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CHEMISTRY – Chapter 6

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  1. CHEMISTRY – Chapter 6 Chemical Bonding

  2. Chapter 6 – Chemical Bonding Objectives: • Define chemical bond. • Explain why most atoms form chemical bonds. • State the octet rule. • Discuss ionic, covalent, and metallic bonding. • Classify bonding type according to electronegativity differences. • Draw complete Lewis dot diagrams. • Properly assign bonds and show valence electrons for most compounds.

  3. Chemical Bonding • There are more than 100 elements • From these thousands of compounds are formed w/bonds • Chemical bond – a mutual electrical attraction btwn nuclei and valence e- of different atoms • Nucleus are + and e- are – • Two atoms coming together • Why do elements want to bond together? • To become stable • To complete their valence shell

  4. Lewis Dot Diagrams • The chemical symbol represents the kernel (center of the atoms) • Each dot represents a valence e- • Ions are put in brackets [Na]+

  5. Lewis Dot Diagrams (cont.) ex. Na Mg Al Si P S F Ar

  6. Octet Rule • The tendency of atoms to rearrange themselves into a stable octet (8 e-) • All atoms want to complete their “octet” and be configured like a Noble Gas. • Example: Neon • For this to happen the position of 1 or more e- in the valence shell is altered • Lost, gained, or shared

  7. Ionic Bonding • A bond formed between a metal and a non-metal • Metals lose e- • Non-metals gain e- • e- in valence shell of metals are TRANSFERRED to valence shell of non-metals • Electrostatic force of attraction btwn oppositely charged particles (+ and -) • Want a valence shell of s2p6 • Except He and H • Formula unit – simplest collection of atoms from which an ionic compound can be formed

  8. Formation of a sodium cation Sodium atom Sodium ion Na  – e Na + [Ne]2s1 [Ne] 11 p+ 11 p+ 11 e- 10 e- 01+

  9. Formation of a chloride anion unpaired electron completed octet 1- :Cl  + e :Cl:  [Ne] 2s2 2p5 [Ne] 2s2 2p6 9 p+ 9 p+ 9 e- 10 e- 0 1 - ionic charge

  10. Learning Check A. Number of valence electrons in aluminum 1) 1 e- 2) 2 e- 3) 3 e- B. Change in electrons for octet 1) lose 3e- 2) gain 3 e- 3) gain 5 e- C. Ionic charge of aluminum 1) 3- 2) 5- 3) 3+

  11. Ionic bonds • Bond resulting from electrostatic attraction btwn cations and anions • electrons are TRANSFERRED • + and – attract together • Cations – lose e- (+) • Anions – gain e- (-) • As a result of electrostatic attraction ex. NaCl, MgF2, KCl

  12. Examples K + Cl → K Cl → KCl Mg + F → F Mg F → MgF2

  13. Covalent Bonds • Bond resulting from the SHARING of e- pairs btwn 2 atoms • Orbital of e- from 1 element overlaps orbital of e- from another element • e- are SHARED

  14. Ionic bond Strong + and - High mp Because of strength Hard and brittle Conduct electricity in H2O Break into ions Covalent bond Weaker molecules Low mp B/o weak bonds Gases Do NOT conduct Don’t break up into ions Ionic bonds vs. Covalent bonds

  15. Consider Chloride

  16. HONC 1234 • Remember the HONC 1234 Rule • H = 1 bond • O = 2 bonds • N = 3 bonds • C = 4 bonds

  17. Polar-covalent • bond in which the atoms have an unequal attraction for the shared e- • One end is more + and the other end is more -

  18. Examples Nitrogen reacts with fluoride to form nitrogen trifluoride Carbon reacts with chloride to form carbon tetrafluoride

  19. Ionic, covalent, or polar? • ionic or covalent bonds can be estimated by calculating the difference in electronegativities • Difference of 1.7 or less is covalent • Difference of 0 to 0.3 are nonpolar covalent • Difference of 0.3 to 1.7 is polar covalent • Difference of higher than 1.7 is ionic

  20. Examples

  21. Metallic Bonding • A metal and a metal • Outer E level orbitals overlap and allow e- to roam freely throughout the entire metal • Sea of e-

  22. Characteristics of metallic bonds • Freedom of e- to move accounts for high electrical and thermal conductivity • Many orbitals, separated by small E differences allows metals to absorb a wide range of light frequencies • e- fall back to lower E levels and give off E as light • Shiny (luster)

  23. Characteristics of metallic bonds • Malleability – ability to be hammered or beaten into thin sheets • Ductility – ability to be drawn or pulled into a wire • This can happen b/c metallic bonding is the same in all directions • One plane of atoms can slide past another w/out breaking bonds • Strength varies w/nuclear charge and # of e- in the sea

  24. Polyatomic Ions • Charged group of covalently bonded atoms • An ion made up of more than one atom • Very strong covalent bonds hold them together • Behave as a SINGLE atom w/a charge ex. hydroxide OH- sulfate SO4-2 nitrate NO3- carbonate CO3-2 phosphate PO4-3 ammonium NH4+

  25. Lewis Dot Diagrams for molecules and polyatomic ions Steps • 1. Count total valence e- for all atoms • CO2 C = 4 e-(1) = 4 e- O = 6 e-(2) = 12 e- Total = 16 e- • 2. Count the # of octet e- each atom wants • All want 8 e- except H • C is 1 octet = 8 • O is 1 octet = 8(2) • Total = 24 e-

  26. Lewis Dot Diagrams for molecules and polyatomic ions • 3. Subtract valence e- from octet e- and this equals bonding e- • 4. Divide bonding e- by 2 • This gives you the number of bonds • 5. Draw atoms in the arrangement that they bond • 6. Find the lone pairs by subtracting bonding e- from valence e- • 7. Make sure that each e- is satisfied with an octet • Bonding: Lewis Dot Structures Practice

  27. Examples