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Chapter 8 Bonding and Molecular Structure

Chapter 8 Bonding and Molecular Structure. Chemical Bonding. attractive force holding two or more atoms together. Chemical bond :. Covalent bond:. a sharing electrons between the atoms. non–metal / non–metal bonds.

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Chapter 8 Bonding and Molecular Structure

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  1. Chapter 8Bonding and Molecular Structure

  2. Chemical Bonding attractive force holding two or more atoms together Chemical bond: Covalent bond: a sharing electrons between the atoms. non–metal / non–metal bonds. electrostatic in nature, transfer of electrons from a metal to a nonmetal. (cation + anion) Ionic bond: Metallic bond: attractive force holding pure metals together. Cations in a “sea” of electrons.

  3. Two Extreme Forms of Connecting or Bonding Atoms Ionic Bonding: complete transfer of 1 or more electrons from one atom to another. (metals and non-metals) Covalent Bonding:valence electrons shared between atoms. (non-metals and non-metals) Most bonds are somewhere in between.

  4. Ionic Bonding e– Ionic Bondingresults when an electron or electrons are transferred from one atom to another. The transfer results in each attaining an octet or Noble gas electron configuration. 3s1 3s23p5 Na + Cl  Na+ Cl 2s22p6 3s23p6 Na+ Cl- [Ne] [Ar] Noble gas electron configurations 

  5. Covalent Bonding • There are many examples of compounds having covalent bonds, including the gases in our atmosphere (O2, N2, H2O, and CO2), common fuels (CH4), and most of the compounds in your body. • Covalent bonding is also responsible for the atom-atom connections in polyatomic ions.

  6. Covalent Bond Formation A covalent bond results from a overlapof valence orbitals on neighboring atoms. Valence Electrons and Lewis Symbols for Atoms:

  7. Covalent Bonding & Lewis Structures • The American chemist Gilbert Newton Lewis (1875–1946) introduced a useful way to represent the valence shell electrons of an atom. • The element’s symbol represents the nucleus including the core electrons. • Up to four valence electrons, indicated by dots, are placed one at a time around the symbol. • If any valence electrons remain, they are paired with ones already there or shared electrons up to a total of eight. G. N. Lewis 1875 - 1946

  8. Lewis Dot Symbols Notice that the group numbers of the main group elements represents the number of valence electrons (dots) for each element.

  9. Valence Electrons • Electrons are divided between core and valence electrons B 1s2 2s2 2p1 • Core = [He] , valence = 2s2 2p1 Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10, valence = 4s2 4p5

  10. Lewis Dot Structures & the Octet Rule H H H • Covalent bonding, a bond results when one or more electron pairs are shared between two atoms. • The electron pair bond between the two atoms of an H2 molecule is represented by a pair of dots or a line. • The representation of a molecule in this fashion is called a Lewis electron dot structureor just a Lewis structurein honor of G. N. Lewis. H-atom: H2 molecule: H–H

  11. Lewis Dot Structures & the Octet Rule • The number of unpaired valence electrons gives a general indication as to the number of bonds an atom will likely form. • Hydrogen has only one electron and therefore can only make one covalent bond. • Gr 7A has only one unpaired electron, so it generally forms one bond. • Gr 6A had two unpaired electrons, thus the likelihood of two bonds. ... and so on.

  12. Lewis Dot Structures & the Octet Rule H F H F H– “single bond” When other covalent species form, there are additional electron pairs that do not participate in bonding. These are called “lone pairs” (lp) F  + one bonding pair & three lone pairs (octet) hydrogen fluoride: HF

  13. The Octet Rule • No. of valence electrons of an atom = Group number • For Groups 1A-4A, no. of bond pairs = group number • For Groups 5A -7A, BP’s = 8 - Grp. No. • Hydrogen can only form one bond and never has lone pairs. • Boron often forms only 3 bonds. • Elements in the 3rd period and beyond can exceed the octet rule. • Some stable molecules can form with an odd number of electrons.

