Chemical Bond A chemical bond is a strong

Chemical Bond A chemical bond is a strong attractive force that exists between certain atoms in a substance. There are three types of chemical bonds: Ionic bonds Covalent bonds Metallic bonds קשר כימי 8 - 1

An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions. קשר כימי 8 - 2

Ionic Bond An ionic bond forms when one or more electrons are transferred from the valence shell of one atom to the valence shell of another atom. The atom that transferred the electron(s) becomes a cation. The atom that gained the electron(s) becomes an anion. קשר כימי 8 - 3

A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the chemical symbol of the element. Note – Dots are placed one to each side, until all four sides are occupied. קשר כימי 8 - 4

Lewis electron-dot symbol Table 9. 1 illustrates the Lewis electron-dot symbols for second- and third-period atoms. קשר כימי 8 - 5

Lewis electron-dot symbol Use Lewis electron-dot symbols to represent the transfer of electrons from magnesium to fluorine atoms to form ions with noble-gas configurations. קשר כימי 8 - 6

Ionic Bond An ionic bond between two atoms is formed in two steps: 1. An electron is transferred between the two separate atoms to give ions. 2. The ions attract one another, forming an ionic bond. Both steps occur simultaneously. קשר כימי 8 - 7

Coulomb's Law The potential energy obtained in bringing two charges Q 1 and Q 2, initially apart, up to a distance r apart is directly proportional to the product of the charges and inversely proportional to the distance between them. k is a physical constant (8. 99 × 109 J·m/C 2). קשר כימי 8 - 8

The lattice energy is the change in energy that occurs when an ionic solid is separated into gasphase ions. Na. Cl(s) → Na+(g) + Cl-(g) It is very difficult to measure lattice energy directly. קשר כימי 8 - 9

Trends in Lattice Energy: Ion Size • The force of attraction between charged particles is inversely proportional to the distance between them. • Larger ions mean the center of positive charge (nucleus of the cation) is farther away from the negative charge (electrons of the anion). 10 קשר כימי 8 -

Lattice Energy versus Ion Size 11 קשר כימי 8 -

Trends in Lattice Energy: Ion Charge • The force of attraction between oppositely charged particles is directly proportional to the product of the charges. • Larger charge means the ions are more strongly attracted. – Larger charge = stronger attraction – Stronger attraction = larger lattice energy • Of the two factors, ion charge is generally more important. 12 קשר כימי 8 -

Properties of Ionic Compounds Ionic solids are high-melting substances. This is due to the strong interactions between the ions. Charge on the ions of Mg. O is double the charge on the ions of Na. Cl thus, the force will be four times stronger. The size of Na+ is larger than that of Mg 2+; the size of Cl- is larger than that of O 2 -. The distance between Mg 2+ and O 2 - is smaller than the distance between Na+ and Cl-, the force between Mg 2+ and O 2 - will be greater. Thus, Mg. O has a higher melting point than Na. Cl. 13 קשר כימי 8 -

Properties of Ionic Compounds Based on the higher charge and the smaller distance for Mg. O, its melting point Mg. O should be significantly higher than the melting point of Na. Cl. The actual melting point of Na. Cl is 801°C; the melting point of Mg. O is 2800°C. 14 קשר כימי 8 -

Properties of Ionic Compounds • Hard and brittle crystalline solids – All are solids at room temperature. • Melting points generally > 300 C • The liquid state conducts electricity. – The solid state does not conduct electricity. • Many are soluble in water. – The solution conducts electricity well. 15 קשר כימי 8 -

Properties of Ionic Compounds 16 קשר כימי 8 -

Ionic Bonding Model versus Reality Lewis theory implies that if the ions are displaced from their positions in the crystal lattice, repulsive forces should occur. This predicts the crystal will become unstable and break apart. Lewis theory predicts ionic solids will be brittle. Ionic solids are brittle. When struck, they shatter. 17 קשר כימי 8 -

Conductivity of Na. Cl In Na. Cl(s), the ions are stuck in position and not allowed to move to the charged rods. 18 In Na. Cl(aq), the ions are separated and allowed to move to the charged rods. קשר כימי 8 -

