Structure and Molecular Bonding CAPTER 9 1 Introduction Bonds - Attractive forces that hold atoms together in compounds Valence Electrons - The electrons involved in bonding are in the outermost (valence) shell. 2 Valence Electrons core and Electrons are divided between core and valence electrons B 1s 2 2s 2 2p 1 Core = [e], valence = 2s 2 2p 1 Br [Ar[ Ar] ] 3d 10 4s 2 4p 5 Core = [Ar[ Ar] ] 3d 10, valence = 4s 2 4p 5 3
Rules of the Game No. of valence electrons of a main group atom = Group number For Groups 1A-4A, no. of bond pairs = group number. For Groups 5A -7A, BP s s = 8 - Grp.. No. 4 Rules of the Game 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 s = 8 - Grp.. No. Except for (and sometimes atoms of 3rd and higher periods), BP s s + LP s s = 4 This observation is called the OCTET RULE 5 Lewis Dot Formulas of Atoms... e............... Li Be B.. C.. N.. O........ F. Ne.. 6
Ionic Bonding Formation of Ionic Compounds An ion is an atom or a group of atoms possessing a net electrical charge. 1. positive (+) ions or cations These atoms have lost 1 or more electrons. 2. negative (-) ions or anions These atoms have gained 1 or more electrons. 7 Formation of Ionic Compounds Monatomic ions consist of one atom. Examples: Na +, Ca 2+, Al 3+ - cations Cl -, O 2-, N 3- -anions Polyatomic ions contain more than one atom. N 4+ - cation NO 2-,CO 3 2-, SO 4 2- - anions 8 Formation of Ionic Compounds Reaction of Group IA Metals with Group VIIA Nonmetals G -1metal G -17 nometal 2 Li silver solid (s) + F yellow gas 2(g) 2 LiF (s) white solid with an 842 o melting point C 9
Formation of Ionic Compounds 1s 2s 2p Li F These atoms form ions with these configurations. Li + F - same configuration as [Ne] Li. +..... F. Li +... F... same configuration as [e] [ ] 10 Formation of Ionic Compounds General trend Cations become isoelectronic with the preceding noble gas. Anions become isoelectronic with the following noble gas. 11 Formation of Ionic Compounds In general for the reaction of 1 metals and 17 nonmetals, the reaction equation is: 2 M (s) + X 2 2 M + X - (s) where M is the metals Li to Cs and X is the nonmetals F to I. Electronically this is occurring. ns np ns np M M + X X - 12
Formation of Ionic Compounds Next we examine the reaction of 2 metals with 17 nonmetals. One example is the reaction of Be and F 2. Be (s) + F 2(g) BeF 2(g) 13 Formation of Ionic Compounds The valence electrons in these two elements are reacting in this fashion. 2s 2p 2s 2p Be [e] Be 2+ F [e] F - 14 Formation of Ionic Compounds... F... Be... Be 2+ 2. Ḟ..... F... The remainder of the 2 metals and 17nonmetals react similarly. Symbolically this can be represented as: M (s) + X 2 M 2+ X 2 - M can be any of the metals Be to Ba. X can be any of the nonmetals F to Cl. 15
Formation of Ionic Compounds For the reaction of 1 metals with 16 nonmetals, a good example is the reaction of lithium with oxygen. The reaction equation is: 4 Li (s) + O 2(g) 2 Li + 2 O 2 ( - s) 16 Formation of Ionic Compounds Draw the electronic configurations for Li, O, and their appropriate ions. You do it! 2s 2p 2s 2p Li [e] Li 1+ O [e] O 2- Draw the Lewis dot formula representation of this reaction. 17 Formation of Ionic Compounds Simple Binary Ionic Compounds Table Reacting Groups Compound General Formula Example 1 + 17 MX NaF 2 + 17 MX 2 BaCl 2 3 + 17 MX 3 AlF 3 1 + 16 M 2 X Na 2 O 2 + 16 MX BaO 3 + 16 M 2 X 3 Al 2 S 3 18
Formation of Ionic Compounds Reacting Groups Compound General Formula Example 1 + 15 M 3 X Na 3 N 2 + 15 M 3 X 2 Mg 3 P 2 3 + 15 MX AlN, a nonmetal, forms ionic compounds with 1 and 2 metals for example, Li, K, Ca 2, and Ba 2. Other hydrogen compounds are covalent. 19 Formation of Ionic Compounds Coulomb s Law inverse square law + ( q )( q ) F 2 d where F = force of attraction between ions q = magnitude of charge on ions d = distance between center of ions 20 Formation of Ionic Compounds Force - Small ions with high ionic charges >> large ions with small ionic charges Al 3+ 2-2+ 2-1+ - 2 O3 > Ca O > K Cl 21
