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CHAPTER 8

CHAPTER 8. Bonding and Molecular Structure. Introduction. Bonds: Attractive forces that hold atoms together in compounds Valence Electrons: the outermost electrons -These e - are involved in bonding. Valence Electrons. Electrons are divided between core and valence electrons

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CHAPTER 8

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  1. CHAPTER 8 Bonding and Molecular Structure

  2. Introduction • Bonds: Attractive forces that hold atoms together in compounds • Valence Electrons: the outermost electrons • -These e- are involved in bonding

  3. 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

  4. Valence Electrons -The number of valence electrons of a main group atom is the Group number -For Groups IA-IVA, number of bonding (unpaired) electrons is equal to the group number -For Groups VA -VIIA, number of bonding (unpaired) electrons is equal to 8 - group number

  5. Valence Electrons -Except for H (and sometimes atoms of the 3rd group and higher) -The total number of valence electrons around a given atom in a molecule will be eight: OCTET RULE - (with the exception of hydrogen) atoms in molecules prefer to be surrounded by 8 electrons (or have 4 bonds = 8 electrons)

  6. Lewis Dot Formulas of Atoms IA IIA IIIA IVA VA VIA VIIA VIIIA H He Li Be B C N O F Ne

  7. Ionic Bonding An ion is an atom or a group of atoms possessing a net electrical charge • -cations: positive (+) ions • These atoms have lost 1 or more electrons • -anions: negative (-) ions • These atoms have gained 1 or more electrons

  8. Formation of Ionic Compounds • Monatomic ions consist of one atom • Examples: • Na+, Ca2+, Al3+ - cations • Cl-, O2-, N3- -anions • Polyatomic ions contain more than one atom Examples: • NH4+ - cation • NO2-,CO32-, SO42- - anions

  9. Formation of Ionic Compounds • General trend: • metals become isoelectronic with the preceding noble gas electron configuration • nonmetals become isoelectronicwith the following noble gas electron configuration

  10. Formation of Ionic Compounds • Reaction of Group IA Metals with Group VIIA Nonmetals

  11. Formation of Ionic Compounds • 1s2s2p Li  . . F    These atoms form ions with these configurations. Li+. same configuration as [He] F- same configuration as [Ne]

  12. Formation of Ionic Compounds • In general: the reaction of IA metals and VIIA nonmetals: 2 M(s) + X2 2 MX(s) • where M is the metals Li to Cs • and X is the nonmetals F to I Electronically it looks like: nsnp nsnp M   M+ __ __ __ __ X   X- 

  13. Formation of Ionic Compounds reaction of IIA metals with VIIA nonmetals: Be(s) + F2(g)BeF2(g)

  14. Formation of Ionic Compounds The valence electrons in these two elements react like: 2s2p2s2p Be [He]  __ __ __  Be2+ __ __ __ __ F [He]    F-  Lewis dot structure representation:

  15. Formation of Ionic Compounds The remainder of the IIA metals and VIIA nonmetals react similarly: M(s) + X2 MX2 M can be any of the metals Be to Ba X can be any of the nonmetals F to I

  16. Formation of Ionic Compounds For the reaction of IA metals with VIA nonmetals: Draw the valence electronic configurations for Li, O, and their appropriate ions

  17. Formation of Ionic Compounds • Draw the electronic configurations for Li, O, and their appropriate ions You do it! 2s2p2s2p Li [He]   Li1+ O [He]   O2-  Draw the Lewis dot formula representation of this reaction

  18. Formation of Ionic Compounds Simple Binary Ionic Compounds Table • Reacting GroupsGeneral FormulaExample IA + VIIA MX NaF IIA + VIIA MX2 BaCl2 IIIA + VIIA MX3 AlF3 IA + VIA M2X Na2O IIA + VIA MX BaO IIIA + VIA M2X3 Al2S3

  19. Formation of Ionic Compounds • Reacting Groups General FormulaExample IA + VA M3X Na3N IIA + VA M3X2 Mg3P2 IIIA + VA MX AlN -H forms ionic compounds when bound to metals (IA and IIA metals For example: LiH, KH, CaH2, and BaH2 -When H is bound to nonmetals, the compounds are covalent in nature

  20. Formation of Covalent Bonds • -potential energy of an H2 molecule as a function of the distance between the two H atoms

  21. Covalent Bonding • Atoms share electrons • If the atoms share: • 2 electrons a single covalent bond is formed • 4 electrons - a double covalent bond • 6 electrons - a triple covalentbond Atoms have a lower potential energy when bound…this is a more favorable situation (why?)