Covalent Bonding in Molecules Covalent Bond: A bond that results from the sharing of electrons between atoms 19 קשר כימי 8 -

Covalent Bond To consider how a covalent bond forms, we can monitor the energy of two isolated hydrogen atoms as they move closer together. The potential energy decreases—first gradually, and then more steeply—to a minimum. As the atoms continue to move closer, it increases dramatically. The distance between the atoms when energy is at a minimum is called the bond length. קשר כימי 8 - 20

Covalent Bonding קשר כימי 8 - 21

Covalent Bonding G. N. Lewis (1916): Some atoms share e- to form bonds. • Number of bonds = Number shared e- pairs. • Molecules with shared e- are covalent compounds with covalent bonds. 22 קשר כימי 8 -

Single Covalent Bonds Lewis structures: dot = 1 e-. Line = 1 pair of e- H H H−H Single bond: one shared pair of e-. Octet rule To form bonds, elements gain, lose, or share e- to achieve 8 valence e-. • Doesn’t apply if Z ≤ 5 (H, He, Li, Be and B) קשר כימי 8 - 23

Single Covalent Bonds Bonding pairs = e- pairs shared between atoms. Lone pairs = unshared e- pair. bonding lone pair H O H bonding lone pair קשר כימי 8 - 24

Single Covalent Bonds F F 2 F = 2(7) = 14 valence e. Share 2 e- to form octets O + 2 H = 6 + 2(1) = 8 val. e. Two O-H bonds H N H H H O H N + 3 H = 5 + 3(1) = 8 val. e 3 N-H bonds קשר כימי 8 - 25

A coordinate covalent bond is formed when both electrons of the bond are donated by one atom. The two electrons forming the bond with the hydrogen on the left were both donated by the nitrogen. Once shared, they are indistinguishable from the other N—H bonds. קשר כימי 8 - 26

Single Covalent Bonds Those examples reveal a general rule for main-group elements: Number of e- shared = 8 – (A group number) Group Number of Valence e- Number of e- shared Example 4 A 4 8– 4=4 C in CH 4 5 A 5 3 N in NH 3 6 A 6 2 O in H 2 O 7 A 7 1 F in HF קשר כימי 8 - 27

Guidelines for Writing Lewis Structures 1. Count the valence e- in the molecule. 2. Draw a skeleton structure. Join atoms with single lines (pairs of e-). 3. Add e- pairs to form octets (except H). Start with terminal atoms. Extra e- ? Place around the central atom. Too few e-? Convert lone pairs into multiple bonds. קשר כימי 8 - 28

Writing Lewis Structures Phosphorus trifluoride, PF 3 1. PF 3 = 5 + 3 (7) = 26 valence e. P (group 5 A) 3 x F (group 7 A) 2. Skeleton (X is central in XYn ). F P F F קשר כימי 8 - 29

Writing Lewis Structures 3. Build octets – start with terminal atoms. F P F 6 e- used in 3 bonds, 20 e- remain (10 pairs) F 6 + 20 = 26 e- קשר כימי 8 - 30

Writing Lewis Structures Phosphate ion, PO 431. PO 43 - = 5 + 4(6) + 3 = 32 P (grp 5 A) 4 x O (grp 6 A) charge (-3) 3 e- O 2. Skeleton; P is central O P O O קשר כימי 8 - 31

Writing Lewis Structures 3. Add e- pairs: 8 e- used in 4 bonds, 24 e- remain (12 pairs) O O P O 32 e- used. O Add brackets and overall charge to show this is an ion. O 3 - O P O O קשר כימי 8 - 32

Multiple Covalent Bonds Too few “dots” to complete all the octets? Convert lone pairs to shared pairs. Methanal (formaldehyde) H 2 CO 1. Valence e- = 2(1) + 4 + 6 = 12 2. Skeleton 3. 6 e- in bonds. Add the other 3 pairs to O (outer atom). • Each H shares 2 e • C only “has” 6 e-. H H C O קשר כימי 8 - 33