Covalent Bonding Atoms share electrons. If the atoms share 2 electrons a single covalent bond is formed. 4 electrons - a double bond. 6 electrons - a triple bond. The atoms have a lower potential energy when bound. 22 Formation of Covalent Bonds This figure shows the potential energy of an 2 molecule as a function of the distance between the two atoms. 23 Writing Lewis Formulas: 1. Sum the number of valence electrons for atoms present. 2. Add or subtract electrons for the charge. 3. Identify the central atom (one that requires more e- to complete octet less e-neg if in same group) and draw a skeletal structure. 4. Place a bond between each atom (2 e- per) 5. Fill in octet of outer atoms. 6. Complete octet of central atom if deficient make multiple bonds. 24
Formation of Covalent Bonds Use Lewis dot formulas 1. molecule formation representation.. +.. or 2 2. Cl molecule formation..... + Cl..... Cl or Cl.... 25 Writing Lewis Formulas: The Octet Rule Lewis octet rule - representative elements usually attain stable noble gas electron configurations in most of their compounds. Distinguish between bonding (or shared) electrons and nonbonding (or unshared or lone pairs) of electrons. 26 Lewis Formulas for Molecules and Polyatomic Ions omonuclear diatomic molecules. 1. Two atoms of the same element. 1. ydrogen molecule, 2.. or 2. Fluorine, F 2...... F..... F. or... F F....... 3. Nitrogen, N 2. N N or N N 27
Lewis Formulas for Molecules and Polyatomic Ions Next, look at heteronuclear diatomic molecules. 1. hydrogen fluoride, F. F or F 2. hydrogen chloride, Cl. Cl or Cl 3. hydrogen bromide, Br. Br or Br 28 Lewis Formulas for Molecules and Polyatomic Ions Water, 2 O Ammonia molecule, N 3 29 Lewis Formulas for Molecules and Polyatomic Ions Water, 2 O O Ammonia molecule, N 3 N 30
Lewis Formulas for Molecules and Polyatomic Ions Polyatomic ions. One example is the ammonium ion, N 4+. + N Notice that the atoms other than in these molecules have eight electrons around them. 31 Writing Lewis Formulas: The Octet Rule Example: Write Lewis dot and dash formulas for the sulfite ion, SO 2-3. 32 Writing Lewis Formulas: The Octet Rule Sulfite ion, SO 3 2-. O S O O 2- or O S O O 2-33
Double and even triple bonds are commonly observed for C, N, P, O, and S 2 CO SO 3 C 2 F 4 O C O 34 Lewis Structures Example: Write Lewis dot and dash formulas for sulfur trioxide, SO 3. 35 Lewis Structures Example: Write Lewis dot and dash formulas for sulfur trioxide, SO 3. O S O or O S O O O 36
Resonance There are three possible structures for SO 3.. O S O O O S O O O S O O otwo or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent resonance structures to show the molecule s structure. odouble-headed arrows are used to indicate resonance formulas. 37 Resonance Resonance is a flawed method of representing molecules. There are no single or double bonds in SO 3. O S O O 38 Sulfur Dioxide, SO 2 1. Central atom = S 2. Valence electrons = 18 or 9 pairs O S O 3. Form double bond so that S has an octet but note that there are two ways of doing this. bring in left pair O S O OR bring in right pair 39
Sulfur Dioxide, SO 2 bring in left pair O This leads to the following resonance structures. S OR bring in right pair O O S O O S O 40 Writing Lewis Formulas: Limitations of the Octet Rule There are some molecules that violate the octet rule. For these molecules the N - A = S rule does not apply: 1. - Be. 2. - Group - 13. 3. -Odd number of total electrons. 4. -Central element must have a share of more than 8 valence electrons to accommodate all of the substituents. (I.e. some S, P) 41 Writing Lewis Formulas: Limitations of the Octet Rule Example: Write Lewis formula for BBr 3. Br. B Br. B Br or Br Br B Br Br 42
Sulfur Tetrafluoride,, SF 4 Central atom = Valence electrons = or pairs. Form sigma bonds and distribute electron pairs. F S F F F 5 pairs around the S atom. A common occurrence outside the 2nd period. 43 Writing Lewis Formulas: Limitations of the Octet Rule Example: Write dot and dash formulas for AsF 5. 44 Formal Atom Charges Atoms in molecules often bear a charge (+ or -). The predominant resonance structure of a molecule is the one with charges as close to 0 as possible. Formal charge = Group number 1/2 (no. of bonding electrons) - (no. of LP electrons) 45
