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Chemical Bonding I: Basic Concepts

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  1. Chemical Bonding I:Basic Concepts Chapter 9 – (Topic 4 and 14)

  2. 4.1 – Ionic Bonding

  3. Group e- configuration # of valence e- ns1 1 1A 2A ns2 2 3A ns2np1 3 4A ns2np2 4 5A ns2np3 5 6A ns2np4 6 7A ns2np5 7 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that particpate in chemical bonding. 9.1

  4. Lewis Dot Symbols for the Representative Elements & Noble Gases 9.1

  5. Why do substances bond? • More stability • Atoms want to achieve a lower energy state

  6. Ionic Bonding • Between a metal and a non-metal. • Metals lose electrons becoming a cations, while non-metals gain electrons becoming anions. • An ionic bond is an electrostatic attraction between the oppositely charged ions.

  7. - - - - + Li+ Li Li Li+ + e- e- + Li+ Li+ + F F F F F F The Ionic Bond [He] [Ne] 1s22s1 1s22s22p5 1s2 1s22s22p6 9.2

  8. E = k Q+Q- r lattice energy cmpd MgF2 2957 Q= +2,-1 MgO 3938 Q= +2,-2 LiF 1036 LiCl 853 Electrostatic (Lattice) Energy Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. Q+ is the charge on the cation Q- is the charge on the anion r is the distance between the ions Lattice energy (E) increases as Q increases and/or as r decreases. r F- < r Cl- 9.3

  9. 9.3

  10. Ionic Structures • In an ionic compound (solid), the ions are packed together into a repeating array called a crystal lattice. • The simplest arrangement is one in which the spheres in the base are packed side by side. Opposite charges are attracted to each other. • Its called simple cubic packing (NaCl is an example)

  11. 4.5 Physical Properties

  12. General physical properties • Depend on the forces between the particles • The stronger the bonding between the particles, the higher the M.P and BP • MP tends to depend on the existence of a regular lattice structure

  13. Impurities and Melting points • An impurity disrupts the regular lattice that its particle adopts in the solid state, so it weakens the bonding. • They always LOWER melting points • Its often used to check purity of a known molecular covalent compound because its MP will be off, proving its contamination

  14. 4.2 – Covalent Bonding

  15. Why should two atoms share electrons? + 8e- 7e- 7e- 8e- F F F F F F F F lonepairs lonepairs single covalent bond single covalent bond lonepairs lonepairs A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Lewis structure of F2 9.4

  16. single covalent bonds H H H H or H H O 2e- 2e- O 8e- O C O C O O double bonds 8e- 8e- 8e- double bonds O N N triple bond N N triple bond 8e- 8e- Lewis structure of water + + Double bond – two atoms share two pairs of electrons or Triple bond – two atoms share three pairs of electrons or 9.4

  17. Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond 9.4

  18. 9.4

  19. F H F H Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms (electrons are shared unequally) electron rich region electron poor region e- poor e- rich d+ d- 9.5

  20. F H Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electronegativity - relative, F is highest electron rich region electron poor region 9.5

  21. The Electronegativities of Common Elements 9.5

  22. Increasing difference in electronegativity Nonpolar Covalent Polar Covalent Ionic partial transfer of e- Unequal sharing share e- equally transfer e- Classification of bonds by difference in electronegativity Difference Bond Type 0 Nonpolar Covalent  2 Ionic 0 < and <2 Polar Covalent 9.5

  23. Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 NonPolar Covalent 9.5

  24. DH0 = 436.4 kJ H2 (g) H (g) + H (g) DH0 = 242.7 kJ Cl2 (g) Cl (g) + Cl (g) DH0 = 431.9 kJ HCl (g) H (g) + Cl (g) DH0 = 498.7 kJ O2 (g) O (g) + O (g) O DH0 = 941.4 kJ N2 (g) N (g) + N (g) O Bond Energies Single bond < Double bond < Triple bond N N The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy 9.10

  25. Coordinate Covalent or Dative Bond • A covalent bond in which one of the atoms donates both electrons. • Properties do not differ from those of a normal covalent bond.