Multiple Covalent Bonds Convert lone pairs to bond pairs. H H C O Each H shares 2 e. C shares 8 O “has” 8 4 shared + 2 lone pairs קשר כימי 8 - 34

Multiple Covalent Bonds Carbon dioxide CO 2 1. 4 + 2(6) = 16 e 2. Skeleton O C O 3. 4 e- in bonds. Add 3 pairs to each O. 4. Convert lone pairs to bond pairs. O C O קשר כימי 8 - 35

Multiple Covalent Bonds A single bond is a covalent bond in which one pair of electrons is shared by two atoms. A double bond is a covalent bond in which two pairs of electrons are shared by two atoms. ethene C 2 H 4 H H C C H H (ethylene) קשר כימי 8 - 36

Multiple Covalent Bonds A triple bond is a covalent bond in which three pairs of electrons are shared by two atoms. ethyne (acetylene) H-C≡C-H C 2 H 2 Double bonds form primarily with C, N, O, and S atoms. Triple bonds form primarily with C and N atoms. קשר כימי 8 - 37

A polar covalent bond (or polar bond) is a covalent bond in which the bonding electrons spend more time near one atom than near the other atom. Electronegativity, X, is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. קשר כימי 8 - 38

Electronegativity increases from left to right and from bottom to top in the periodic table. F, O, N, and Cl have the highest electronegativity values. קשר כימי 8 - 39

Polar Covalent Bond The difference in electronegativity between the two atoms in a bond is a rough measure of bond polarity. A small difference results in a nonpolar bond. A large difference results in a polar bond. A very large difference can result in an ionic bond. קשר כימי 8 - 40

Polar Covalent Bond Using electronegativities, arrange the following bonds in order by increasing polarity: P—H, H—O, C—Cl. The difference for P—H is 0. 0 The difference for H—O is 1. 4 The difference for C—Cl is 0. 5 The order is P—H < C—Cl < H—O קשר כימי 8 - 41

Polar Covalent Bond During bond formation, electrons are piled toward the more electronegative atom. The direction of electrons can be predicted using the electronegativity scale. In H–Cl bond, electrons are pulled toward the Cl atom (X = 3. 0) rather than the H atom (X = 2. 1). The Cl atom acquires a partial negative charge. The H atom acquires a partial positive charge. Thus, the HCl molecule is a polar molecule. קשר כימי 8 - 42

Bond Properties: Electronegativity Differences, ΔEN, determine bond polarity: EN ΔEN 4. 0 2. 1 2. 6 2. 1 4. 0 0. 9 4. 0 F–F H–Br H–F Na+ F- 0 0. 5 1. 9 3. 1 קשר כימי 8 - 43

Formal Charge The charge a bonded atom would have if its bonding e - were shared equally. Used to study charge distribution in a molecule. Method 1. Find the number of e- assigned to each atom: e- “on” an atom = (lone pair e-) + ½ (bonding e-) 2. Formal charge of each atom = (# of valence e-) – (e- “on” the atom). Note: sum of formal charges = molecular charge קשר כימי 8 - 44

Formal Charge [O C N] Valence e. Lone pair e½ shared e. Formal Charge O 6 6 1 -1 – C 4 0 N 5 2 3 0 Check: (formal charges) = ion charge = -1 קשר כימי 8 - 45

Formal Charge If there is choice between Lewis structures: • Smaller formal charges are favored. • Negative formal charges should be on the most EN atoms • Like charges should not be on adjacent atoms Which N 2 O structure is preferred? O N N Formal charges: -1 +1 0 Preferred. ENO > ENN O N N 0 +1 -1 קשר כימי 8 - 46

Formal Charge Which Cl. O 2 - structure is preferred? Formal charges: O Cl O -1 -1 0 0 +1 -1 Preferred (Smaller charges) קשר כימי 8 - 47

Lewis Structures & Resonance Ozone has 2 equivalent structures: O O O Both: • obey the octet rule • have the same number and types of bonds • have the same formal charges Experiments show that the OO bonds are identical. קשר כימי 8 - 48

