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PERIODICITY PowerPoint Presentation
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PERIODICITY

PERIODICITY

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PERIODICITY

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  1. PERIODICITY

  2. Periodic Table • Dmitri Mendeleev developed the modern periodic table. Argued that element properties are periodic functions of their atomic weights. • We now know that element properties are periodic functions of their ATOMIC NUMBERS.

  3. Periods in the Periodic Table

  4. Groups in the Periodic Table

  5. Regions of the Periodic Table

  6. 6.1 Metals, Nonmetals, and Metalloids • Three classes of elements are metals, nonmetals, and metalloids. • Across a period, the properties of elements become less metallic and more nonmetallic.

  7. 6.1 Metals, Nonmetals, and Metalloids • Metals, Metalloids, and Nonmetals in the Periodic Table

  8. 6.1 Metals, Nonmetals, and Metalloids • Metals, Metalloids, and Nonmetals in the Periodic Table

  9. 6.1 Metals, Nonmetals, and Metalloids • Metals, Metalloids, and Nonmetals in the Periodic Table

  10. 6.1 Metals, Nonmetals, and Metalloids • Metals, Metalloids, and Nonmetals in the Periodic Table

  11. 6.1 Metals, Nonmetals, and Metalloids • Metals • Metals are good conductors of heat and electric current. • 80% of elements are metals. • Metals have a high luster, are ductile, and are malleable.

  12. 6.1 Metals, Nonmetals, and Metalloids • Uses of Iron, Copper, and Aluminum

  13. 6.1 Metals, Nonmetals, and Metalloids • Uses of Iron, Copper, and Aluminum

  14. 6.1 Metals, Nonmetals, and Metalloids • Uses of Iron, Copper, and Aluminum

  15. 6.1 Metals, Nonmetals, and Metalloids • Nonmetals • In general, nonmetals are poor conductors of heat and electric current. • Most nonmetals are gases at room temperature. • A few nonmetals are solids, such as sulfur and phosphorus. • One nonmetal, bromine, is a dark-red liquid.

  16. 6.1 Metals, Nonmetals, and Metalloids • Metalloids • A metalloid generally has properties that are similar to those of metals and nonmetals. • The behavior of a metalloid can be controlled by changing conditions.

  17. Element Abundance C O Al Si Fe http://www.webelements.com/webelements/elements/text/Si/geol.html

  18. Hydrogen Shuttle main engines use H2 and O2 The Hindenburg crash, May 1939.

  19. Group 1A: Alkali Metals Reaction of potassium + H2O Cutting sodium metal

  20. Group 2A: Alkaline Earth Metals Magnesium Magnesium oxide

  21. Calcium Carbonate—Limestone Champagne cave carved into chalk in France The Appian Way, Italy

  22. Group 3A: B, Al, Ga, In, Tl Aluminum Boron halides BF3 & BI3

  23. Gems & Minerals • Sapphire: Al2O3 with Fe3+ or Ti3+ impurity gives blue whereas V3+ gives violet. • Ruby: Al2O3 with Cr3+ impurity

  24. Group 4A: C, Si, Ge, Sn, Pb Quartz, SiO2 Diamond

  25. Group 5A: N, P, As, Sb, Bi White and red phosphorus Ammonia, NH3

  26. Phosphorus • Phosphorus first isolated by Brandt from urine, 1669

  27. Group 6A: O, S, Se, Te, Po Sulfuric acid dripping from snot-tite in cave in Mexico Sulfur from a volcano

  28. Group 7A: F, Cl, Br, I, At

  29. XeOF4 Group 8A: He, Ne, Ar, Kr, Xe, Rn • Lighter than air balloons • “Neon” signs

  30. Transition Elements Lanthanides and actinides Iron in air gives iron(III) oxide

  31. Electron Configurations of Groups

  32. Lithium Group 1A Atomic number = 3 1s22s1 ---> 3 total electrons

  33. Beryllium Group 2A Atomic number = 4 1s22s2 ---> 4 total electrons

  34. Boron Group 3A Atomic number = 5 1s2 2s2 2p1 ---> 5 total electrons

  35. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 ---> 6 total electrons

  36. Nitrogen Group 5A Atomic number = 7 1s2 2s2 2p3 ---> 7 total electrons

  37. Oxygen Group 6A Atomic number = 8 1s2 2s2 2p4 ---> 8 total electrons

  38. Fluorine Group 7A Atomic number = 9 1s2 2s2 2p5 ---> 9 total electrons

  39. Neon Group 8A Atomic number = 10 1s2 2s2 2p6 ---> 10 total electrons

  40. Colors of Transition Metal Compounds Nickel Cobalt Copper Zinc Iron

  41. PERIODIC TRENDS

  42. Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity • Electronegativity

  43. Effective Nuclear Charge, Z* • Explains why E(2s) < E(2p) • Z* is the nuclear charge experienced by the outermost electrons. Is the result of the nuclear attraction being blocked by the core electrons. Nuclear attraction increases with an increase in protons • Estimate Z* by --> [ Z - (no. core electrons) ] • Charge felt by 2s e- in Li Z* = 3 - 2 = 1 • Be Z* = 4 - 2 = 2 • B Z* = 5 - 2 = 3 and so on!

  44. Effective Nuclear Charge, Z* • Shielding effect remains constant across a period. As the nuclear attraction increases across the shielding effect is less effective. • Shielding effect increases down a group thus effectively blocking any increase in nuclear attraction. • Electrons with a higher quantum number have more kinetic energy and thus are less affected by the nuclear charge. • Each of these forces need to be accounted for in each trend.

  45. Effective Nuclear Charge, Z* • Atom Z* Experienced by Electrons in Valence Orbitals • Li +1.28 • Be ------- • B +2.58 • C +3.22 • N +3.85 • O +4.49 • F +5.13 Increase in Z* across a period

  46. Periodic Trend in the Reactivity of Alkali Metals with Water Lithium Sodium Potassium

  47. Atomic Size • Size goes UP on going down a group. • Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy.

  48. Atomic Size • Size goes UP on going down a group. • Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy.

  49. General Outline for Trends • Trend-define • Down a group • Nuclear attraction-define once • Trend, effect • Shielding effect-define once • Trend, effect • Kinetic energy-define once • Trend, effect • Across a period • Nuclear attraction • Trend, effect • Shielding effect • Trend, effect • Kinetic energy • Trend, effect

  50. Atomic Radius • Atomic radius is the distance from the nucleus to the valance electrons. • Nuclear attraction (the attraction of the protons in the nucleus on valance electrons) increases going down a group. This should pull the electrons in closer to the nucleus. • Shielding effect (the blocking of nuclear attractions by core electrons) Shielding effect increases down a group offsetting the increase in nuclear attraction. • Kinetic energy (the energy of valance electrons associated with principle energy levels) increases down a group allowing the valance electrons to orbit farther from the nucleus increasing atomic radius.