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Chapter 9 Chemical Bonding I: Lewis Theory

Chapter 9 Chemical Bonding I: Lewis Theory. Read/Study: Chapter 9 MGC Homework: Due April 17, 2008 at 11:50 p.m. MGC Quiz: Due by April 20, 2008 at 11:50 p.m. Chapters 7, 8, and 9. Chapters 9 and 10. Chemical Bond Types Ionic Pure Covalent Polar Covalent Coordinate Covalent

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Chapter 9 Chemical Bonding I: Lewis Theory

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  1. Chapter 9 Chemical Bonding I: Lewis Theory Read/Study:Chapter 9 MGC Homework: Due April 17, 2008 at 11:50 p.m. MGC Quiz: Due by April 20, 2008 at 11:50 p.m. Chapters 7, 8, and 9

  2. Chapters 9 and 10 • Chemical Bond Types • Ionic • Pure Covalent • Polar Covalent • Coordinate Covalent • Metallic • Properties of Chemical Bonds • Bond Length • Bond Strength • Bond Angles • Bond Order

  3. Chapters 9 and 10 • Chemical Bond Types • Ionic • Pure Covalent • Polar Covalent • Coordinate Covalent • Metallic

  4. Your Basic Chemical Bonding Tool Kit • Lewis Symbols • The Octet Rule • Lewis Structures • Resonance Structures • Formal Charges Chapter 9 • Your Advanced Chemical Bonding Tool Kit • Valence Shell Electron Pair Repulsion • Valence Bond Theory (VB) • Molecular Orbital Theory (MO) Chapter 10

  5. Key Questions • Why do atoms combine to form compounds and molecules? • Why do they combine in definite proportions by mass? • Why do molecules have characteristic shapes? • Important Terms • Valence - The combining capacity of an element. • Chemical Bond - A “strong” link among atoms in a molecule or crystal.

  6. Overview of Bonding Types Ionic Bonding - An ionic bond is a chemical bond that results from an electrostatic attraction among oppositely charged ions in a compound. They form when electrons are transferred from one atom to another to form ions. Na [Ne] 3s1 + Cl [Ne] 3s2 3p5 Na [Ne]+ + Cl [Ne]- Cl- Chapter 9 Na+

  7. Overview of Bonding Types Covalent Bonding - A covalent bond is a chemical bond that results from a sharing of electrons among the atoms in a compound. e- e- + + + Energy H H Chapter 9 e- H2 + e- +

  8. Overview of Bonding Types Polar Covalent Bonding - A covalent bond that occurs when the atoms unequally share one or more pairs of electrons. This happens when the atoms have different electronegativities. e- e- e- e- + F H e- e- Energy e- e- e- e- Chapter 9 e- e- e- e- d+ d- H F e- e-

  9. Overview of Bonding Types Coordinate Covalent Bonding (Dative) - A dative bond is a covalent bond that occurs when the two shared electrons are donated to the bond bythe same atom. The donating atom is the donor or Lewis Base and the accepting atom is the acceptor or Lewis Acid. .. .. :F: F: :F: B:F: :F: .. .. .. .. .. .. :F: B:F: .. : .. .. .. .. .. :F: Donor .. .. Acceptor Tetrafluoroborate ion

  10. Overview of Bonding Types Metallic Bonding - The force of attraction that holds metal atoms together in a metallic lattice. It results from the fact that the valence electrons (outer shell electrons) are NOT bound to a particular atom but are free to be shared by all of the atoms - they are delocalized. + + + + + + + + + + + + + + + + + +

  11. Summary of Bonding Types Ionic Bonding Polar Covalent Bonding Covalent Bonding Electropositive + Electronegative Electronegative + Electronegative Metallic Bonding Electropositive + Electropositive

  12. Assignment: State whether each of the following has ionic, covalent, polar covalent, or metallic bonding: H2O NaI S4 HCl Polar Ionic Covalent Polar K NH4Cl Rb C (diamond) Metallic Dative Polar and ionic Metallic Covalent Network

  13. Your Basic Chemical Bonding Tool Kit • Lewis Symbols • The Octet Rule • Lewis Structures • Resonance Structures • Formal Charges Chapter 9 • Your Advanced Chemical Bonding Tool Kit • Valence Shell Electron Pair Repulsion • Valence Bond Theory (VB) • Molecular Orbital Theory (MO) Chapter 10