Resonance 49 קשר כימי 8 -

Lewis Structures & Resonance structures are used to show O 3 is a mixture of both: O O O A resonance hybrid best represents O 3. Each bond is ≈1½ bonds. Ozone does NOT “flip” back and forth. Valid resonance structures: • Bonding and e- pair positions are changed. • Atom positions must not change. קשר כימי 8 - 50

Lewis Structures & Resonance in CO 32 - Experiment: All three CO bonds = 129 pm Typical bond lengths C-O =143 pm; C=O = 122 pm. קשר כימי 8 - 51

Lewis Structures and Resonance Benzene, C 6 H 6, is a ring compound best described using resonance ideas: Experiments: All C-C bonds have identical lengths. קשר כימי 8 - 52

Resonance & Structure of Benzene is usually drawn without its hydrogen atoms shown, or with a ring representing six delocalized electrons spread evenly over the carbon atoms: or Solid ring shows resonance קשר כימי 8 - 53

Metallic Bonds • The low ionization energy of metals allows them to lose electrons easily. • The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared by all atoms/ions in the metal. – An organization of metal cation islands in a sea of electrons – Electrons delocalized throughout the metal structure • Bonding results from attraction of the cations for the delocalized electrons. 54 קשר כימי 8 -

Exceptions to the Octet Rule Be and B form e- deficient compounds: H Be H 2 + 2(1) = 4 valence e- F F B 3 + 3(7) = 24 valence e- F קשר כימי 8 - 55

Fewer than Eight Valence Electrons Often very reactive F F B F H + N H H F F H B N F H H קשר כימי 8 - 56

Odd Number of Valence Electrons Some stable molecules have an odd number of e-. Examples NO 5 + 6 = 11 valence e- NO 2 5 + 2(6) = 17 valence e- N O O N O Free radical = atom or molecule with unpaired e-. Very reactive. Most stable molecules have paired e-. קשר כימי 8 - 57

More Than Eight Valence Electrons “Expanded octets” are relatively common. Low-lying d orbitals can accept extra e(only 3 rd period and beyond). Examples • 5 bonds (5 e- pairs) around P in PF 5 • NF 5 does not exist • 4 bonds and 1 lone-pair (5 e- pairs) around S in SF 4 • OF 4 does not exist. קשר כימי 8 - 58

More Than Eight Valence Electrons Cl. F 3 7 + 3(7) = 28 val. e- Make octets on F 24 e- used, 4 remain [bonds (3 x 2); Lone pairs (3 x 6)] F F Cl F Add 2 lone pairs to Cl – the 3 rd period element קשר כימי 8 - 59

Bond Properties: Bond Length Atom size is important. Remember the periodic table trends: < increasing size < קשר כימי 8 - 60

Bond Length Average single bond lengths (pm) I Br Cl S P Si F O N C H H 161 142 127 132 138 145 92 94 98 110 74 C 210 191 176 181 187 194 141 143 147 154 N 203 184 169 174 180 187 134 136 140 O 199 180 165 170 176 173 130 132 F 197 178 163 168 174 181 128 Si 250 231 216 221 227 234 P 243 224 209 214 220 S 237 218 203 208 Cl 232 213 200 Br 247 228 I 266 קשר כימי 8 - 61

Bond Length More bonds = greater e- density between atoms = atoms pulled together more strongly Multiple bonds (pm): N-N 140 C-C 154 C-N 147 N=N 120 C=C 134 C=N 127 N≡N 110 C≡C 121 C≡N 115 קשר כימי 8 - 62

Bond Enthalpy Average bond enthalpies (k. J/mol) I Br Cl S P Si F O N C H H 299 366 431 347 322 323 566 467 391 416 436 C 213 285 327 272 264 301 486 336 285 356 193 ~200 335 272 201 160 205 ~340 173 190 146 490 582 158 N O 201 F 255 326 226 Si 234 310 391 P 184 264 319 213 255 242 S Cl 209 217 Br 180 193 I 151 Multiple bonds: 226 209 226 Shorter bond = stronger bond. N-N 160 N=N 418 C-C 356 C=C 598 C≡C 813 C-N 285 C=N 616 C≡N 866 N≡N 946 קשר כימי 8 - 63

