6.1 - The Periodic Table: A History

1. 6.1 - The Periodic Table: A History http://www.woodrow.org/teachers/ci/1992/MENDELEEV.GIF

2. Jöns Jakob Berzelius 1828 Swedish chemist - developed a table of atomic weights - introduced letters to symbolize elements • made the task easier 33 elements known by 1800

3. Jöns Jakob Berzelius

4. Johann Döbereiner 1829 German Chemist Triads 53 known elements http://www.glogster.com/media/1/6/49/85/6498532.jpg

5. History of the Periodic Table • Döbereiner a) - described triads of elements (e.g. Cl, Br, I; Ca, Ba, Sr; Li, Na, K) - first indication that elements are related to one another - atomic mass is related to chemical properties – the mass of the center element was halfway between the masses of the other two elements, all three have similar properties

6. History of the Periodic Table

7. History of the Periodic Table 1848 57 elements

8. 1860 Karlsruhe Congress (big Chemistry Conference) Germany

9. John Newlands 1865 • English Chemist • Arranged elements by atomic mass • Described the “Rule of octaves” • 62 elements http://www.rsc.org/education/teachers/learnnet/periodictable/scientists/newlands.jpg

11. Lothar Meyer 1870 • German Chemist • Arranged elements based on atomic mass • Discovered periodic properties related to atomic volume • Established concept of valency http://www.chemistrydaily.com/chemistry/Lothar_Meyer

12. Meyer’s Data

13. It’s in the Cards Pre-Lab • Ionization energy = the amount of energy, in J or kJ, required to remove 1 electron from an atom in the gaseous state • Atomic radius = the distance between the nuclei of two adjacent atoms of the same kind, divided by 2, measured in pm • Melting point = the temperature at which a solid becomes a liquid, measured in oC

14. It’s in the Cards Pre-Lab • Average atomic mass = the weighted average of the masses of all known isotopes of an element, measured in amu (or g) • Density = ratio of mass divided by volume, g/mL or g/cm3 • Electronegativity = a measure of the relative ability of an atom to attract electrons in the context of a chemical bond, Paulings or none

15. Dmitri Mendeleev

16. Dmitri Mendeleev 1869 Russian chemist Wrote elements and properties on notecards Arranged by atomic mass and properties Noted repetition of properties every 8 or 18 elements http://anhso.net/data/69/X_kun/571478/mendeleev18371.jpg

17. Dmitri Mendeleev 1869 • Predicted properties of 3 elements! • eka-aluminum, eka-boron, eka-silicon • Problems: Ar/K, Te/I, Co/Ni • First element of each pair has greater atomic mass

18. Properties of Some Elements Predicted by Mendeleev

19. Properties of Some Elements Predicted by Mendeleev

20. Properties of Some Elements Predicted by Mendeleev

24. Review • Döbereiner 1829 • Arranged by atomic mass • Triads: [Cl Br I], [Ca Ba Sr], [Li Na K] • Newlands 1865 • Arranged by atomic mass • Rule of Octaves • Meyer 1870 • Arranged by atomic mass, periodic trend with atomic volume • Established concept of valency • Mendeleev • Arranged by atomic mass • Repetition every 8 or 18 elements • Predicted 3 elements not yet discovered: eka-aluminium - gallium, eka-silicon - germanium and eka-boron - scandium

25. Discovery of the Noble Gases1890s • Lord Rayleigh (physicist) and Sir William Ramsay (chemist) • 1894 - Argon “the lazy one”, discovered when Ramsay was trying to isolate nitrogen • 1895 - Helium – found on earth in uranium minerals (found in the sun in 1868) • 1898 - Neon “the new one” Krypton “the hidden one” Xenon “the alien one” • 1910 – Radon Properties: Largely unreactive 8 electrons in valence shell Low boiling and melting points

26. Nucleus discovered – 1910 Rutherford predicted that the charge of an atom is proportional to its mass

27. Henry Moseley 1913 • English Physicist • worked with Rutherford – was given the task of testing his prediction about charge vs. mass • Periodic Law: Properties of elements are periodic functions of their atomic numbers http://www.explicatorium.com/images/Personalidades/Henry_Moseley.jpg

28. History of the Periodic Table Ön of emitted X-rays corresponded to # protons  atomic number “Do other properties match atomic numbers?” Yes! • arranged the periodic table by atomic #’s, not mass

29. Law of Atomic Numbers - the properties of elements are periodic functions of their atomic numbers (not atomic mass)  corrected incorrect placement of cobalt and nickel, and iodine and tellurium

30. Glenn Seaborg 1940’s • American Scientist at UC Berkeley • Nobel Prize in Physics, 1951 • Discovered 7 elements beyond U • Developed actinide series and added it to PT • Seaborgium the only element publicly named after a living person

32. Letter to Seaborg Seaborgium Lawrencium Berkelium, Californium Americium

33. Trends of the Periodic Table “periodic” = repeating pattern • Overall theme = electrons’ positions relative to each other and the nucleus determine the following properties.