  14. Lewis Electron Dot Symbols A symbol for an atom or ion consisting of the chemical symbol for the element surrounded by a number of dots equal to the number of valence electrons in the atom. Alkaline Earth Metals!! H 1s1 H • Li [He] 2s1 Li • Na [Ne] 3s1Na • K [Ar] 4s1 K • Rb [Kr] 5s1 Rb • Cs [Xe] 6s1Cs • Be [He] 2s2Be: Mg [Ne] 3s2Mg: Ca [Ar] 4s2 Ca: Sr [Kr] 5s2 Sr: Ba [Xe] 6s2Ba:

  15. Lewis Electron Dot Symbols .. Assignment: F [He] 2s22p5:F: Cl [Ne] 3s23p5:Cl: Br [Ar] 4s24p5:Br: I [Kr] 5s25p5:I: . Draw the Lewis Symbols for the following - .. . .. C P Tl Ge As Po . .. . Used Primarily with Main Group Elements

  16. Ionic Bonding Ionic Bonding - An ionic bond is a chemical bond that results from an electrostatic attraction among oppositely charged ions in a compound. They form when electrons are transferred from one atom to another to form ions. Na [Ne] 4s1 + Cl [Ne] 3s2 3p5 Na[Ne]+ + Cl [Ne]- Cl- Na+

  17. Ionic Bonding - “What’s yours is mine.” “What’s mine you can have” 1. Metal Ion Formation - Metals in Groups 1 and 2 tend to give up one and two electrons, respectively, to form 1+ and 2+ ions that are isoelectronic with the preceding noble gas. Positive ions are called “cations”. Li• Li+ + e- [He] 2s1 [He]+ Sr2+ + 2 e- Sr: 2+ [Kr] 5s2 [Kr]

  18. 2. Negative Ion Formation - The atoms of non- metals in Groups 15, 16, and 17 tend to gain electrons as to fill their valence shell s and p orbitals. Thus, they become isoelectronic with the noble gas at the end of their respective period. .. : Br : + e- .. : Br : - . .. 2- .. : Se : .. : Se : + 2 e- .. . : P : + 3 e- .. : P : 3- ..

  19. 3. The Octet Rule - A “rule” expressing the tendency of some main group elements to obtain a total of 8 electrons in their valence shell. There are MANY exceptions! : Sn2+ + 2e- [Kr] 4d10 5s2 : Sn : [Kr] 4d10 5s2 5p2 : Tl+ + e- [Xe] 4f14 5d106s2 : Tl • [Xe] 4f14 5d106s2 6p1

  20. 4. Formation of Ionic Bonds • Transfer of electrons to form ions. • Electrostatic attraction among the ions to form an ionic crystal lattice. .. Ca : + : O : .. Ca2+ + : O : 2- .. - + - - 6:6 Coordination; Octahedral Arrangement - + - - -

  21. Electrolytes

  22. Covalent Bonding - “Share and share alike!” Covalent bonding involves the sharing of electrons. 1.Lewis Structures of Molecules: H · + · H H:H H · + · O · + · H H:O:H Octet .. .. .. .. Shared Lone Pairs Pairs The total number of electrons in a Lewis structure of a molecule is the sum of the valence electrons in the individual atoms.

  23. H .. H : : H:O:H C::C H:C:::C:H .. : : H H Single Bonds Double Bond Triple Bond H H H-O-H C=C H-CC-H H H • Find the number of electrons in the Lewis structure by adding up the valence electrons of all the atoms in the molecule or ion; add one extra electron for each negative charge; subtract one electron for each positive charge.

  24. 1 C @ 4 valence electrons = 4 V.E. 2 O @ 6 valence electrons = 12 V.E. Total Valence Electrons = 16 V.E. CO2 • Draw a skeletal structure of the molecule or ion by arranging atoms and putting one single bond between atoms that are bonded to each other. O-C-O or O:C:O • Distribute the remaining electrons so as to satisfy the octet rule as closely as possible. : O=C=O: .. ..