Molecular Geometry • Diatomic molecules are the easiest to visualize in three dimensions • HCl • Cl 2 • Diatomic molecules are linear Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 64

Two Electron Pairs • Linear • Bent • Bond angles • Always 180° in

Valence Shell Electron Pair Repulsion Theory • Ideal geometry of a molecule is determined by the way the electron pairs orient themselves in space • Orientation of electron pairs arises from electron repulsions • Electron pairs spread out so as to minimize repulsion Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 66

VSEPR Electron-Pair Geometries • Fundamental geometry exists that corresponds to the total number of electron pairs around the central atom: 2, 3, 4, 5, and 6 linear trigonal planar tetrahedral trigonal bipyramidal Copyright © 2016 Cengage Learning. All Rights Reserved. octahedral קשר כימי 8 - 67

Molecular Geometries • Trigonal planar • Electron pairs form an equilateral triangle around the central atom • Bond angles are 120° • Tetrahedral • Bond angles are 109. 5° Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 68

Molecular Geometries • Trigonal bipyramid • Bond angles vary • In the trigonal plane, 120° • Between the plane and the apexes, 90° • Between the central atom and both apexes, 180° Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 69

Molecular Geometries • Octahedron • Is a square bipyramid • Bond angles vary • 90° in and out of plane • 180° between diametrically opposite atoms and the central atom Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 70

Molecular Geometry Summarized Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8

Molecular Geometry Summarized Copyright © 2016 Cengage Learning. All Rights Reserved.

Unshared Pairs and Geometry • Electron-pair geometry • Same as that when single bonds are involved • Bond angles are a smaller than the ideal values • Molecular geometry • Quite different when one or more unshared pairs are present • While describing, only the positions of the bonded atoms are referred Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 73

Two Ways of Showing the Geometry of the NH 3 Molecule Copyright © 2016

Geometries with Two, Three, or Four Electron Pairs Around a Central Atom Copyright ©

Nonbonding Pairs and Bond Angle • Nonbonding pairs are physically larger than bonding pairs. • Therefore, their repulsions are greater; this tends to compress bond angles. 76 קשר כימי 8 -

Multiple Bonds and Bond Angles • Double and triple bonds have larger electron domains than single bonds. • They exert a greater repulsive force than single bonds, making their bond angles greater. 77 קשר כימי 8 -

Example Predict the geometry of NH 4+, BF 3, and PCl 3 Copyright ©

Example • Strategy: • Start by writing Lewis structures for each species • Focus on the central atom, then decide what species type (AX 2, AX 3, . . . ) the molecule or ion is • A represents the central atom • X represents the terminal atoms • E represents the unshared electron pairs • Match the species type with the molecular geometry and ideal bond angles for the species Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 79

Example Solution: • Lewis structure for NH 4+ Species type A = N, X

Example Lewis structure for BF 3 Species type A = B, X = F(3),

Example Lewis structure for PCl 3 Species type A = P, X = Cl(3), no E = 1→ AX 3 E Geometry trigonal pyramid (The ideal bond angles are 109. 5° but actually are 104°) Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 82

The A-X-E Notation • • A denotes a central atom X denotes a terminal atom E denotes a lone pair Example • Water • H 2 O • O is central • Two lone pairs • Two hydrogens • AX 2 E 2 Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 83

Lone Pairs and Expanded Octets • Where expanded octets are possible, place the extra lone pairs on the central atom • Example: Xe. F 4 Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 84

Multiple Bonds • For molecular geometry purposes, multiple bonds behave the same as single bonds • Electron pairs are located in the same place (between the nuclei) • Geometry of the molecule is determined by the number of terminal atoms • Not affected by the presence of a double or triple bond Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 85

Molecular Geometries for Molecules with Expanded Octets Copyright © 2016 Cengage Learning. All Rights

Example – – Predict the geometries of the Cl. O 3 ion, the NO 3 ion, and the N 2 O molecule, which have the following Lewis structures, respectively Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 88

Example • Strategy • Classify each species as AXm. En • Multiple bonds count as single bonds. It is the number of terminal atoms (X) that are counted, not the number of bonds Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 89