  22. Writing Lewis Formulas: • 1. Add the number of valence electrons for all the atoms that are present in the molecule • 2. Add or subtract electrons based on the molecule’s (or ion’s) charge • 3. Identify the central atom and draw a skeletal structure: • -the one that requires the most e- to complete octet • -the less electronegative • 4. Place a bond between each atom (1 bond = 2 e-) • 5. Fill in octet of outer atoms first • 6. Finish by completing the octet of central atom • – if you run out of e- then multiple bonds must be created between the central atom and atoms bound to it

  23. Writing Lewis Formulas octet rule: representative elements usually attain stable noble gas electron configurations (8 valence e-) in most compounds You must distinguish the difference between: • -bonding electrons and nonbonding electrons -shared (paired) and unshared (unpaired) electrons

  24. Formation of Covalent Bonds • Lewis dot structures: • 1. H2 molecule formation: 2. HCl molecule formation:

  25. Lewis Structures • Homonuclear diatomic molecules • 1. Two atoms of the same element, H2: 2. Fluorine, F2: • Nitrogen, N2:

  26. Lewis Structures heteronuclear diatomic molecules 1. hydrogen fluoride, HF 2. hydrogen chloride, HCl 3. hydrogen bromide, HBr

  27. Lewis Structures • Water, H2O • Ammonia molecule , NH3

  28. Lewis Structures • Polyatomic ions: • ammonium ion NH4+ Notice that the N-atom in this molecule has eight electrons around them (H does not)

  29. Writing Lewis Formulas • Sulfite ion, SO32-.

  30. H2CO Double and even triple bonds are commonly observed for C, N, P, O, and S SO3 C2F4

  31. Lewis Structures • Example: Write Lewis dot and dash formulas for sulfur trioxide, SO3

  32. Resonance • There are three possible structures for SO3: • -Two or more Lewis formulas are necessary to show the bonding in a molecule • -use equivalent resonance structures to show the molecule’s structure • -Double-headed arrows are used to indicate resonance formulas

  33. Resonance Resonance is a flawed method of representing molecules • -There are no single or double bonds in SO3

  34. Sulfur Dioxide, SO2 1. Central atom = 2. Valence electrons = ___ or ___ pairs 3. Write the Lewis structure 4. Form double bond so that S has an octet — but note that there are two ways of doing this.

  35. Limitations of the Octet Rule • There are some molecules that violate the octet rule: • - Be • - Group IIIA • -Odd number of total electrons. • -Central element must have a share of more than 8 valence electrons to accommodate all of the substituents. (i.e. S and P)

  36. Limitations of the Octet Rule • Example: Write Lewis formula for BBr3.

  37. Sulfur Tetrafluoride, SF4 Central atom = Valence electrons = ___ or ___ pairs. Form sigma bonds and distribute electron pairs. 5 pairs around the S atom. A common occurrence outside the 2nd period.

  38. Limitations of the Octet Rule • Example: Write dot structures for AsF5.

  39. Formal Atomic Charges • Atoms in molecules often bear a charge (+ or -) • The predominant resonance structure of a molecule is the one with charges on atoms as close to 0 as possible • Formal charge = Group number – 1/2 (# of bonding electrons) - (# of Lone electrons) • = Group number – (# of bonds) • – (# of Lone electrons)

  40. Formal Charge CO2 . . . . . . . .

  41. • • • • • • S C N S C N • • • • • • • • • • • • S C N • • • • • • Formal Charge Thiocyanate Ion, SCN- Which is the most stable resonance form?

  42. Theories of Covalent Bonding • Valence Shell Electron Pair Repulsion Theory • Commonly designated as VSEPR • Principal originator • R. J. Gillespie in the 1950’s • Valence Bond Theory (Chapter 9) • Involves the use of hybridized atomic orbitals • Principal originator • L. Pauling in the 1930’s & 40’s

  43. VSEPR Theory electron densities around the central atom are arranged as far apart as possible to minimize repulsions (why?) • Five basic molecular shapes: • Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral

  44. VSEPR Theory • Two regions of high electron density around the central atom.

  45. VSEPR Theory 2. Three regions of high electron density around the central atom.

  46. VSEPR Theory 3. Four regions of high electron density around the central atom.

  47. VSEPR Theory 4. Five regions of high electron density around the central atom.

  48. VSEPR Theory 5. Six regions of high electron density around the central atom.

  49. VSEPR Theory • Electronic geometry(family): locations of regions of electron density around the central atom(s) • Molecular geometry:arrangement of atoms around the central atom(s) Electron pairs are not used in the molecular geometry determination

  50. VSEPR Theory Lone pairs (unshared pairs) of electrons require more volume than shared pairs • -there is an ordering of repulsions of lone electrons around central atom Criteria for the ordering of the repulsions: 1. Lone pair to lone pair is the strongest repulsion. 2. Lone pair to bonding pair is intermediate repulsion. 3. Bonding pair to bonding pair is weakest repulsion.

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