Calculated Partial Charges in CO 2 Yellow = negative & red = positive Relative size = relative charge 46 Thiocyanate Ion, SCN - 6 - (1/2)(2) - 6 = -1 5 - (1/2)(6) - 2 = 0 S C N 4 - (1/2)(8) - 0 = 0 47 Thiocyanate Ion, SCN - S C N S C N S C N Which is the most stable resonance form? 48
Calculated Partial Charges in SCN - All atoms negative, but most on the S S C N 49 Dipole Moments Asymmetric charge distribution The dipole moment has the symbol µ. µ is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q. 50 Dipole Moments For example, F and I: a + - δ - Fδ 1.91 Debye units a + - δ - I δ 0.38 Debye units 51
Dipole Moments There are some nonpolar molecules that have polar bonds. There are two conditions that must be true for a molecule to be polar. 1. There must be at least one polar bond present or one lone pair of electrons. 2. The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another. 3. Examples (water, CF 4, CO 2, N 3, N 4+ ) 52 Polar Molecules: The Influence of Molecular Geometry Molecular geometry affects molecular polarity. Due to the effect of the bond dipoles and how they either cancel or reinforce each other. A B A linear molecule nonpolar A B A angular molecule polar 53 Polar Molecules: The Influence of Molecular Geometry Polar Molecules must meet two requirements: 1. One polar bond or one lone pair of electrons on central atom. 2. Neither bonds nor lone pairs can be symmetrically arranged that their polarities cancel. 54
Polar and Nonpolar Covalent Bonds Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds. Nonpolar covalent bonds have a symmetrical charge distribution. N N or N N. or 55 Polar and Nonpolar Covalent Bonds Polar covalent bonds - electrons are not shared equally -different electronegativities. Electronegativities 2.1 14243 4.0 Difference = 1.9 1.9 F very polar bond 56 Polar and Nonpolar Covalent Bonds Compare F to I. I Electronegativities 2.1 14243 2.5 Difference = 0.4 slightly polar bond 0.4 57
Two Simple 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 10) Involves the use of hybridized atomic orbitals Principal originator L. Pauling in the 1930 s & 40 s 58 VSEPR Theory VSEPR - electron density around the central atom are arranged as far apart as possible to minimize repulsions. Five basic molecular shapes 59 VSEPR Theory 1 Two regions of high electron density around the central atom. 60
VSEPR Theory 2 Three regions of high electron density around the central atom. 61 VSEPR Theory 3 Four regions of high electron density around the central atom. 62 VSEPR Theory 4 Five regions of high electron density around the central atom. 63
C VSEPR Theory 5 Six regions of high electron density around the central atom. 64 VSEPR Theory 1. Electronic geometry - locations of regions of electron density around the central atom(s). 2. Molecular geometry - arrangement of atoms around the central atom(s). Electron pairs are not used in the molecular geometry determination just the positions of the atoms in the molecule are used. 65 VSEPR Theory An example of a molecule that has the same electronic and molecular geometries is methane - C 4. Electronic and molecular geometries are tetrahedral. 66
C VSEPR Theory An example of a molecule that has different electronic and molecular geometries is water - 2 O. Electronic geometry is tetrahedral. Molecular geometry is bent or angular. 67 VSEPR Theory Lone pairs of electrons (unshared pairs) require more volume than shared pairs. Consequently, there is an ordering of repulsions of electrons around central atom. Criteria for the ordering of the repulsions: 68 VSEPR Theory 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. Mnemonic for repulsion strengths lp/lp > lp/bp > bp/bp Lone pair to lone pair repulsion is why bond angles in water are less than 109.5 o. 69