Example • Solution For Cl. O 3–: Species type Geometry For NO 3–: Species type Geometry A = Cl, X = O =3, E = 1 → AX 3 E trigonal pyramid, ideal bond angles are 109. 5° A = N, X = O = 3, E = 0 → AX 3 trigonal planar, ideal bond angles are 120° Copyright © 2016 Cengage Learning. All Rights Reserved. קשר כימי 8 - 90

Example For N 2 O: Species type Geometry A = N, X = N

Describing the Geometry of Glycine 92 קשר כימי 8 -

Dipoles • When two equal, but opposite, charges are separated by a distance, a dipole forms. • A dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated: =Qxr • It is measured in debyes (D). • 1 D = 3. 34 X 10 -30 C m. (coulomb-meters) 93 קשר כימי 8 -

In part A, there is no electric field; molecules are oriented randomly. A Dipoles In part B, there is an electric field; molecules align themselves against the field. B קשר כימי 8 - 94

Molecular Polarity The H─Cl bond is polar. The bonding electrons are pulled toward the Cl end of the molecule. The net result is a polar molecule. 95 קשר כימי 8 -

Molecular Polarity The O─C bond is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule. 96 קשר כימי 8 -

Molecular Polarity The H─O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. Because the molecule is bent, not linear, the net result is a polar molecule. 97 קשר כימי 8 -

Vector Addition 98 קשר כימי 8 -

Polarity of Molecules קשר כימי 8 - 100 Copyright © 2016 Cengage Learning. All

Bond and Molecular Polarity 101 קשר כימי 8 -

The dipole moment of a molecule can affect its properties cis-1, 2 -dichloroethene The net polarity is down; this is a polar molecule. Boiling point 60°C. trans-1, 2 -dichloroethene There is no net polarity; this is a nonpolar molecule. Boiling point 48°C. קשר כימי 8 - 102

Valence Bond Theory: A quantum mechanical model that shows how electron pairs are shared in a covalent bond 103 קשר כימי 8 -

Valence Bond Theory: A quantum mechanical model that shows how electron pairs are shared in a covalent bond 104 קשר כימי 8 -

Valence Bond Theory • Covalent bonds are formed by overlap of atomic orbitals, each of which contains one electron of opposite spin. • Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms. • The greater the amount of overlap, the stronger the bond. 105 קשר כימי 8 -

Hybridization and sp 3 Hybrid Orbitals How can the bonding in CH 4 be explained? 4 valence electrons 2 unpaired electrons 106 קשר כימי 8 -

Hybridization and sp 3 Hybrid Orbitals How can the bonding in CH 4 be explained? 4 valence electrons 2 unpaired electrons 107 קשר כימי 8 -

Hybridization and sp 3 Hybrid Orbitals How can the bonding in CH 4 be explained? 4 nonequivalent orbitals 108 קשר כימי 8 -

Hybridization and sp 3 Hybrid Orbitals How can the bonding in CH 4 be explained? 4 equivalent orbitals 109 קשר כימי 8 -

Hybridization and sp 3 Hybrid Orbitals 111 קשר כימי 8 -

Other Kinds of Hybrid Orbitals-sp 2 112 קשר כימי 8 -

Other Kinds of Hybrid Orbitals-sp 2 113 קשר כימי 8 -

Bond Rotation 114 קשר כימי 8 -

Double Bonds & Isomerism Bond rotation: CHCl=CHCl is “locked” - 2 isomers are possible. 115 קשר כימי 8 -

Double Bonds & Isomerism 116 קשר כימי 8 -

Other Kinds of Hybrid Orbitals-sp 117 קשר כימי 8 -

Other Kinds of Hybrid Orbitals-sp 118 קשר כימי 8 -

Consider a Molecule Like PCl 5 Phosphorous: valence electron configuration of 3 s 23 p 3. (5 electrons) Each of the five electrons forms a single bond with a chlorine atom. This means that central atom in the molecule needs 5 bonding orbitals to achieve the trigonal bipyramidal electronic geometry. This cannot happen with sp 3 hybridization… 3 s mix or “hybridize ” 3 p sp 3 hybrid valence bond orbitals קשר כימי 8 - 119