  25. .. But what about…. :O:::C:O: ??? .. Both structures obey the octet rule. Which one is RIGHT??? 2. Formal Charges - A positive, negative, or zero value assigned “formally” to the atoms in a Lewis structure; it is calculated for each atom using the following formula: Formal Charge = V.E. - [(unshared electrons) + 1/2(shared electrons)] Purpose - To provide a method for predicting the “best” Lewis structure for molecules with more than one bonding possibility.

  26. .. .. .. .. .. .. or :O=S-O: :O-S=O: .. .. FCOL = 6 - [6 + 1/2(2)] = 6 - [6 + 1] = 6 - 7 = -1 FCS = 6 - [2 + 1/2(6)] = 6 - [2 + 3] = 6 - 5 = +1 FCOR = 6 - [4 + 1/2(4)] = 6 - [4 + 2] = 6 - 6 = 0 .. .. .. .. .. .. :O-S=O: :O=S-O: .. .. -1 +1 +1 -1 Resonance Structures

  27. Rules for Using Formal Charges • The most stable structures have the least formal charge. • Structures in which adjacent atoms have formal charges of the same sign tend to be unstable. • Structures in which positive charges are on more electronegative atoms are not as stable. Since both of the SO2 structures are equivalent (not “identical”), they both contribute equally and should be shown as resonance forms.

  28. 3. Resonance Structures - Two or more equivalent • ways to depict the bonding in a molecule or ion; • two or more equivalent and legitimate Lewis • structures for the same molecule or ion. Assignment: Draw all appropriate resonance structures for the nitrate ion, NO3-, the carbonate ion, CO3-, and the sulfur trioxide molecule, SO3. Class Exercise:Construct appropriate Lewis structures for CO2, calculate formal charges for the atoms, and determine if resonance structures are important.

  29. 4. Exceptions to the Octet Rule • Atoms with more than 8 electrons - this is possible for elements in row 3 and beyond in the periodic table due to the fact that d-orbitals become available to handle addi- tional electrons. .. .. :Cl: :Cl: P Cl: :Cl: :Cl: .. :F: :F F: S :F F: :F: .. .. .. .. .. .. .. .. .. .. .. .. ..

  30. Atoms with fewer than 8 electrons - This occurs in “electron deficient” compounds. .. :F: B :F: :F: +1 - 1 :F: B :F: :F: .. .. .. .. .. .. .. .. :Cl Be Cl: :Cl = Be = Cl: .. .. +1 - 2 +1

  31. Free Radicals - An atom, molecule, or ion that contains an odd number of unpaired electron. .. .. .. :O - Cl - O: .. . ..

  32. Polar Covalent Bonding - “All are equal but • some are more ‘equal’ than others” • Bond Polarity - The result of an uneven charge distribution between two atomic nuclei that bonded to each other. It is due to the fact that different elements have different levels of attraction - electronegativity - for the electron pairs being shared. • Non-Polar Bond - A “pure” covalent bond wherein the two atoms sharing the electrons have identical attraction for them - they have the same electronegativity.

  33. .. .. H:H :Cl:Cl: :N:::N: .. .. Non-Polar Bonds (Pure Covalent Bonds) • Electronegativity - The general tendency of an atom or group of atoms to attract SHARED electrons. 1. Electronegativity Scales - Relative values assigned to the elements in the Periodic Table to represent their attraction for electrons in a CHEMICAL BOND. (This is NOT the same as ELECTRON AFFINITY.)

  34. 2. Electronegativity Differences - The greater the difference in electronegativity between two atoms that are bonded together, the more polar the bond will be. E.N. for Cl = 3.2* E.N. for H = 2.2 D E.N. for HCl = 1.0  +  - H - Cl *Pauling Scale • 3. Polar Molecules - A molecule is polar IF it has • one or more polar bonds AND is unsymmetrical • in charge distribution.

  35. Assignment: State whether or not each of the following molecules is polar: N2 H2O SF6 CCl4 SO2 Non-Polar Polar Non-Polar Non-Polar Polar

  36. Coordinate Covalent Bonding (Dative Bonding) - “I have plenty so you may share” This type of covalent bonding involves a sharing of two electrons that have BOTH been provided by only one of the atoms. H O: H+ H3O+ H .. H3N:BF3 H3N: BF3

  37. Chapter 10 Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory Read/Study:Chapter 10 MGC Homework: Due April 23, 2008 at 11:50 p.m. MGC Quiz: Due April 25, 2008 at 11:50 p.m.

