Evolution of the Periodic Table: From Mendeleev to Modern Design

1. The periodic law Chapter 5

2. Why do we need a table? • To organize the elements • To show trends

3. Periodic • A repeating pattern

4. Mendeleev’s table • 1869 – Dmitri Mendeleev – Russian • Arranged the elements in order of increasing mass and noticed that chemical properties were periodic • Put the elements into groups according to properties

5. Mendeleev vs. Meyer • 1860s Mendeleev and German Lothar Meyer each made an eight column table. • Mendeleev left some blanks in his table in order for all the columns to have similar properties – he predicted elements that hadn’t been discovered yet.

6. Why similar properties? • Why did they group according to properties and mass and not atomic number or number of outer level electrons?

7. Germanium • Mendeleev’s blank spots and his ability to predict future elements helped his table win acceptance.

8. Mendeleev’s table • Elements arranged in order of increasing mass. • Properties are repeated in an orderly, periodic, fashion. • Mendeleev’s periodic law – the properties of the elements are a periodic function of their masses.

10. Mass mistakes? • In order for Mendeleev to arrange his elements by properties, he had to put tellurium and iodine in the wrong order. • He explained this by assuming that their masses hadn’t been measured very accurately.

11. More mass mistakes? • Nickel and cobalt • Argon and potassium • Better mass measurements just confirmed the discrepancy

12. Explanation • 1913 – Henry Moseley • X-ray experiments revealed the atomic number was the number of protons • Modern periodic law – the properties of the elements are a periodic function of their atomic numbers

13. Modern periodic table • An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.

14. Noble gases • Not discovered on Earth until 1894 - 1900. • Group 18 was added to the table

15. Lanthanides • Hard to separate • All have similar properties • Added to the table in the early 1900s

16. Actinides • Discovered later • Also all have similar properties

17. Periodicity • Elements in the same group (column) have similar properties.

18. Chemical properties of an element • Are governed by the electron configuration of an atom’s highest energy level

19. Period length • Determined by the number of electrons than can occupy the sublevels being filled in that period. • Table 5-1

20. Full periodic table • Table with f-block in place

21. 1st period • 1s sublevel being filled • 1s can hold 2 electrons, so there are 2 elements in the 1st period.

22. 2nd and 3rd periods • 2s and 2p or 3s and 3p being filled • s and p sublevels can hold 8 total, so there are eight elements in these periods

23. 4th and 5th periods • Add d sublevels, which can hold 10 electrons • Need to fill 4s, 3d, and 4p – 18 electrons • 18 elements in each period

24. 6th and 7th periods • Add f-block, which holds 14 electrons • Fill 6s, 5d, 4f, 6p • Need 32 electrons • 32 elements in each period

25. Figure 5-5 • Shows blocks

26. Electron configurations • Elements in columns 1, 2, and 13-18 have their last electron added in an s or p orbital. • Elements in columns 3-12 have their last electron added in a d level.

27. The s-block elements: Groups 1 and 2 • Chemically reactive metals • Group 1 • Have 1 electron in outer s orbital • Coefficient represents period • Row 2: 2s1, Row 3: 3s1, etc. (ns1) • Group 2 • Have 2 electrons in outer s orbital • Coefficient represents period • Row 2: 2s2, Row 3: 3s2, etc. (ns2)

28. Alkali metals • Metals in group 1 • Have silvery appearance • Soft enough to cut with a knife • Not found alone in nature • React violently with nonmetals • Melting point decreases as you go down the table

29. Alkaline-earth metals • Group 2 • Harder, denser, and stronger than alkali metals • Higher melting points than alkalis • Less reactive • Not found alone in nature

30. Hydrogen and helium • Hydrogen • Located above group 1 because of its electron configuration • Not really in group 1, because its properties don’t match • Helium • Has an electron configuration like group 2 elements • In group 18 because it is unreactive

31. Discuss • Page 133 • Sample problem 5-1 and practice problems

32. Discuss • Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [Rn] 7s1 is located. • Group 1, 7th period, s block • Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [He]2s2 is located. • Group 2, second period, s block

33. d-block elements: Groups 3-12 • End in d1 to d10. • Coefficients are one less than the period • Example: Fe is in the 6th column of transition elements in the 4th period, ends in 3d6

34. Transition elements • Groups 3-12 • Typical metallic properties • Good conductors • High luster • Less reactive than alkalis and alkaline-earths • Some are unreactive enough to appear in nature

35. p-block elements: groups 13-18 • End in p1 to p6. • Coefficients are the same as the period • ns2np1 • Always have a full s-sublevel

36. p-block elements • Properties vary greatly • Includes all nonmetals except hydrogen and helium • Solids, liquids and gases • Includes all the metalloids • Between metals and nonmetals • Brittle solids • Semiconductors – can conduct under certain conditions • Includes some metals • Less reactive than alkalis and alkaline-earths

37. Halogens • Group 17 • Most reactive nonmetals • Form compounds called salts

38. f-block elements • Lanthanides and actinides • Endings are f1 to f14 • Coefficients are two less than the period • All actinides are radioactive • Those after neptunium are synthetic

39. Discuss • Sample problems and practice problems on pages 136, 138, and 139 • With your group first, then join with another group. • Do you have any questions?

40. Atomic radius • Ideally, the distance from the center of the atom to the edge of it’s orbital. • But, atoms are “fuzzy”, not clearly defined. • Defined as one-half the distance between the nuclei of identical atoms that are bonded together.

41. Period trends – see figure 5-13 • As we move from left to right across the table, we gain protons. • There is a greater positive charge on the nucleus. • This greater charge pulls harder on the outer electrons, pulling them in closer. • The atom gets smaller.

42. Group trends • As we move down the table, the principle quantum number increases. • When the principle quantum number increases, the electron cloud gets bigger. • The size of the atoms gets bigger.

43. Discuss • Which of the elements Li, Rb, K, and Na has the smallest atomic radius? Why? • Li, it is highest on the table • Which of the elements Zr, Rb, Mo, and Ru has the largest atomic radius? Why? • Rb, it is farthest to the left on the table

44. Ion • An atom or group of bonded atoms that has a positive or negative charge

45. Ionization • Any process that makes ions

46. Ionization energy (IE) • First ionization energy (IE1) – the energy required to remove the most loosely held electron. • Measured in kJ/mol

47. Ionization energy – see figure 5-15 • Experimentally determined. • From isolated atoms in the gas phase • Tends to increase as you move across a row from left to right • Why group 1 is most reactive • Caused by higher charge • Tends to decrease as you move down a column • Electrons are farther from nucleus • Shielding from inner electrons

48. Other Ionization Energies – see Table 5-3 • Energy required to remove other electrons from positive ions. • IE2, IE3, etc • Get higher as you remove more electrons • Less shielding

49. Noble Gases • Have High ionization energies • When a positive ion of another element reaches a noble gas configuration, its ionization energy goes up. • Example: When K loses one electron, it has Ar’s electron configuration • This makes it stable • Its IE2 is much higher than its IE1

50. Discuss • State in words the general trends in ionization energies down a group and across a period of the periodic table.