The Periodic Table

1. The Periodic Table Unit 3

2. Overview • Discovery of table • Introduction to table • States of matter, radioactivity, synthetics • Metals, nonmetals, metalloids • Location and properties • Periods versus groups • Group Location and Properties • Alkali metals, alkaline earth metals, transition metals, halogens, noble gases, rare earth elements • Periodic Trends • Effective nuclear charge • Atomic radius • Ionization energy • Electron affinity • Electronegativity • Ionic Radii

3. Discovery of the Elements

4. Early Element Tables (1778-1808)

5. Dmitri Mendeleev’s Periodic Table • 1869

6. Other tables following Mendeleev

7. Henry Moseley (1887-1915) • Periodic Law (atomic numbers)

8. Modern Periodic Table

9. 117 Known Elements • 83 are stable and found in nature. • Many of these a very rare. • 7 are found in nature but are radioactive. • 27+ are not natural on the earth. • 2 or 3 of these might be found in stars.

10. Why do we have the rows at the bottom? This arrangement takes too much space and is hard to read.

11. Radioactive Elements • Radioactive elements • Elements with an unstable nucleus • Spontaneously give off radiation and/or subatomic particles as nucleus “breaks apart” • Can occur naturally be man-made • Used in various forms of technology • Medicine, food preservation, weaponry, etc.

12. Man-made (Synthetic) Elements • Do not occur naturally on Earth • Artificially created • Unstable and decaying with half-lives between years and milliseconds • Made by bombarding any known element with a form of proton species such as helium ions or the use of proton bombardment via the use of a cyclotron to accelerate both of these types into another element to create other more heavy elements

13. Why make elements that last such a short time? • To support theories of the nature of matter. • The standard model of the nature of matter predicts that elements with roughly 184 neutrons and 114 protons would be fairly stable. • 288115, which lasted a relatively long time, has 115 protons and 173 neutrons. • The technology developed to make new elements is also being used for medical purposes. • Heavy-ion therapy as a treatment for inoperable cancers • Beams of carbon atoms shot at tumor.

14. Solid Liquid Gas States of Matter (Room Temperature) H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe * Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn + Fr Ra Lr * La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No +

15. He B C N O F Ne Si P S Cl Ar Ge As Se Br Kr At Rn Metals H Li Be Na Mg Al K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te Te Xe I * Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po + Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb * Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Ac +

16. Properties of Metals • Have a shiny metallic luster • Conduct heat well and conduct electric currents in the solid form • Malleable • For example, gold, Au, can be hammered into very thin sheets without breaking. • Ductile • Can be stretched into wiring without breaking • Solid at room temperature (except mercury, Hg)

17. Nonmetals H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe * Cs Ba Ly Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn + Fr Ra Lr Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb * La Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Ac +

18. Properties of Nonmetals • Dull • Poor conductors of heat and electricity • Brittle in the solid form • For example, if you hit a piece of sulfur, S, with a hammer it will shatter instead of flattening out • Not ductile • Cannot be stretched into wiring • Exist as solids, liquids, and gases at room temperature

19. Metalloids H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te Xe I * Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn + Fr Ra Lr * La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No +

20. Properties of Metalloids • Mixed properties of metals and nonmetals • Shiny or dull • More conductive than nonmetals but less than metals • Solid at room temperature

21. Periods are assigned numbers Periods H He 1 2 3 4 5 6 7 Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

22. A group or family I A II A III A IV A V A VI A VIIA 0 H He Groups are assigned Roman numerals with an A or B Li Be B C N O F Ne Na Mg Al Si P S Cl Ar III B IVB V B VIB VIIB VIII IB IIB K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

23. Group Numbers

24. Alkali metals Alkaline earth metals Halogens Noble gases Groups I A II A III A IV A V A VI A VIIA 0 H He Transition Metals Li Be B C N O F Ne Na Mg Al Si P S Cl Ar III B IVB V B VIB VIIB VIII B IB IIB K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr ( ) La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Rare Earth Elements Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

25. Groups • Elements in the same group have similar properties • Due to valence electrons being the same for each element in the group

26. Alkali Metals (Group 1) • Most reactive metals • Never found free in nature • React explosively with water • Extremely soft (can be cut with knife)

27. Alkaline Earth Metals (Group 2) • Less reactive than metals • Not found free in nature • Slightly react with water • Soft but not as soft as alkali metals • Often used in fireworks

28. Transition Metals (Groups 3-12) • Great conductors of heat/electricity • Hard metals • High melting/boiling points • Malleable and ductile

29. Halogens (Group 17) • Most reactive nonmetals • Elements exist as all three solid, liquid, and gas • Elements in pure form are diatomic (F2, Cl2, Br2, I2)

30. Noble Gases (Group 18) • Nonreactive (except for a few recently found compounds) • Usually only exist in free state in nature • Most are found in atmosphere in trace amounts) • All gases at room temperature • Colorless and odorless

31. Rare Earth Elements (Lanthanides/Actinides) • Usually silver, silvery-white, or gray metals • High luster but tarnish readily in air • Highly conductive • Many have similar or the same properties so difficult to distinguish one from the other • Occur naturally in minerals • Not particularly rare but prior to 1945 it was a long, tedious process to obtain pure samples of the elements

32. Effective Nuclear Charge • Many properties of atoms are due to the average distance of the outer electrons from the nucleus and to the effective nuclear charge, Zeff, experienced by these electrons. • Electrons are simultaneously attracted to nucleus and repelled by the other electrons. • Can estimate net attraction of each electron to nucleus by considering its interaction with average environment created by nucleus and other electrons • Inner electrons “shield” valence electrons from attraction of nucleus • Electrons in the same shell do not screen each other effectively

33. Effective Nuclear Charge • Interaction of charges: Zeff = Z – S • Zeff = effective nuclear charge • Z = number of protons in the nucleus • S = Shielding effect from other electrons • Zeffincreases moving left to right across a period • Zeffstays the same moving down a group

34. Atomic Radius • Half of the distance between nuclei in covalently bonded diatomic molecule • Radiusdecreasesacross a period • Increased effective nuclear charge due to decreased shielding • Radius increases down a group • Each row on the periodic table adds a “shell” or energy level to the atom

35. Atomic Radius

36. Ionization Energy • The energy required to remove an electron from an atom • Increases for successive electrons taken from the same atom due to the increased effective nuclear charge • The first ionization energy of an atom is the minimum energy needed to remove an electron from an atom • The second ionization energy is the energy needed to remove a second electron, • The third is to remove the third electron and so forth…

37. Ionization Energy • Tends to increase across a period • Electrons in the same quantum level do not shield as effectively as electrons in inner levels • Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove • Tends to decrease down a group • Outer electrons are farther from the nucleus and easier to remove

38. Trends in first ionization energy

39. Electron Affinity • The energy change associated with the addition of an electron • Affinity tends to increase across a period • Affinity tends to decrease as you go down in a period • Electrons farther from the nucleus experience less nuclear attraction • Some irregularities due to repulsive forces in the relatively small p orbitals

40. Electron Affinity

41. Electronegativity • A measure of the ability of an atom in a chemical compound to attract electrons • Electronegativity tends to increase across a period • As radius decreases, electrons get closer to the bonding atom’s nucleus • Electronegativity tends to decrease down a group or remain the same • As radius increases, electrons are farther from the bonding atom’s nucleus

42. Electronegativity

43. Summary of Periodic Trends

44. Ionic Radii • Cations • Positively charged ions form when an element loses an electron • Cation is smaller than corresponding atom • Anions • Negatively charged ions form when an element gains an electron • Anion is larger than corresponding atom