  14. N N           : : N N Triple bond! Example: N2 Each nitrogen atom needs 8 electrons to complete an octet… But each nitrogen has only 5 valence electrons! As a result, the electrons must be shared. There are 10 electrons available           N2 needs 16 electrons (28) There are 10 valence electrons present Therefore 6 electrons must be shared. The remaining 4 electrons are the lone pairs

  15. Building a Lewis Dot Structure Ammonia, NH3 1. Decide on the central atom; the central atom is generally the atom of lowest affinity for electrons & never hydrogen. Therefore, N is the central atom. 2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs

  16. Building a Lewis Dot Structure H H N H •• H H N H 3. Form a single bonds between the central atom and each surrounding atom. 4. The remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair.

  17. Sulfite Ion, SO32− Step 1. Central atom = S Step 2. Count valence electrons: S = 6 3 x O = 3 x 6 = 18 Negative charge = 2 TOTAL = 26 e- or 13 pairs Step 3. Form bonds 10 pairs of electrons are now left.

  18. Sulfite Ion, SO32− •• O • • • • •• •• O S O • • • • •• •• • The remaining pairs become lone pairs, first on outside atoms and then on central atom. •• Each atom is surrounded by an octet of electrons.

  19. Carbon Dioxide, CO2 1. Central atom = _______ 2. Valence electrons = __ or __ pairs 3. Form bonds. This leaves 6 pairs. 4. Place lone pairs on outer atoms.

  20. Carbon Dioxide, CO2 4. Place lone pairs on outer atoms. 5. So that C has an octet, we shall form DOUBLE BONDS between C and O. The second bonding pair forms adoublebond.

  21. Formaldehyde, CH2O C H O 12 electronsavailable 4 + 2 × 1 + 6 = C Carbon is the least electronegative, hydrogen can form only one bond.

  22. Formaldehyde, CH2O C H O 12 electronsavailable 4 + (2x1) + 1x6 = H H C O Carbon is the least electronegative, hydrogen can form only one bond.

  23. Formaldehyde, CH2O C H O 12 electronsavailable 4 + 2 × 1 + 6 = H H C O Carbon is the least electronegative, hydrogen can form only one bond.

  24. Formaldehyde, CH2O C H O 12 electronsavailable H H C O Carbon is the least electronegative, hydrogen can form only one bond.

  25. Formaldehyde, CH2O C H O 12 electronsavailable H H C O : : Carbon is the least electronegative, hydrogen can form only one bond.

  26. Formaldehyde, CH2O H O H : : C H H C Additional Stuctures: O : :

  27. Formaldehyde, CH2O H O H : : C : : H O H C H H C Additional Stuctures: O : : : :

  28. Formaldehyde, CH2O H O H : : C : : H O H C H H C Additional Stuctures: correct structure O : : Not this one… : : or this one…

  29. Formaldehyde, CH2O H O H : : C : : H O H C H H C Additional Stuctures: correct structure O : : nature likes symmetry! Not this one… : : or this one…

  30. Formaldehyde, CH2O H O H : : C : : H O H C H H C Additional Stuctures: correct structure O : : nature likes symmetry! Not this one… or this one… You will see why the first structure is favored by formal chargers later.

  31. Double and even triple bonds are commonly observed for C, N, P, O, and S Double & Triple Bonds H2CO SO3 C2F4

  32. Lewis Structures of Common Oxoacids & Their Anions

  33. Common Isoelectric Molecules & Ions Molecules and ions having the same number of valence electrons and the same Lewis structures are said to be isoelectronic

  34. Atom Formal Charges in Covalent Molecules & Ions • The formal charge is the charge that would reside on an atom in a molecule or polyatomic ion if we assume that all bonding electrons are shared equally. • The formal charge for an atom in a molecule or ion is calculated based on the Lewis structure of the molecule or ion, using: • NVE = the number of valence electrons in the uncombined atom (and equal to its group number in the periodic table). • LPE = number of lone pair electrons on an atom. • BE = number of bonding electrons around an atom. FC = NVE  [LPE + ½ BE]

