1 Structure and Bonding 1 n Organic Chemistry

1. Structure and Bonding 1

n • Organic Chemistry : Chemistry of Compounds of Carbon n • Review Electronic Structure of the Atom : Quantum Mechanical Model or Schrodinger Model; atomic orbitals (s, p, d, f); energies of the orbitals; aufbau principle (Pauli Exclusion Priniciple, Hund’s Rule). n • Review Octet Rule; Ionic Bonds, Covalent Bonds; n • C is tetravalent n • Review Lewis Structures=Electron-dot structures; n Kekule. Structures=Line-bond structures; bond pairs and lone pairs; n • Review VSEPR Theory; Molecular Geometries, Bond Angles 2

Bonds: Ionic bonds, Covalent Bonds, Metallic Bonds Covalent Bond: Sharing of pairs of electrons between 2 atoms Single Bond: sharing of 1 pair Double Bond: sharing of 2 pairs Triple Bond: sharing of 3 pairs Covalent Bonds hold atoms together in molecular compounds and in polyatomic ions. Ionic Bond: Attraction between a positively charged ion (cation) and a negatively charged ion (anion) in ionic compounds Tetravalent: forms 4 bonds : C, Si Trivalent: forms 3 bonds: N, P Divalent: forms 2 bonds: O, S Monovalent: forms 1 bond: Halogens, H Octet : 8 valence electrons; filled valence s and p subshells 3

Electronic Structure of the Atom • Electrons are arranged in atomic orbitals starting with the lowest energy orbital: aufbau principle • Arrangement of electrons having the lowest energy: ground state electron configuration. • Partial list of Atomic orbitals in increasing order of energy: 1 s < 2 p < 3 s < 3 p < 4 s < 3 d < 4 p < 5 s < 4 d < 5 p < 6 s • There is only 1 orbital of each of the following: 1 s, 2 s, 3 s, 4 s, 5 s, 6 s, 7 s • There are 3 orbitals of each of the following: 2 p, 3 p, 4 p, 5 p, 6 p, 7 p • There are 5 orbitals of each of the following: 3 d, 4 d, 5 d, 6 d • Pauli Exclusion Principle limits the maximum number of electrons in any orbital to 2 electrons. • The 3 orbitals of 2 p are said to make up the 2 p subshell. • Similarly, the 5 orbitals of 3 d are said to makeup the 3 d subshell. • Orbitals in a subshell have the same energy: they are said to be degenerate. • For example, all 3 2 p orbitals have the same energy; all 5 3 d orbitals are degenerate • Hund’s rule: Degenerate orbitals are filled 1 electron at a time before 4 pairing

• When atoms form covalent bonds, they share valence electrons. • Lewis Symbols of an atom: Its chemical symbol with all valence electrons represented as dots. The dots can be placed on any or all 4 sides of the symbol. Dots may be placed in pairs or singly. • Number of valence electrons for s & p block elements = group no. in the North American scheme (not the IUPAC) • For example, Lewis symbol of H: H●; Lewis symbosl of alkaline earth metals have 2 dots; Lewis symbols of halogens have 7 dots; lewis symbol of C has 4 dots etc. • For molecules & polyatomic ions, Lewis structures are drawn combining the Lewis symbols of the constituent atoms. • In Lewis structures, shared electron pairs are represented by lines connecting 2 atoms. Therefore, a single bond is represented by 1 line (), a double bond by 2 parallel lines (=) and a triple bond by three lines ( ). Nonbonding electrons or lone pairs are left as dots. 5

Rules for Drawing Lewis Structures: • Add up the valence electrons for all constituent atoms. For cations subtract as many electrons from this total as the charge. For anion, add as many electrons to the total as the charge. For example, the total number of valence electrons in CH 4 is 4 + (1 x 4) = 8; the total number of valence electrons in CO 3 -2 is 4 + (3 x 6) + 2 = 24; the total number of valence electrons in CH 3+ is 4 + (1 x 3) -1 = 6 • Connect all atoms to the central atom (there may be more than 1) usingle bonds. Recall: C is always a central atom; H & F are always end atoms • Complete the octets of all end atoms except H • Put any remaining electrons on the central atom • If the central atom still does not have an octet, make one or more lone pairs (if available) on end atoms bond pairs between the end atom & the central atom till octet is achieved. • Resonance structures are equivalent Lewis structures that differ ONLY in the distribution of multiple bonds (double, triple). The connectivity of atoms is not different. • When more than 1 resonance structure can be drawn for a formula, the actual structure is a hybrid of all the resonance structures and is called a resonance hybrid. 6

• For example, the actual structure of the carbonate ion is a hybrid of all three resonance structures (drawn in your notes). The result is that all three CO bonds are identical. They are longer and weaker than a pure C=O and they are shorter and stronger than a pure C-O. • Often, one resonance structure may contribute more to the hybrid than the other remaining resonance structures. Such a resonance structure is called a major resonance structure. The remaining structures are called minor resonance structures. In the case of the carbonate ion, all three resonance structures contribute equally. But this is not the case for the cyanate ion, CNO-. • The resonance structures of the cyanate ion are O C-N (I); O=C=N (II); and O-C N (III) • Formal charges are required to distinguish between major and minor resonance structures when they do not contribute equally. • The formal charge of an atom in a Lewis structure = Number of valence electrons in the isolated atom – Number of electrons assigned to the atom in the Lewis structure. = Number of valence electrons – (No. of dots on atom + number of lines attached to atom) • The major resonance structure is one with the maximum number of atoms with zero formal charge. Therefore the major resonance structure of CNO- is either structure II or III or both. The major resonance structure is also one in which negative formal charges are on more electronegative atoms. Therefore III is 7 the major resonance structure.

