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Bonding and Periodic Trends Bonding and Periodic PowerPoint Presentation
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Bonding and Periodic Trends Bonding and Periodic

Bonding and Periodic Trends Bonding and Periodic

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Bonding and Periodic Trends Bonding and Periodic

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  1. Bonding and Periodic Trends Bonding and Periodic

  2. Noble Gases • Noble gases (group 18 elements) are chemically stable because their last energy level is full. • All elements except Hydrogen and Helium want 8 electrons in their last energy level in order to be chemically stable like the noble gases. (Hydrogen & Helium want 2 electrons) • Noble gas configuration - Elements other than the Noble Gases will gain, lose, or share electrons in order to become chemically stable.

  3. Octet Rule In forming compounds, atoms gain, lose, or share one or more valence electrons in such a way that they achieve the electron configuration of the nearest noble gas in the periodic table. • Two types of Chemical Bonds • Ionic Bonds (metal + nonmetal)-transfer of electrons • Covalent Bonds (nonmetal + nonmetal)-sharing of electrons

  4. Ionic Bonding • Atoms lose or gain electrons to form ions • Cations are positive ions • metals generally form cations • Anions are negative ions • non-metals generally form anions • Ionic compounds are held together by electrostatics- the positive charge of the cation attracting the negative charge of the anion.

  5. Ionic Compound • METAL + NONMETAL • Important metal ions • group 1 - lose 1 electron (+1) • group 2 - lose 2 electrons (+2) • group 13 - lose 3 electrons (+3) • Transition metals also form ionic compound • Their behavior is less predictable • They usually lose electrons to form cations

  6. Important Nonmetal Ions • Group 17 gain 1 electron (-1) • Group 16 gain 2 electrons (-2) • Group 15 gain 3 electrons (-3) Main group elements gain or lose electrons so that they have 8 valence electrons. (Octet Rule).

  7. Oxidation Numbers • The number of electrons an atom gains or loses to have a full outer energy level. • The charge an atom acquires when it becomes an ion. • Includes a + or – sign following by a number.

  8. IONS • Atoms that have a charge because they have lost or gained electrons • Cation – has lost one or more electron – has positive charge • Anion – has gained one or more electron – has negative charge

  9. sodium ion Na+ e- e- e- e- 11p+ e- e- e- e- e- e- Formation of Cation sodium atom Na e- e- e- e- e- e- loss of one valence electron 11p+ e- e- e- e- e- e-

  10. chloride ion Cl1- e- e- e- e- e- e- e- e- 17p+ e- e- e- e- e- e- e- e- e- e- Formation of Anion e- chlorine atom Cl gain of one valence electron e- e- e- e- e- e- e- 17p+ e- e- e- e- e- e- e- e- e- e-

  11. Group 13 e e e e e e e e e Li Li+ F- F 64 152 60 136 Na Na+ Al Cl- Cl 50 95 143 99 186 181 K K+ Br- Br 114 133 195 227 Trends in Atomic and Ionic Size Metals Nonmetals Group 1 Group 17 Al3+ Cations are smaller than parent atoms Anions are larger than parent atoms

  12. 5

  13. Relative Size of Atoms

  14. Atomic Radius • One half the distance between the nuclei of identical atoms that are bonded together. • Trends in atomic radius is the same as trends in atomic size.

  15. Ionic Crystals • Ionic bonds form crystal shaped structures • Metal + Nonmetal • (+ )charged metal attracts (-) charge nonmetal • Ions packed close together in a crystal lattice • Cations are surrounded by anions, anions are surrounded by cations • Ionic crystals are brittle

  16. Lewis Dot Diagrams • Lewis Dot Diagrams show the valence electrons around an atom.

  17. Valence ElectronsMain (A) Group Elements 6 8 1 5 4 7 3 2 Valence electrons are the electrons in the atom’s highest numbered energy level. Group 18 Group 14 Group 17 Group 15 Group 1 Group 2 Group 16 Group 13

  18. Electron Dot Diagrams of Ions • Atoms • Metals form (+) ions • Nonmetals form (-) ions • Ions Na Na Cl Cl Cl

  19. Ionization Energy • 1st Ionization Energy – The energy required to remove one valence electron (outermost electron) • 2nd Ionization Energy – The energy required to remove the 2nd valence electron after the first has already been removed.

