Chapter 6

1. Chapter 6 The Periodic Table & Periodic Law

2. Section 6.1—Development of the Modern Periodic Table • In the late 1790s, French scientist Antoine Lavoisier compiled a list of elements known at the time, 23 of them.

3. Section 6.1—Development of the Modern Periodic Table • In 1864, English chemist John Newlands proposed that the elements be arranged in increasing atomic mass, and if so, their properties repeated every 8th element.

4. Section 6.1—Development of the Modern Periodic Table • He called this the law of octaves, which was later disproved because it did not work for all of the elements, • After a few years passed, it was shown that he was basically correct; the properties of the elements do repeat in a periodic way.

5. In 1869, Russian chemist Dmitri Mendeleev arranged the elements by increasing atomic mass into columns with similar properties into the 1st periodic table.

6. Section 6.1—Development of the Modern Periodic Table • He left blank spaces for undiscovered elements of which he predicted scandium, gallium, & germanium.

7. Section 6.1—Development of the Modern Periodic Table • In 1913, English chemist Henry Moseley noted that if you arrange the elements in order of the number of protons they contain instead of atomic masses, they result in a clear periodic pattern of properties.

8. Section 6.1—Development of the Modern Periodic Table • The statement that there is a periodic repetition of chemical & physical properties of the elements when they are arranged by increasing atomic number is called the periodic law.

9. Section 6.1—Development of the Modern Periodic Table

10. Section 6.1—Development of the Modern Periodic Table • The boxes are arranged into a series of columns called groups, or families, and rows called periods.

11. The representative elements (the A group) are often referred to as the main group because they possess a wide range of chemical & physical properties. • The transition elements are the elements in the B group.

12. Section 6.1—Development of the Modern Periodic Table • There are 3 main classifications for the elements—metals, nonmetals, & metalloids.

13. Section 6.1—Development of the Modern Periodic Table • Metals—elements that are generally shiny when smooth & clean, solid at room temperature, & good conductors of heat & electricity. Most are ductile & malleable. • They also lose electrons

14. Section 6.1—Development of the Modern Periodic Table • Metals located on left of stair-step line(except for hydrogen) • The group 1A elements are known as the alkali metals • The group 2A elements are known as the alkaline earth metals • Both are chemically reactive, with the alkali being more reactive

15. The transition elements are divided into the transition metals & the inner transition metals • The inner transition metals are the lanthanide & actinide series that make up the bottom of the periodic chart.

17. Section 6.1—Development of the Modern Periodic Table • Nonmetals—are elements that are generally gases or brittle, dull-looking solids. They are poor conductors of heat & electricity. The nonmetals occupy the upper right side of the periodic table. • The only nonmetal that is liquid at room temp is bromine.

18. The highly reactive 7A group is known as the halogens. • The highly unreactive 8A group are known as the noble gases. • Gain e-

19. Nonmetals

20. Section 6.1—Development of the Modern Periodic Table • Metalloids—are elements with physical & chemical properties of both metals & nonmetals. • These elements border the stair-step line.

21. stop

22. Section 6.2—Classification of the Elements • The elements are organized by their electron configuration. • Atoms in the same group have similar chemical properties because they have the same number of valence electrons.

23. Section 6.2—Classification of the Elements • Group 1 elements have a configuration of s1. • Group 2 elements have a configuration of s2.

25. The energy level of an element’s valence electrons indicates the period on the periodic table in which it is found.

26. Section 6.2—Classification of the Elements • [Ar]4s23d104p1 Gallium belongs to the 4th period.

27. Section 6.2—Classification of the Elements • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 • Period? • 5 • Group? • 2A • Element? • Strontium

28. Examples. Give the element. • …4s2 • …6s1 • …2s2 • …1s2

29. Section 6.2—Classification of the Elements • Remember electron dot??? Does not apply to the transition elements. • Group 1  valence electrons = 1 • Group 2  valence electrons = 2 • Group 13  valence electrons = 3 • Group 14  valence electrons = 4 • Group 15  valence electrons = 5 • Group 16  valence electrons = 6 • Group 17  valence electrons = 7 • Group 18  valence electrons = 8 (exception is He)

31. How many valence e-? • Silicon? • Calcium? • Bromine? • Krypton? • Potassium? • Helium?

32. Section 6.2—Classification of the Elements • Remember the orbitals??? s, p, d, & f

33. Section 6.2—Classification of the Elements • s-block elements Group 1 & 2 elements & helium. • Valence electrons only in “s” orbitals

34. p-block elements Groups 13-18 (except He) • valence electrons are filling or have filled the “p” orbitals • Group 18  noble gases  s & p orbitals are completely filled & are very stable elements

35. Give the element • …3p1 • …4p6 • …5p2 • …6p2 • …3s2 • …3d10

36. Section 6.2—Classification of the Elements • d-block elements Groups 3-12, the transition metals • Period 4 is filling the 3d orbital • Period 5 is filling the 4d orbital • Period 6 is filling the 5d orbital

37. f-block elements inner transitional metals; lanthanide & actinide series • Period 6 is filling the 4f orbital • Period 7 is filling the 5f orbital

38. Give the element • …3d2 • …5d9 • …4d10 • …5f2 • …4f1 • …5f1

40. Section 6.3—Periodic Trends • Trends describe predictable changes in properties of the elements as you move across a period or down a group.

41. Atomic Radius (Atomic Size) • There is a general decrease in atomic radii as you move across a period. • There is an increase as you move down a group.

43. C or N, which is larger? • C • Ca or Mg, which is smaller? • Mg • Fr or Ra, which is larger? • Fr

44. Ionization energy is the amount of energy needed to remove an electron.

47. Section 6.3—Periodic Trends • A high ionization energy indicates the atom has a strong hold on its electrons.

48. As you go across a period, the ionization energy increases.

49. Section 6.3—Periodic Trends • As you go down a group, the ionization energy decreases.