1 / 41

Chapters 8

Chapters 8. Basic Concepts of Chemical Bonding. General Rule : All of the elements want to have the same number of electrons as their nearest noble gas neighbor. H He Li Be B ____ ____ ____ ____ ____ O F Ne Na Mg Al ____ ____ ____ ____ ____ ____ S Cl Ar K Ca

reginaa
Télécharger la présentation

Chapters 8

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapters 8 Basic Concepts of Chemical Bonding

  2. General Rule: All of the elements want to have the same number of electrons as their nearest noble gas neighbor. H He Li Be B ____ ____ ____ ____ ____ O F Ne Na Mg Al ____ ____ ____ ____ ____ ____ S Cl Ar K Ca ____ ____ ____ ____ ____ * * *

  3. Lewis symbols: The symbols represent the nucleus and all inner energy level e-. Dots represent valence e- (the outer energy level e-).

  4. Lewis structures: 1. Use the symbol to represent the nucleus and all inner energy level e- 2. Use dots to represent all valence e-, the outer energy level e-

  5. Octet Rule Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons. To achieve the same number of electrons as the noble gas closest to them in the Periodic Table. Except hydrogen, which has only the first energy level that can hold 2 electrons.

  6. Types of Bonding • Ionic Bonding – between a metal & non-metal Transfer of e- resulting in a positive ion and a negative ion that are attracted to each other. Metal – valence orbitals are emptied Non – metal – valence orbitals fill to have the configuration of the next Noble gas.

  7. Remember that you have learned many types of energy! NaCl(s)  Na(s) + ½ Cl2 (g)DHf = 411 KJ Na(s)  Na(g)DH = 108 KJ Na(g)  Na+1(g) + e-DH = 496 KJ ½ Cl2(g)  Cl(g)DH = 122 KJ Cl(g) + e-  Cl-1(g)DH = -349 KJ NaCl(s)  Na+1(g) + Cl-1(g)DHlattice = 788 KJ Ionization Energy e- Affinity When atoms are far apart, they DON’T interact.

  8. Lattice Energy is the change in the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.NaCl(s)  Na+1(g) + Cl-1(g)DHlattice = 788 KJ/mol Lattice Energy = k (Q1Q2/d )k = proportionality constant that is dependent on the structure of the solid and e- configurationQ1 and Q2 are charges on ionsd = distance between the centers of the ionsLattice energy increases: as the charges on the ions increase as their radii decreases.

  9. Transition metals lose valence e- from the s valence 1st, then enough d valence e- as required to reach the ionic charge. Ag = [Kr] 4d10 5s1 Fe = [Ar] 3d6 4s2 Fe = [Ar] 3d6 4s2 Ag+1 = [Kr] 4d10 Fe+2 = [Ar] 3d6 Fe+3 = [Ar] 3d5

  10. Covalent Bonding: between Non-metals A sharing of e- in such a way as to complete the valence orbitals of both (all) atoms. All atoms end up with the e- configuration of Noble gases.

  11. Remember that all atoms need to have an octet of e- when bonding except for hydrogen. It only has the first energy level which can hold only 2 e-.

  12. Bond polarity is a measure of how equal or unequal the e- in any covalent bond are shared. Nonpolar Bond = e- shared equally Polar Bond = one atom exerts a greater attraction for the e-

  13. Electronegativity (EN) The ability of an atom in a molecule to attract shared electrons to itself. Ranges from 0.7 to 4.0 Highest value F = 4.0 (upper right corner) Lowest value Fr = 0.7 (lower left corner) * Increases as move across the table  * Decreases as go down the table To determine is the e- are shared evenly between the bonding atoms, use:DEN = EN atom 1 – EN atom 2Use absolute valueDEN is less than 0.5 = Non – polar covalent bondDEN is 0.5 to 1.9 = Polar covalent bondDEN is greater than 1.9 = Ionic bond Page 310 

  14. 0.93.0 DEN = 3.0 – 0.9 = 2.1 (use absolute values) 2.1 = ionic bond

  15. Dipolar (or Dipole moment) Molecule with a center of positive charge and a center of negative charge. H – F: H - F: • d + d - • DEN = 4.0 – 2.1 = 1.9 a polar covalent bond • :O – H : O – H • H H • DEN = 3.5 – 2.1 = 1.4 a polar covalent bond or d - . . . . d +

  16. Dipole Moment – a measure of the charge of the dipole m = Q r If 2 equal but opposite charges Q+ and Q- are separated by a distance r, the magnitude of the dipole moment is the product measured in Debyes.

  17. Red represents an e- density therefore negative. Moving to the blue, which is positive due to a lack of e- density.

  18. The opposed bond polarities cancel out. There is no dipole moment, a nonpolar molecule.

  19. The bond polarities going in all directions cancel out. This is a non-polar molecule.

  20. Steps for determining covalent bonding: • Determine the total number of valence e- • Write the skeleton structure and join the atoms in this structure by a single covalent bond. • With the remaining valence e-, first complete the octets for all atoms attached to a central atom, then the central atoms. • Place any leftover e- on the central atoms. • 5. If there are not enough e- to give the cental atom an octet, make multiple bonds. • (Double bond = 4 e- or 2 pair of e-, Triple bond = 6 e- or 3 pair of e-)

  21. Formal charges: These are not actually charges on an atom, but an “accounting” procedure that helps to determine the most plausible Lewis structure. • Whenever formal charges are necessary to determine structures, these should be as small as possible. • Negative (-) formal charges should be on the most electronegative atom. • Positive (+) formal charges should be on the least electronegative atom. • Formal charges must be: Equal to zero for a neutral molecule. Equal to the ionic charge on a polyatomic ion.

  22. Formal charge accounting process Formal charge = Total number of valence e- minus (1/2 the bonding e-) minus (number of lone pair e-) Lone pair e-

  23. Resonance theory: when there are more than one plausible (valid) Lewis structure that can be drawn to describe the structure of a molecule, however the “correct” structure cannot. All of these are valid structures, none correct. These show 2 long single bonds and one short double bond. Actually, there are 3 bonds which appear the same length.

  24. Determine the correct resonant structureH – N = C: H – C = N :

  25. Write the correct Lewis structure for NOCl

  26. In polyatomic ions, the sum of the formal charges must equal the charge on the ion.

  27. Exceptions to the Octet Rule of bonding • Incomplete Octets: Limited to some Be, B, and Al compounds.

  28. Expanded Octets: Bonding with 10 or 12 e-Limited to Non-metals of period 3 (P or S) or greater. It is assumed that sulfur has 3s, and 3p full with the extra e- going into the empty 3d orbitals.

  29. PCl5 with 40 valence e-

  30. When it is necessary to exceed the octet rule, assume that the extra e- should be placed on the central atom.

  31. Odd e- spin: Free radicals are molecule fragments with 1 or more unpaired e-. These are paramagnetic.NO has 11 valence e-N = O

  32. Bond Energy • Bond breaking is an endothermic process (+DH) • Bond forming is an exothermic process (-DH) • DH = (S BE bond breaking) - (S BE bond forming)

  33. Page 326

  34. Using bond energy, calculate the DH associated with the following:CH4 (g) + Cl2 (g) CH3CL (g) + HCl (g)

  35. Longest = single bond Medium = double bond Shortest = triple bond Bond lengths are the distance between the center of the 2 nuclei of the bonding atoms. 1 angstrom = 100 pm

More Related