  26. Coordinate covalent bonds (dative) • A covalent bond that occurs between two atoms in which both electrons shared in the bond come from the same atom. • Both electrons from the nitrogen are shared with the upper hydrogen • Ammonium has 3 polar covalent bonds and 1 coordinate (dative) covalent bond.

  27. Hydronium (H3O+) Carbon monoxide (CO) Examples

  28. Writing Lewis Structures • Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. • Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. • Complete an octet for all atoms except hydrogen • If structure contains too many electrons, form double and triple bonds on central atom as needed. 9.6

  29. Write the Lewis structure of nitrogen trifluoride (NF3). F N F F Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons 9.6

  30. Write the Lewis structure of the carbonate ion (CO32-). O C O 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 O Total = 24 Step 1 – C is less electronegative than O, put C in center • Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) • -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e- 9.6

  31. What are the resonance structures of the carbonate (CO32-) ion? - - + + O O O O O O O O O C C C O O O - - - - O O O - - A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. 9.8

  32. Resonance is possible whenever a Lewis structure has a multiple bond and an adjacent atom with at least one lone pair. • The following is the general form for resonance in a structure of this type.

  33. Be – 2e- 2H – 2x1e- H Be H 4e- B – 3e- 3 single bonds (3x2) = 6 3F – 3x7e- F B F 24e- Total = 24 9 lone pairs (9x2) = 18 F Exceptions to the Octet Rule The Incomplete Octet BeH2 BF3 9.9

  34. N – 5e- S – 6e- N O 6F – 42e- O – 6e- 48e- 11e- F 6 single bonds (6x2) = 12 F F Total = 48 S 18 lone pairs (18x2) = 36 F F F Exceptions to the Octet Rule Odd-Electron Molecules (radicals -very reactive) NO The Expanded Octet (central atom with principal quantum number n > 2) SF6 9.9

  35. # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry Class linear linear B B Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. AB2 2 0 10.1

  36. Cl 0 lone pairs on central atom Be Cl 2 atoms bonded to central atom 10.1

  37. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar Arrangement ofelectron pairs Molecular Geometry Class VSEPR AB2 2 0 linear linear AB3 3 0 10.1

  38. 10.1

  39. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement ofelectron pairs Molecular Geometry Class tetrahedral tetrahedral VSEPR AB2 2 0 linear linear AB4 4 0 10.1

  40. 10.1

  41. 14.1 - Molecules with more than 4 electron pairs • Molecules with more than 8 valence electrons [expanded valence shell] • Form when an atom can ‘promote’ one of more electron from a doubly filled s- or p-orbital into an unfilled low energy d-orbital • Only in period 3 or higher because that is where unused d-orbitals begin

  42. Why does this ‘promotion’ occur? • When atoms absorb energy (heat, electricity, etc…)their electrons become excited and move from a lower energy level orbital to a slightly higher one. • How many new bonding sites formed depends on how many valence electrons are excited.

  43. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement ofelectron pairs Molecular Geometry Class trigonal bipyramidal trigonal bipyramidal VSEPR AB2 2 0 linear linear tetrahedral tetrahedral AB4 4 0 AB5 5 0 10.1

  44. 10.1

  45. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement ofelectron pairs Molecular Geometry Class trigonal bipyramidal trigonal bipyramidal AB5 5 0 octahedral octahedral VSEPR AB2 2 0 linear linear tetrahedral tetrahedral AB4 4 0 AB6 6 0 10.1

  46. 10.1

  47. 10.1

  48. lone-pair vs. lone pair repulsion lone-pair vs. bonding pair repulsion bonding-pair vs. bonding pair repulsion > >

  49. # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry Class trigonal planar bent VSEPR trigonal planar trigonal planar AB3 3 0 AB2E 2 1 10.1

  50. # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry Class trigonal pyramidal tetrahedral VSEPR tetrahedral tetrahedral AB4 4 0 AB3E 3 1 10.1