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Chapters 1 & 2

Chapters 1 & 2. General Chemistry Review Electronic Structure and Bonding. Molecular Representations. ~ 0.1 nm. Anders Jöns Ångström (1814-1874) 1 Å = 10 picometers = 0.1 nanometers = 10 -4 microns = 10 -8 centimeters. Nucleus = 1/10,000 of the atom. 1 nm = 10 Å

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Chapters 1 & 2

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  1. Chapters 1 & 2 General Chemistry Review Electronic Structure and Bonding Molecular Representations

  2. ~ 0.1 nm Anders Jöns Ångström (1814-1874) 1 Å = 10 picometers = 0.1 nanometers = 10-4 microns = 10-8 centimeters Nucleus = 1/10,000 of the atom • 1 nm = 10 Å • An atom vs. a nucleus • ~10,000 x larger

  3. Question 1.1 • What is the electronic configuration of carbon? • A) 1s2 2s2 2px2 • B) 1s2 2s2 2px1 2py12pz0 • C) 1s2 2s2 2px12py12pz1 • D) 1s2 1px1 1py12s2

  4. Electron Configurations Noble Gases and The Rule of Eight • When two nonmetalsreact to form a covalent bond: They share electrons to achieve a Noble gas electron configuration. • When anonmetaland a metal react to form an ionic compound: Valence electrons of the metalare lost and the nonmetal gains these electrons.

  5. Notes from Lewis’s notebook and his “Lewis” structure. G.N. Lewis Photo Bancroft Library, University of California/LBNL Image Library Footnote: G.N. Lewis, despite his insight and contributions to chemistry, was never awarded the Nobel prize. http://chemconnections.org/organic/Movies%20Org%20Flash/LewisDotStructures.swf

  6. Ionic Compounds • Ionic compounds are formed when electron(s) are transferred. • Electrons go from less electronegative element to the more electronegative forming ionic bonds. Worksheet 1: Bonds, Formulas, Structures & Shapes http://chemconnections.org/organic/chem226/226assign-12.html#Worksheets

  7. Covalent Compounds • Share electrons. • 1 pair = 1 bond. • Octet rule (“duet” for hydrogen) • Lewis structures: Notice the charges: In one case they balance, can you name the compound? In the other they do not, can you name the polyatomic ion? More about “formal” charge to come.

  8. Covalent Bonding – Simple Lewis Structures • For simple Lewis structures: • Draw the individual atoms using dots to represent the valence electrons. • Put the atoms together so they share PAIRSof electrons to make complete octets. • Take NH3, for example: • Practice with SKILLBUILDER 1.3.

  9. Question 1.2 • Select the correct Lewis structure for methyl fluoride (CH3F). • A) B) • C) D)

  10. Important Bond Numbers (Neutral Atoms / Normal electron distribution / Free electron pairs not shown) O two bonds three bonds N four bonds C one bond H F Cl Br I

  11. Question 1.3 • What is the best Lewis structure for formaldehyde (H2CO)? • A) B) • C) D)

  12. Question 1.4 • Which of the following contains a triple bond? • A) SO2 • B) HCN • C) C2H4 • D) NH3

  13. Formal Charge Formal charge is the charge of an atom in a Lewis structure which has a different than normal distribution of electrons. Klein: 2.4, 2.9

  14. O two bonds three bonds N four bonds C Important Bond Numbers (Neutral Atoms / Normal electron distribution) one bond H F Cl Br I

  15. Important Bond Numbers (Neutral Atoms / Normal electron distribution)

  16. Formal Charge Formal charge = number of valence electrons – (number of lone pair electrons +1/2 number of bonding electrons) • Equals the number of valence electrons (Group Number of the free atom) minus [the number of unshared valence electrons in the molecule + 1/2 the number of shared valence electrons in the molecule]. • Moving/Adding/Subtracting atoms and electrons.

  17. HNO3 Nitric Acid

  18. Complete the following table. It summarizes the formal charge on a (“central”) atom for the most important species in organic chemistry.  

  19. Formal Charge = # of valence e-s -[1/2(# of bonding e-s ) + # of non-bonding e-s ] =? =? Worksheet #1 & Worksheet #4 http://chemconnections.org/organic/chem226/226assign-12.html#Worksheets

  20. Question 1.5 • What is the formal charge of the carbon atom in the Lewis structure? • A) -1 • B) 0 • C) +1 • D) +2 C

  21. Question 1.6 What is the formal charge of the oxygen atom in the Lewis structure? A) -1 B) 0 C) +1 D) +2

  22. Resonance Klein: 2.7 – 2.12

  23. Resonance Resonance is a very important intellectual concept that was introduced by Linus Pauling in 1928 to explain experimental observations. Eg. SO2 Bond order  1.5 Bond length > double bond; Bond length< single bond TUTORIAL http://www.nku.edu/~russellk/tutorial/reson/resonance.html

  24. Resonance • Two or more Lewis structures may be legitimately written for certain compounds (or ions) that have double bonds and/or free pairs of non-bonded electrons • It is a mental exercise in “pushing” or moving electrons.