  38. Your Basic Chemical Bonding Tool Kit • Lewis Symbols • The Octet Rule • Lewis Structures • Resonance Structures • Formal Charges Chapter 9 Chapter 10 • Your Advanced Chemical Bonding Tool Kit • Valence Shell Electron Pair Repulsion • Valence Bond Theory (VB) • Molecular Orbital Theory (MO)

  39. Valence Shell Electron Pair Repulsion : A model of chemical bonding that allows the shapes of molecules to be predicted by making the logical assumption that electron pairs in molecules tend to stay as far apart as possible. VSEPR is a sophisticated use of Lewis structures to determine the geometry of polyatomic ions and molecules.

  40. Central Atom - An atom in a molecule or ion that is bonded to two or more other atoms. • Molecules with one Central Atom 1. Write the Lewis structure. .. .. :O = C = O: 2. Count the VSEPR electron pairs on the • Central Atom. • A. Count each bond - single, double, or • triple - as ONEVSEPR pair.

  41. B. Count each lone pair as one • VSEPR pair. .. .. Total of 2 VSEPR pairs on Carbon! :O = C = O: • C. Place VSEPR pairs around the Central • Atom so that they are as far apart as • possible. : C : 180o

  42. D. Using ATOMS, not electron pairs, • determine the geometry. O - C - O Geometry is Linear! • E. In determining geometry, take into • consideration the electron pair re- • pulsions: • (1) Bond Pair - Bond Pair Repulsion • (2) Bond Pair - Lone Pair Repulsion • (3) Lone Pair - Lone Pair Repulsion BP-BP < BP-LP < LP-LP Greater Repulsion

  43. Geometry Types # VSEPR Pairs Geometry 2 Linear 3 Trigonal Planar 4 Tetrahedral 5 TrigonalBipyramidal 6 Octahedral

  44. Assignment: Determine the geometry of the following species: Ammonia - NH3 Sulfur Hexafluoride - SF6 Nitrate Ion - NO3- Water - H2O Ammonium Ion - NH4+ Sulfur Dioxide - SO2 Iodine Heptafluoride - IF7 Chlorite Ion - ClO2-

  45. Your Basic Chemical Bonding Tool Kit • Lewis Symbols • The Octet Rule • Lewis Structures • Resonance Structures • Formal Charges Chapter 9 Chapter 10 • Your Advanced Chemical Bonding Tool Kit • Valence Shell Electron Pair Repulsion • Valence Bond Theory (VB) • Molecular Orbital Theory (MO)

  46. Valence Bond Theory • 1. Limitations of Lewis Structures and VSEPR • Gives only information about geometry. • Is based on the “octet rule” which has many exceptions. • Cannot adequately explain bonding in species such as Li2 and H2+. • Does not reflect the QUANTUM nature of electrons.

  47. 2. Quantum Mechanical Theories of Bonding • Valence Bond Theory - Involves the over- lapping of atomic orbitals from THE SAME atom. • Molecular Orbital Theory - Involves the formation of MOLECULAR orbitals around two or MORE nuclei in a molecule. The molecular orbitals are formed by the over- lapping of atomic orbitals from DIFFERENT atoms.

  48. 3. Valence Bond Theory • A. The Simplest View - Bonds are formed by • the simple overlap of atomic orbitals from • two different atoms. Chapter 10 H H H 1s1 H-H s1s2 H 1s1 I 5 p1x I 5 p1x I-I s5p2

  49. A single bond consists of 2 electrons of opposite spin. The electrons are in a Sigma Bond (s). Sigma Bond - A bond resulting from the overlap of two atomic orbitals from DIFFERENT atoms, resulting in the build-up of electron density along the interatomic axis. +500 Repulsion H + H Energy 0 ATTRACTION H + H 74 pm; - 436 kJ/mol

  50. B. Orbital Hybridization Cl Be Cl [He] 2s2 [Ne] 3s2 3p5 [Ne] 3s2 3p5 How do we explain the bonding in BeCl2 using Valence Bond Theory? We must invoke….

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