  35. Calculating the Formal Charge on Each Atom in a Covalent Molecule H F H F Hydrogen Fluoride: FC = NVE  [LPE + ½ BE] H F Number of valence electrons: 1 7 Number of bonding electrons: 2 2 Number of lone pair electrons: 0 6 0 0 Formal charge: 0 0

  36. Calculating the Formal Charge on Each Atom in a Covalent Molecule H F When the sum of the formal charges on the atoms in a molecule equals the expected overall charge on the molecule, the Lewis structure is valid. HF is expected to be a neutral compound, and the formal charges validate this Lewis dot structure. + 0 = 0 0

  37. Carbon Dioxide, CO2 • • • • O C O • • • • O-atom: FC = 6 – [4 + ½  4] = 0 0 + 0 + 0 = 0 C-atom: FC = 4 – [0 + ½  8] = 0

  38. What if this was the structure of CO2? O-atom: FC = 6 – [6 + ½ 2] = -1 C-atom: FC = 4 – [0 + ½ 4] = +2 It is possible but structure, but it might not be the most stable.

  39. Use Formal Charges to Predict Which Atom Carries the Negative Charge on CN− [ :C  N: ]– –1 0 FC(C) FC(N) The negative charge resides on the carbon atom. As expected the overall charge is 1

  40. CH2O Formal Charges H O H : : C H C H : : H O H C O 0 0 0 Generally, the structure with the lowest formal charges on each atom or negative charges on the atoms with the highest EA’s are often most favored. This one is most favorable! : : 0 0 +2 0 0 0 not favorable! -1 +1 -2 not favorable!

  41. Resonance : : : : : : : : : : O N O O N O O N O : : : : : : : : : : : : O O O – – – Consider the nitrate anion: The double bond does not have to be on the vertical O-atom. There are three equivalent structures that can be drawn. We see that the double bond moves around or “resonates” between the structures indicating that the molecule exhibits what is called “Resonance” .

  42. Resonance : : : : : : : : : : : O N O O N O O N O : : : : : : : : : : : : O O O – – – All three of the Lewis structures are equivalent Resonance Structures. Resonance stabilizes the energy (lowers) by distributing electron density over the entire molecule. The measured fHo is more exothermic than the predicted value.

  43. Resonance : : : O N O : : : : O – Identifying Resonance: Compounds that exhibit resonance will have fractional bond orders.

  44. Exceptions to the Octet Rule BF3 SF4 Certain elements can violate the octet rule. Boron may form stable compounds with only 6 valence electrons. Compounds in the 3rd period and beyond can have more than 8 electrons.

  45. Exceptions to the Octet Rule • There are three classes of exceptions to the octet rule: • Molecules with an odd number of electrons. • 2. Molecules in which one atom has less than an octet • 3. Molecules in which one atom has more than an octet.

  46. Odd Numbers of Electrons Examples:Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons. NO (nitrogen monoxide) Available electrons= 11 There must be an odd electron! Molecular Orbital Theory will explain why NO is stable later. or

  47. Incomplete Octet H H B H • Relatively rare. • Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A Most typical example is BH3. Available Electrons= 3 + 3(1) = 6 There are only 6 electrons on the central atom. Hydrogen may only form one bond. The formal charges (all zero) support the drawn structure.

  48. Central Atoms that Exceed the Octet Rule • Atoms from the 3rd period onwards can accommodate more than an octet. • Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density. BrF3 Example: Available Electrons = 7 + 3(7) = 28 Each fluorine atom can only share one electron. What does this mean??

  49. BrF3 Bromine has 7 valence electrons (4s24p5) Fluorine also has 7 (2s22p5) Br has the lowest electronegativity so it will be the central atomin the molecule. : F : : : Each F shares one electron to complete an octet. : Br F : : : : F There are 10 total electrons in the Br valence : : :

  50. SF6 F F F S F F F Available Electrons= 48 Sulfur has 6 valence electrons, fluorine has seven. If each F shares one electron with one of the six from S, F completes its octet. This will yield a compound with S in the center having 6 F’s bonded. 12 valence electrons around the S

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