34. 6.3 - Periodic Trends

35. Trends of the Periodic Table “periodic” = repeating pattern Electron configuration (  reactivity and bonding) 1. Atomic radius 2. Ionization energy 3. Electronegativity

36. Periodic Trends The position of a valence electron and the ability to remove it from an atom are related to • the number of protons in the nucleus • the extent to which the valence electron is shielded from the positively-charged nucleus by the negatively-charged core electrons

37. 1. Atomic Radius Trend across a period: smaller • Add e- tovalence shell, add p+, stronger pull from nucleus draws e-’s closer. • Shielding effect is constant across period • Not as noticeable with heavier elements

38. Atomic Radius

39. Atomic Radius

40. Atomic Radius 1. Which groups and periods of elements are shown in the table of atomic radii? 2. In what unit is atomic radius measured? Express this unit in m. 3. What are the values of the smallest and largest atomic radii shown? What elements have these atomic radii? 4. What happens to atomic radii within a period as the atomic number increases? 5. What accounts for the trend in atomic radii within a period? 6. What happens to atomic radii within a group? 7. What accounts for the trend in atomic radii within a group? 8. a) Divide the atomic radius of Cs by the atomic radius of Li and round to 2 significant figures. Cs:Li b) Divide the atomic radius of Cs by the atomic radius of Rn and round to 2 significant figures. Cs:Rn c) Summarize your findings about a) and b) here:

41. Atomic Radius • Which groups and periods of elements are shown in the table of atomic radii? groups 1A-8A; periods 1-6 2. In what unit is atomic radius measured? pm Express this unit in m. 10-12 m • What are the values of the smallest and largest atomic radii shown? What elements have these atomic radii? 31 pm – helium; 265 pm - cesium • What happens to atomic radii within a period as the atomic number increases? The atomic radius of the elements within a period generally decreases as the atomic number of the elements increases. • What accounts for the trend in atomic radii within a period? With increasing atomic number, the increased positive charge of the nucleus pulls more strongly on the outermost electrons, pulling them closer to the nucleus. The size of the shield stays the same, so becomes less effective. Consequently, the atomic radius decreases. 6. What happens to atomic radii within a group? The atomic radius within a group generally increases as the atomic number of the elements increases. 7. What accounts for the trend in atomic radii within a group? With increasing atomic number, the increased pull by the larger positive charge of the nucleus is offset by the outer electrons’ larger orbitals and by shielding by inner electrons. Consequently, the atomic radius increases. 8. a) Divide the atomic radius of Cs by the atomic radius of Li and round to 2 significant figures. Cs:Li 1.7 X b) Divide the atomic radius of Cs by the atomic radius of Rn and round to 2 significant figures. Cs:Rn 1.9 X c) Summarize your findings about a) and b) here:

42. 2. Ionization Energy • the energy required to remove an electron from an atom in the gas phase (in J or kJ) • there is a series of ionization energies for each atom (since > 1 electron can be removed) • removing each subsequent electron requires more energy

43. Diagram from Document Camera

44. Ionization Energy

45. Successive Ionization Energies

46. Successive Ionization Energies 1. What happens to the values of the successive ionization energies of an element? 2. How is a jump in ionization energy related to the valence electrons of the element?

47. Successive Ionization Energies • What happens to the values of the successive ionization energies of an element? The values of the successive ionization energies increase. • How is a jump in ionization energy related to the valence electrons of the element? The jump occurs after the valence electrons have been removed.

48. First Ionization Energy

49. Ionization Energies 1. What is meant by first ionization energy? • Which element has the smallest first ionization energy? The largest? What are their values? • What generally happens to the first ionization energy of the elements within a period as the atomic number of the elements increases? • What accounts for the general trend in the first ionization energy of the elements within a period? • Based on the graph, rank the group 2A elements in periods 2-5 in decreasing order of first ionization energy. • What generally happens to the first ionization energy of the elements within a group as the atomic number of the elements increases? • What accounts for the general trend in the first ionization energy of the elements within a group?

50. Ionization Energies 1. What is meant by first ionization energy? First ionization energy is the energy required to remove the first electron from a gaseous atom. • Which element has the smallest first ionization energy? The largest? What are their values? rubidium – about 400 kJ/mol; helium – about 2375 kJ/mol • What generally happens to the first ionization energy of the elements within a period as the atomic number of the elements increases? The first ionization energy of the elements within a period generally increases as the atomic number of the elements increases. • What accounts for the general trend in the first ionization energy of the elements within a period? With increasing atomic number, the increased positive charge of the nucleus produces an increased hold on the valence electrons. Consequently, the first ionization energy increases. • Based on the graph, rank the group 2A elements in periods 1-5 in decreasing order of first ionization energy. beryllium, magnesium, calcium, strontium 6. What generally happens to the first ionization energy of the elements within a group as the atomic number of the elements increases? The first ionization energy of the elements within a group generally decreases as the atomic number of the elements increase. 7. What accounts for the general trend in the first ionization energy of the elements within a group?