends for Selected Nonmetals Group 7A: The Halogens Chlorine is the most industrially useful halogen. It is produced by the electrolysis of brine (NaCl): 2NaCl(aq) + 2H 2 O(l)  2NaOH(aq) + H 2 (g) + Cl 2 (g). The reaction between chorine and water produces hypochlorous acid (HOCl) which disinfects pool water: Cl 2 (g) + H 2 O(l)  HCl(aq) + HOCl(aq). Hydrogen compounds of the halogens are all strong acids with the exception of HF.

44. Prentice Hall © 2003Chapter 7 Group Trends for Selected Nonmetals Group 8A: The Noble Gases

45. Chapter 7 Group Trends for Selected Nonmetals Group 8A: The Noble Gases These are all nonmetals and monatomic. They are notoriously unreactive because they have completely filled s and p sub-shells. In 1962 the first compound of the noble gases was prepared: XeF 2, XeF 4, and XeF 6. To date the only other noble gas compounds known are KrF 2 and HArF.