Lewis structures with formal charges larger than ± 1 are not satifactory representations. Therefore structure I is not a satisfactory Lewis structure for the CNO- because the formal charge of N is -2. 8

Valence Shell Electron Pair Repulsion Theory: VSEPR Theory VSEPR theory allows the following: Lewis Structure Predict the Molecular Geometries (shapes of molecules) The electron pairs around a central atom (A) in a Lewis structure arrange themselves so as to minimize the electrostatic repulsions among them. The resulting arrangement is called ELECTRON DOMAIN GEOMETRY Some of the electron pairs around A may be shared pairs and the remaining lone pairs. If we ignore the lone pairs and consider just the arrangement of bond pairs, the resulting geometry is called MOLECULAR GEOMETRY and that is the shape of the molecule. The angle between bond pairs is called the bond angle. 9

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal planar bent AB 3 3 0 trigonal planar AB 2 E 2 1 trigonal planar 10. 1

VSEPR Class AB 4 # of atoms bonded to central atom # lone pairs on central atom 4 0 AB 3 E 3 AB 2 E 2 2 1 Arrangement of electron pairs Molecular Geometry tetrahedral trigonal pyramidal 2 bent O H H

n • Review depiction of tetrahedral geometry using solid, dashed, wedged lines. Solid : in plane; wedged: in front; dashed: behind n • Review Valence Bond Theory; hybridization; sigma ( ) and pi ( ) bonds. n • Discuss Bonding, hybridization in ethane, ethylene, acetylene n http: //www. mhhe. com/physsci/chemistry/essentialche mistry/flash/hybrv 18. swf n http: //www. mhhe. com/physsci/chemistry/animations/ch ang_7 e_esp/bom 5 s 2_6. swf n • Review Molecular Orbital Theory; Bonding and Antibonding MOs n • Review Condensed Structural Formulas, Skeletal Structural Formulas 12

• Valence bond theory – bonds are formed by sharing of e- from overlapping atomic orbitals. • In the context of valence bond theory, hybridization is introduced to explain the geometries of molecules. • Hybridization is the mixing of 2 or more VALENCE orbitals on an atom to create AN EQUAL NUMBER of DEGENERATE (having the same energy) HYBRID ORBITALS. • The type of hybridization is determined by the sum of the lone pairs + number of atoms attached to the atom in question. 13

How do I predict the hybridization of the central atom? 1. Draw the Lewis structure of the molecule. 2. Count the number of lone pairs AND the number of atoms bonded to the central atom # of Lone Pairs + # of Bonded Atoms Hybridization 2 (AB 2) sp 3 (AB 3, AB 2 E) sp 2 4 (AB 4, AB 3 E, AB 2 E 2) sp 3 Examples Be. Cl 2 CH 3+, CH 2 CH 4, NH 3, H 2 O 10. 4

1. 11 Hybridization of Nitrogen and Oxygen n Elements other than C can have hybridized orbitals n H–N–H bond angle in ammonia (NH 3) 107. 3° n N’s orbitals (sppp) hybridize to form four sp 3 orbitals n One sp 3 orbital is occupied by two nonbonding electrons, and three sp 3 orbitals have one electron each, forming bonds to H 15

Hybridization of Oxygen in Water n The oxygen atom is sp 3 -hybridized n Oxygen has six valence-shell electrons but forms only two covalent bonds, leaving two lone pairs n The H–O–H bond angle is 104. 5° 16

1. 12 Molecular Orbital Theory n A molecular orbital (MO): where electrons are most likely to be found (specific energy and general shape) in a molecule n Additive combination (bonding) MO is lower in energy n Subtractive combination (antibonding) forms MO is higher 17

Molecular Orbitals in Ethylene n The bonding MO is from combining p orbital lobes with the same algebraic sign n The antibonding MO is from combining lobes with opposite signs n Only bonding MO is occupied 18

Summary n n n Organic chemistry – chemistry of carbon compounds Atom: positively charged nucleus surrounded by negatively charged electrons Electronic structure of an atom described by wave equation n n Electrons occupy orbitals around the nucleus. Different orbitals have different energy levels and different shapes n n n n Covalent bonds - electron pair is shared between atoms Valence bond theory - electron sharing occurs by overlap of two atomic orbitals Molecular orbital (MO) theory, - bonds result from combination of atomic orbitals to give molecular orbitals, which belong to the entire molecule Sigma (s) bonds - Circular cross-section and are formed by head-on interaction Pi ( ) bonds – “dumbbell” shape from sideways interaction of p orbitals Carbon uses hybrid orbitals to form bonds in organic molecules. n n s orbitals are spherical, p orbitals are dumbbell-shaped In single bonds with tetrahedral geometry, carbon has four sp 3 hybrid orbitals In double bonds with planar geometry, carbon uses three equivalent sp 2 hybrid orbitals and one unhybridized p orbital Carbon uses two equivalent sp hybrid orbitals to form a triple bond with linear geometry, with two unhybridized p orbitals Atoms such as nitrogen and oxygen hybridize to form strong, oriented bonds n The nitrogen atom in ammonia and the oxygen atom in water are sp 3 -hybridized 19