Valence Bond Theory (2): Expanded Valence The only way to produce 5 half-filled orbitals on phosphorous is by adding a fifth atomic orbital… 3 d 3 d 3 s 3 p mix the 3 s, the three 3 p and one 3 d new sp 3 d hybrid valence bond orbitals Each of these half–filled sp 3 d orbitals can form a –bond with a chlorine atom in PCl 5. קשר כימי 8 - 120

Additional Example of sp 3 d Hybrid Molecules SF 4 (sulfur tetrafluoride) EPG: Trigonal Bipyramidal MG: See Saw Cl. F 3 (chlorine trifluoride) EPG: Trigonal Bipyramidal MG: T-shape קשר כימי 8 - 121

Valence Bond Theory (2): Expanded Valence What about molecules with 12 electrons in the valence? In order to achieve an expanded valence that can hold six electron pairs (bp & lp) we need to form 6 new hybrid orbitals. This requires the mixing of an s, three p’s and two d–atomic orbitals. 3 d 3 d 3 p mix the 3 s, the three 3 p and two 3 d’s new sp 3 d 2 hybrid valence bond orbitals 3 s קשר כימי 8 - 122

sp 3 d 2 Hybridization: SF 6 3 d 3 d 3 p 3 s sulfur mix the 3 s, the three 3 p and two 3 d’s six new sp 3 d 2 hybrid valence bond orbitals Each of these half–filled sp 3 d 2 orbitals can form a –bond with a fluorine atom. קשר כימי 8 - 123

sp 3 d 2 Hybridization: SF 6 Sulfur has a valence electron configuration of 3 s 23 p 4 (6 electrons). Each of the six electrons forms a single bond with a fluorine atom forming an octahedral MG and EPG. The bonding can be described in terms of sp 3 d 2 hybrid orbitals. SF 6 קשר כימי 8 - 124

Molecular Orbital Theory: The Hydrogen Molecule Atomic Orbital: A wave function whose square gives the probability of finding an electron within a given region of space in an atom Molecular Orbital: A wave function whose square gives the probability of finding an electron within a given region of space in a molecule 125 קשר כימי 8 -

Molecular Orbital Theory: The Hydrogen Molecule σ bonding orbital σ* antibonding orbital Bond order = 126 (# bonding e– – # antibonding e–) 2 קשר כימי 8 -

Molecular Orbital Theory: The Hydrogen Molecule Bond order = 127 2– 0 =1 2 קשר כימי 8 -

Molecular Orbital Theory: The Hydrogen Molecule Bond order: 128 2– 2 =0 2 2– 1 = 1/2 2 קשר כימי 8 -

Interaction of p Orbitals 129 קשר כימי 8 -

Interaction of p Orbitals 130 קשר כימי 8 -

Interaction of p Orbitals 131 קשר כימי 8 -

O 2 • Dioxygen is paramagnetic. • Paramagnetic material has unpaired electrons. • Neither Lewis theory nor valence bond theory predict this result. 133 קשר כימי 8 -

O 2 as Described by Lewis and VB Theory 134 קשר כימי 8 -

Molecular Orbital Theory: Other Diatomic Molecules 135 קשר כימי 8 -

s and p Orbital Interactions • In some cases, s orbitals can interact wit the pz orbitals more than the px and py orbitals. • It raises the energy of the pz orbital and lowers the energy of the s orbital. • The px and py orbitals are degenerate orbitals. 136 קשר כימי 8 -

Heteronuclear Diatomic Molecules • Diatomic molecules can consist of atoms from different elements. • How does a MO diagram reflect differences? • The atomic orbitals have different energy, so the interactions change slightly. • The more electronegative atom has orbitals lower in energy, so the bonding orbitals will more resemble them in energy. 137 קשר כימי 8 -

Benzene The organic molecule benzene (C 6 H 6) has six -bonds and a p orbital on each C atom, which form delocalized bonds using one electron from each p orbital. 138 קשר כימי 8 -