  20. Ionization Energies . Group 1 18 H 1312 He 2372 Symbol First Ionization Energy (kJ/mol) Mg 738 1 1 13 15 17 2 14 16 Li 520 Be 900 B 801 C 1086 N 1402 O 1314 F 1681 Ne 2081 2 2 Na 496 Mg 738 Al 578 Si 787 P 1012 S 1000 Cl 1251 Ar 1521 3 3 12 4 6 8 9 10 11 3 5 7 Period K 419 Ca 590 Sc 633 Ti 659 V 651 Cr 653 Mn 717 Fe 762 Co 760 Ni 737 Cu 746 Zn 906 Ga 579 Ge 762 As 947 Se 941 Br 1140 Kr 1351 4 4 Rb 403 Sr 550 Y 600 Zr 640 Nb 652 Mo 684 Tc 702 Ru 710 Rh 720 Pd 804 Ag 731 Cd 868 In 558 Sn 709 Sb 834 Te 869 I 1008 Xe 1170 5 5 * Cs 376 Ba 503 La 538 Hf 659 Ta 761 W 770 Re 760 Os 839 Ir 878 Pt 868 Au 890 Hg 1007 Tl 589 Pb 716 Bi 703 Po 812 At -- Rn 1038 6 6 y Uuu -- Uub -- Uut -- Uuq -- Uup -- Fr -- Ra 509 Ac 490 Rf -- Db -- Sg -- Bh -- Hs -- Mt -- Ds -- 7 * Ce 534 Pr 527 Nd 533 Pm 536 Sm 545 Eu 547 Gd 592 Tb 566 Dy 573 Ho 581 Er 589 Tm 597 Yb 603 Lu 523 Lanthanide series y Th 587 Pa 570 U 598 Np 600 Pu 585 Am 578 Cm 581 Bk 601 Cf 608 Es 619 Fm 627 Md 635 No 642 Lr -- Actinide series

  21. Factors Affecting Ionization Energy Nuclear Charge The larger the nuclear charge, the greater the ionization energy. Radius The greater the distance between the nucleus and the outer electrons of an atom, the less the ionization energy. Sublevel An electron from a full or half-full sublevel requires additional energy to be removed. Smoot, Price, Smith, Chemistry A Modern Course 1987, page 189

  22. Covalent Bonds • How do two identical atoms bond to form molecules such as H2, N2, or O2? • Neither atom is more likely than the other to transfer an electron • The two atoms have to SHARE electrons • NONMETAL + NONMETAL

  23. Pictorial Diagram of Water Molecule

  24. F Lewis Structures Continued… • Covalent Compounds - Normally the atom with the highest oxidation number, or fewest number of atoms, goes in the center of the structure. • Carbon is always in the center & Hydrogen is always on the outside of the structure. • Ex. PF3 P F F

  25. Drawing Lewis Structures • Arrange symbols with the appropriate element in the middle • Count up total number of valence electrons from each atom • Arrange electrons so that each atom is surrounded by 8 electrons (Hydrogen only gets 2) • You can only use the number of electrons available from your total count • Replace each shared pair with a dash

  26. O O O O O O double bond (2 pairs) N N N N N N N N triple bond (3 pairs) Electron sharing examples O2 Share until octet is complete. N2 octet complete

  27. Lewis Structures Practice • Cl2 • H2S • CO2 • H2O • CH4

  28. Cl Cl NaCl This is the formation of an ionic bond. - + Na Cl electron transfer and the formation of ions Cl2 This is the formation of a covalent bond. sharing of a pair of electrons and the formation of molecules

  29. Comparison of Bonding Types covalent ionic • No ions - molecules • Made from ions • Metal + Nonmetal • Nonmetal + Nonmetal • Non-conductive • Occasionally Flammable • Various Textures • Aqueous solutions & molten salts conductive • Rarely Burn • Hard & Brittle • Sharing of valence electrons • Transfer of valence electrons • Low Melting Point • High Melting Point • Many different appearances • Form Crystals

  30. Polar Bonds

  31. F F In a covalent bond of two identical atoms, the nucleus of each atom pulls on the bonding pair.

  32. F F Both atoms have equal pull, so the bonding pair is shared equally. This is called a nonpolar bond.

  33. H Cl If two different atoms share a bond, one will pull more strongly on the bonding electrons. This is called Polar Bonding.