  25. Rules of Resonance • Step 1: The atoms must stay in the same position. Atom connectivity is the same in all resonance structures. Only electrons move. • NON-Example: The Lewis formulas below are not resonance forms. A hydrogen atom has changed position.

  26. Rules of Resonance • Step 2: Each contributing structure must have the same total number of electrons and the samenetcharge. • Example:All structures have 18 electrons and a net charge of 0.

  27. Rules of Resonance • Step 3: Calculate formal charges for each atom in each structure. • Example:None of the atoms possess a formal charge in this Lewis structure.

  28. Rules of Resonance • Step 4: Calculate formal charges for the second and third structures. • Example:These structures have formal charges. NOTE: They are less favorable Lewis structures.

  29. H H .. .. .. + – : : H C O N O H C O N O .. .. .. .. .. H H “Pushing” Electrons • same atomic positions • differ in electron positions more stable Lewis structure less stable Lewis structure

  30. H H .. .. .. + – : : H C O N O H C O N O .. .. .. .. .. H H “Pushing” Electrons • same atomic positions • differ in electron positions only more stable Lewis structure less stable Lewis structure

  31. Why use Resonance Structures? • Delocalization of electrons and charges between two or more atoms helps explain energetic stability and chemical reactivity. • Electrons in a single Lewis structure are insufficient to show electron delocalization. • A composite of all resonance forms more accurately depicts electron distribution. (HYBRID) NOTE: Resonance forms are not always evenly weighted. Some forms are better than others.

  32. + – •• •• O O O •• •• •• •• Resonance Example • Ozone (O3) • Lewis structure of ozone shows one double bond and one single bond Expect: one short bond and one long bond Reality: bonds are of equal length (128 pm)

  33. + – •• •• O O O •• •• •• •• – + + – •• •• •• O O O O O O •• •• •• •• •• •• •• •• •• Resonance Example • Ozone (O3) • Lewis structure of ozone shows one double bond and one single bond Resonance:

  34. – + + – •• •• •• O O O O O O •• •• •• •• •• •• •• •• •• Resonance Example • Ozone (O3) • Electrostatic potentialmap shows both endcarbons are equivalentwith respect to negativecharge. Middle carbonis positive.

  35. Detailed Resonance Examples

  36. Question 1.7 • Which resonance structure contributes more to the hybrid? • A) B)

  37. VSEPR Klein: 1.10

  38. VSEPR ModelValence Shell Electron Pair Repulsion

  39. VSEPR Model The molecular structure of a given atom is determined principally by minimizing electron pair (bonded &free) repulsions through maximizing separations. Some examples of minimizing interactions.

  40. Predicting a VSEPR Structure • 1. Draw Lewis structure. • 2. Put pairs as far apart as possible. • 3. Determine positions of atoms from the way electron pairs are shared. • 4. Determine the name of molecular structure from positions of the atoms.

  41. Orbital Geometry Molecular Geometry Bond Angle # of lone pairs Chem 226 Linear Linear Trigonal Planar Trigonal Planar Trigonal Planar Bent Tetrahedral Tetrahedral Tetrahedral Trigonal Pyramidal Tetrahedral Bent Trigonal Bipyramidal Trigonal Bipyramidal Trigonal Bipyramidal Seesaw Trigonal Bipyramidal T-shape Trigonal Bipyramidal Linear Octahedral Octahedral Octahedral Square Pyramidal Octahedral Square Planar 0 0 1 0 1 2 0 1 2 3 0 1 2

  42. Molecular Geometry –Summary Practice: SKILLBUILDER 1.8.

  43. Lewis Structures / VSEPR / Molecular Models • Computer Generated Models Ball and stick models of ammonia, water and methane. Worksheet 1: Bonds, Formulas, Structures & Shapes http://chemconnections.org/organic/chem226/226assign-12.html#Worksheets http://chemconnections.org/COT/organic1/VSEPR/

  44. CovalentBond PolarityMolecular PolarityDipole Moment Klein: 1.11

  45. Covalent Compounds • Equal sharing of electrons: nonpolar covalent bond, same electronegativity (e.g., H2) • Unequal sharing of electrons between atoms of different electronegativities: polar covalent bond (e.g., HF)

  46. Polar Covalent Bonds • Electrons tend to shift away from lower electronegativity atoms to higher electronegativity atoms. • The greater the difference in electronegativity, the more polar the bond. Practice: SKILLBUILDER 1.5

  47. Question 1.8 • Which of the following bonds is the most polar? • A) B) • C) D)

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