Molecular Shapes and Bonding In the next sections we will use the following terminology: A = central atom B = bonding pairs around central atom U = lone pairs around central atom For example: AB 3 U designates that there are 3 bonding pairs and 1 lone pair around the central atom. 70 Linear Electronic Geometry:AB 2 Species (No Lone Pairs of Electrons on A) Some examples of molecules with this geometry are: BeCl 2, BeBr 2, BeI 2, gcl 2, CdCl 2 All of these examples are linear, nonpolar molecules. Important exceptions occur when the two substituents are not the same! BeClBr or BeIBr will be linear and polar! 71 Trigonal Planar Electronic Geometry: AB 3 Species (No Lone Pairs of Electrons on A) Some examples of molecules with this geometry are: BF 3, BCl 3 All of these examples are trigonal planar, nonpolar molecules. Important exceptions occur when the three substituents are not the same! BF 2 Cl or BCI 2 Br will be trigonal planar and polar! 72
Tetrahedral Electronic Geometry: AB 4 Species (No Lone Pairs of Electrons on A) Some examples of molecules with this geometry are: C 4, CF 4, CCl 4, Si 4, SiF 4 All of these examples are tetrahedral, nonpolar molecules. Important exceptions occur when the four substituents are not the same! CF 3 Cl or C 2 CI 2 will be tetrahedral and polar! 73 Tetrahedral Electronic Geometry: AB 4 Species (No Lone Pairs of Electrons on A) 74 Tetrahedral Electronic Geometry: AB 3 U Species (One Lone Pair of Electrons on A) Some examples of molecules with this geometry are: N 3, NF 3, P 3, PCl 3, As 3 These molecules are our first examples of central atoms with lone pairs of electrons. Thus, the electronic and molecular geometries are different. All three substituents are the same but molecule is polar. N 3 and NF 3 are trigonal pyramidal, polar molecules. 75
Tetrahedral Electronic Geometry: AB 2 U 2 Species (Two Lone Pairs of Electrons on A) Some examples of molecules with this geometry are: 2 O, OF 2, 2 S These molecules are our first examples of central atoms with two lone pairs of electrons. Thus, the electronic and molecular geometries are different. Both substituents are the same but molecule is polar. Molecules are angular, bent, or V-shaped and polar. 76 Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U2, and AB 2 U 3 Some examples of molecules with this geometry are: PF 5, AsF 5, PCl 5, etc. These molecules are examples of central atoms with five bonding pairs of electrons. The electronic and molecular geometries are the same. Molecules are trigonal bipyramidal and nonpolar when all five substituents are the same. If the five substituents are not the same polar molecules can result, AsF 4 Cl is an example. 77 Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U2, and AB 2 U 3 If lone pairs are incorporated into the trigonal bipyramidal structure, there are three possible new shapes. 1. One lone pair - Seesaw shape 2. Two lone pairs - T-shape 3. Three lone pairs linear The lone pairs occupy equatorial positions because they are 120 o from two bonding pairs and 90 o from the other two bonding pairs. Results in decreased repulsions compared to lone pair in axial position. 78
C Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U2, and AB 2 U 3 AB 4 U molecules have: 1. trigonal bipyramid electronic geometry 2. seesaw shaped molecular geometry 3. and are polar One example of an AB 4 U molecule is SF 4 79 Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U2, and AB 2 U 3 Molecular Geometry 80 Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 AB 3 U 2 molecules have: 1. trigonal bipyramid electronic geometry 2. T-shaped molecular geometry 3. and are polar One example of an AB 3 U 2 molecule is IF 3 81
C C Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U2, and AB 2 U 3 Molecular Geometry 82 Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U2, and AB 2 U 3 AB 2 U 3 molecules have: 1.trigonal bipyramid electronic geometry 2.linear molecular geometry 3.and are nonpolar One example of an AB 3 U 2 molecule is 83 Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U2, and AB 2 U 3 Molecular Geometry 84
Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 Some examples of molecules with this geometry are: SF 6, SeF 6, SCl 6, etc. These molecules are examples of central atoms with six bonding pairs of electrons. Molecules are octahedral and nonpolar when all six substituents are the same. If the six substituents are not the same polar molecules can result, SF 5 Cl is an example. 85 Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 If lone pairs are incorporated into the octahedral structure, there are two possible new shapes. 1. One lone pair - square pyramidal 2. Two lone pairs - square planar The lone pairs occupy axial positions because they are 90 o from four bonding pairs. Results in decreased repulsions compared to lone pairs in equatorial positions. 86 Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 AB 5 U molecules have: 1.octahedral electronic geometry 2.Square pyramidal molecular geometry 3.and are polar. One example of an AB 4 U molecule is IF 5 87
C C Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 Molecular Geometry 88 Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 AB 4 U 2 molecules have: 1.octahedral electronic geometry 2.square planar molecular geometry 3.and are nonpolar. One example of an AB 4 U 2 molecule is XeF 4 89 Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 Molecular Geometry Polarity 90
Compounds Containing Double Bonds Ethene or ethylene, C 2 4, is the simplest organic compound containing a double bond. Compound must have a double bond to obey octet rule. 91 Compounds Containing Double Bonds Lewis Dot Formula C C o r C C 92 Bond Properties What is the effect of bonding and structure on molecular properties? Free rotation around C C C C single bond No rotation around C=C double bond 93
Bond Order # of bonds between a pair of atoms Double bond Single bond Acrylonitrile Triple bond 94 Bond Order Fractional bond orders occur in molecules with resonance structures. Consider NO - 2 N N O O O O The N O N O bond order = 1.5 Bond order = Total # of e - pairs used for a type of bond Total # of bonds of that type Bond order = 3 e - pairs in N O bonds 2 N O bonds 95 Bond Order Bond order is proportional to two important bond properties: (a) bond strength (b) bond length 414 kj 123 pm 110 pm 745 kj 96
Bond Length Bond length is the distance between the nuclei of two bonded atoms. 97 F Bond Length Bond length depends on size of bonded atoms. Cl Bond distances measured in Angstrom units where 1 A = 10-2 pm. I 98 Bond Length Bond length depends on bond order. Bond distances measured in in Angstrom units where 1 A = 10 10-2 -2 pm. 99
Bond Strength measured by the energy req d to break a bond. See Table 9.10. BOND STRENGT (kj/mol) 436 C C 346 C=C 602 C C 835 N N 945 The GREATER the number of bonds (bond order) the IGER the bond strength and the SORTER the bond. 100 101 Bond Strength Measured as the energy req d to break a bond. See Table 9.10 102
Bond Strength Measured as the energy req d to break a bond. See Table 9.10. BOND STRENGT (kj/mol) 436 C C 346 C=C 602 C C 835 N N 945 The GREATER the number of bonds (bond order) the IGER the bond strength and the SORTER the bond. 103 Molecular Polarity Water Boiling point = 100 C Methane Boiling point = -161 C Why do water and methane differ so much in their boiling points? Why do ionic compounds dissolve in water? 104 Bond Polarity Three molecules with polar, covalent bonds. Each bond has one atom with a slight negative charge (-δ)( and and another with a slight positive charge (+ δ) 105
Electronegativity, χ χ is a measure of the ability of an atom in a molecule to attract electrons to itself. Concept proposed by Linus Pauling 1901-1994 1994 106 Electronegativity Figure 9.14 107 Molecular Polarity Molecules will be polar if a) bonds are polar AND b) the molecule is NOT symmetric All above are NOT polar 108
Polar or Nonpolar? Compare CO 2 and 2 O. Which one is polar? 109 Polar or Nonpolar? Consider AB 3 molecules: BF 3, Cl 2 CO, and N 3. 110 F Molecular Polarity, BF 3 F B F B atom is positive and F atoms are negative. B F F bonds in BF 3 are polar. But molecule is symmetrical and NOT polar 111
F Molecular Polarity, BF 2 B F B atom is positive but & F atoms are negative. B F F and B B bonds in BF 2 are polar. But molecule is NOT symmetrical and is polar. 112 Is Methane, C 4, Polar? Methane is symmetrical and is NOT polar. 113 Is C 3 F Polar? C F F bond is very polar. Molecule is not symmetrical and so is polar. 114
C 4 CCl 4 Polar or Not? Only C 4 and CCl 4 are NOT polar. These are the only two molecules that are symmetrical. 115