  34. H Cl The bonding electrons carry negative charge.

  35. H Cl The closer they get to the chlorine atom, the more negative it gets. The farther they get from the hydrogen, the more positive it gets.

  36. + H Cl _ But the charge is only partial. Hydrogen has not lost the electrons as in the formation of an ion. There is an unequal sharing of electrons.

  37. The degree of sharing (equal to unequal) is determined by the electronegativity difference between the two atoms.

  38. Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling 1901 - 1994

  39. H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 K 0.8 Ca 1.0 Sc 1.3 Ti 1.5 V 1.6 Cr 1.6 Mn 1.5 Fe 1.8 Co 1.8 Ni 1.8 Cu 1.9 Zn 1.7 Ga 1.6 Ge 1.8 As 2.0 Se 2.4 Br 2.8 Rb 0.8 Sr 1.0 Y 1.2 Zr 1.4 Nb 1.6 Mo 1.8 Tc 1.9 Ru 2.2 Rh 2.2 Pd 2.2 Ag 1.9 Cd 1.7 In 1.7 Sn 1.8 Sb 1.9 Te 2.1 I 2.5 * Cs 0.7 Ba 0.9 La 1.1 Hf 1.3 Ta 1.5 W 1.7 Re 1.9 Os 2.2 Ir 2.2 Pt 2.2 Au 2.4 Hg 1.9 Tl 1.8 Pb 1.8 Bi 1.9 Po 2.0 At 2.2 y Fr 0.7 Ra 0.9 Ac 1.1 * Lanthanides: 1.1 - 1.3 y Actinides: 1.3 - 1.5 Below 1.0 2.0 - 2.4 1.0 - 1.4 2.5 - 2.9 1.5 - 1.9 3.0 - 4.0 Electronegativities 1A 8A 1 1 3A 5A 7A 2A 4A 6A 2 2 3 3 2B 4B 6B 8B 1B 3B 5B 7B Period 4 4 5 5 6 6 7 Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 373

  40. H At 2.2 2.1 Two atoms with a less than 0.3 difference in electronegativity will share fairly equally, resulting in no charge. NONPOLAR COVALENT BOND

  41. H Cl 3.0 2.1 Two atoms with a small difference (0.4 – 1.6) in electronegativity will share unequally,resulting in partial charge. POLAR COVALENT BOND

  42. 4.0 0.8 K F Two atoms with a large difference in electronegativity (1.7+) will result in a loss of an electron, resulting in a full charge. IONIC BOND

  43. A B A+ B- A B Increasing DEN Increasing polarity 100% covalent 100% ionic Bonding spectrum Transfer Nonpolar Covalent Compounds have less than 5% Ionic Character. Polar Covalent Compounds have 6 – 49% Ionic Character. Ionic Compounds have 50%+ Ionic Characteristics.

  44. Type of bond? – I, PC, or NC Electronegativity Table pg. 344 in Text TiO2 CH4 NaI CS2 CO2 KCl AlCl3 CsF HBr

  45. What does it mean to be reactive? • We will be describing elements according to their reactivity. • Elements that are reactive bond easily with other elements to make compounds. • Some elements are only found in nature bonded with other elements. • What makes an element reactive? • An incomplete valence electron level. • Periodic trends in reactivity

  46. Summary of Periodic Trends Atomic radius decreases Ionization energy increases Electronegativity increases Nuclear charge increases 1A 0 2A 3A 4A 6A 7A 5A Nuclear charge increases Atomic radius increases Ionization energy decreases Electronegativity decreases Ionic size (cations) Ionic size (anions) decreases decreases