Chapter 4 The Periodic Table

1. Chapter 4The Periodic Table Chemistry I 5.0

2. 4-1 How are the elements organized? • Late 1700s – Only 30 elements were identified • Mid 1800s – 65 elements were now identified with the help of spectroscopy.

3. History of Periodic Table • J.W. Dobereiner: Organized the elements into groups with similar properties. • He called these groups triads. • The middle element is often the average of the other two.

4. Triads on the Periodic Table

5. History of the Periodic Table • J.A.R. Newlands • Law of octaves. He said that properties repeated every 8th element.

6. History of the Periodic Table • Mendeleev: Father of the Periodic Table • Organized the 1st periodic table according to increasing atomic mass and put elements with similar properties in the same groups.

7. Mendeleev (cont.) • He rearranged some elements out of atomic mass in order to keep them together with other elements with similar properties. He also left three blanks in his table and correctly identified the properties of these 3 unidentified elements that were later identified and match his predictions.

8. Mendeleev’s Work

10. History of the Periodic Table • Moseley • Each element has a certain amount of positive charge in the nucleus which are called protons. • Moseley reorganized the periodic table by atomic number.

11. Glenn Seaborg “Seaborgium” Sg #106 • Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII by pulling the “f-block” elements out to the bottom of the table. He was the principle or co-discoverer of 10 transuranium elements. He was awarded the Noble prize in 1951 and died in 1999.

12. History of the Periodic Table • The Periodic Law: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a regular (periodic) pattern. • Valence electrons: outermost electrons which are responsible for chemical bonding.

14. Valence Electrons in the Periodic Table

16. Periodic Law • Vertical Column – Group • Similar properties • A.K.A. Family Horizontal Row - Period

17. Properties of Metals, Nonmetals, and Metalloids • Metals: luster, malleable, ductile, high density, solids at room temp., good conductors, react with acids to produce hydrogen gas • Nonmetals: brittle, dull, neither ductile or malleable, nonreactive with acids, nonconductors • Semimetals (A.K.A. Metalloids): properties of both metals and nonmetals

18. Parts of the Periodic Table

19. Aluminum Metal Sulfur Nonmetal Silicon Metalloid

20. 1A 1 8A 18 3A 4A 5A 6A 7A 13 14 15 16 17 2A 2 Transition Group 3B 4B 5B 6B 7B 8B 9B 10B 1B 2B 3 4 5 6 7 8 9 10 11 12 Rare Earth Metals Lanthanide Series Actinide Series 1 = Alkali Metals and Hydrogen Group 16 = Oxygen or Chalogen Group 2 = Alkaline Earth Metals 15= Nitrogen Group 14 = Carbon Group 18 = Noble Gas Group 17 = Halogen Group 13 = Boron Group

21. Trends in the Periodic Table • Atomic Radius • The distance from the nucleus to the outermost electrons. • Atoms get larger going down a group and from right to left in a period.

23. Atomic Radii vs Atomic Number

24. Trends in the Periodic Table • Ionic Size • When atoms gain electrons, they become larger.

25. Trends in the Periodic Table • Ionic Size • When atoms lose electrons, they become smaller. • Ions become larger when you go down a group.

26. Relative Sizes of Positive & Negative Ions The sodium ion lost an electron, and therefore the positive protons in the nucleus exert a stronger pull on the remaining negative electrons, shrinking the orbitals. Thus positive ions are smaller than their atoms. The chloride ion gained an electron, and therefore the fewer positive protons in the nucleus exert a weaker pull on the extra negative electrons, increasing the size of the orbitals. Thus negative ions are larger than their atoms.

27. Trends in the Periodic Table • Ionization Energy • The energy needed to remove electrons from atoms. • Elements that do not want to lose their electrons have high I.E. • I.E. increases going up a group due to electron shielding. • I.E. increases going from left to right in a period.

29. Ionization Energy of the 1st 20 Elements

30. Ionization Energy vs. Atomic Number

31. D. Successive Ionization Energies: • Energy required to remove electrons beyond the 1st electron. • Ionization energies will increase for every electron removed. 3. Na [Ne]3s1 Na• 1st = ____ kJ 2nd = ____ kJ 4. Mg [Ne]3s2 Mg: 1st = ____ kJ 2nd = ____ kJ 3rd = ____kJ 5. Al [Ne]3s23p1 Al: 1st = ____kJ 2nd = ____kJ 3rd = ____kJ 4th = ___kJ 4560 496 738 7730 1450 11,600 1816 2744 577

32. Trends in the Periodic Table • Electronegativity • Reflects an atom’s ability to attract electrons in a chemical bond. • Increases going up a group. • Increases going from left to right in a period.

34. Trends in the Periodic Table • Electron Affinity • Energy change that occurs when an atom gains an electron. • General rule – Nonmetals have more negative electron affinities than metals (except for the Noble Gases)

35. A highly negative electron affinity attracts electrons. (nonmetals) • A positive electron affinity does not attract electrons. (metals)

36. Electron Affinity