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Chapter 1 - 1 Chapter 1 ¥ 3 STRUCTURE AND PROPERTIES IN...

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Unformatted text preview: 1, Chapter 1 ¥ 3.; STRUCTURE AND PROPERTIES IN THIS CHAPTER: V Carbon Compounds V Functional Groups V Formal Charges and Lewis Dot Structures V Atomic Orbitals V Hybridization and Bonding V Electronegatlvlty and Polarity V Resonance and Delocalized TC Electrons V Solved Problems Carbon Compounds Organic chemistry is the study of carbon (C) compounds. molecules which have covalent bonds. Carbon atoms can bond to each other to form open-chain compounds, or cyclic (ring) compounds. Both types 1 2 ORGANlC CHEMISTRY can also have branches of C atoms. Saturated compounds have C’s bonded to each other by single bonds, C—C; unsaturated compounds have C‘s joined by multiple bonds. Examples with double bonds and triple bonds are shown below. Cyclic compounds having at least one atom in the ring other than C (a heteroatom) are called heterocyclics. The heteroaloms are usually oxygen (0), nitrogen (N), or sulfur (S). Most carbon-containing molecules are three~dimensional. 1n methane. the bonds of C make equal angles of 109.5° with each other, and each of the four H's is at a vertex of a regular tetrahedron whose center is occupied by the C atom. Other shapes do occur: ethene. for example, is planar, and ethyne (acetylene) is linear. T H 1095" / l H H \H'>H Methane is a tetrahedron 120° 180° H H \ f/ . fl /C_C\>120 H—CEC—H H H Ethene Ethyne Organic compounds show a widespread occurrence of isomers, which are compounds having the same molecular formula but different struc~ tural formulas. Isomers have different chemical and physical properties. This phenomenon of isomerism is exemplified by isobutane and n- butane. The number of isomers increases as the number of atoms in the molecule increases. are CHgCHZCHZCH3 CHgCHCHa n-Butane Isobutane CHAPTER 1: Structure and Properties 3 Functional Groups Hydrocarbons contain only C and hydrogen (H). HS in hydrocarbons can be replaced by other atoms or groups of atoms. These replacements, called functional groups, are the reactive sites in molecules. Double and triple bonds are considered to be functional groups. Some common functional groups are given in the Functional Group Table. The “R" group is a generic group. and is not part of the functional group of inter~ est. Compounds with the same functional group form a homologous series having similar chemical properties and often exhibiting a regular gradation in physical properties with increasing molecular weight. Formal Charge The formal charge on a covalently bonded atom equals the number of valence electrons of the unbonded atom minus the number of electrons assigned to the atom in its bonded state. The assigned number is one half the number of shared electrons plus the total number of unshared electrons. The sum of all formal charges in a molecule equals the charge on the species. In this outline formal charges and actual ionic charges are both indicated by the signs + and — . The structures shown below are called Lewis dot structures (or simply dot structures). Each dot repre— sents an electron in the outer shell of the atom. These drawings can be highly useful in determining if an atom bears a formal charge. 1-1 G) 'N‘ HINfiH HINIH H H Neutral N Neutral NH3 Ammonium atom (5 e‘) (5 e‘ from N, cation (8 total e" 3 e‘ from HS) 4 belong to N) 4 ORGANIC CHEMISTRY Some Common Functional Groups Alkane -(CH2)— R R 0:0 Alkene R H _ = _‘ ' Alkyne R (3—0 R NOte' Atcohol R—OH T? C Ketone R/ \Rx if C \ Aldehyde R/ H W R/C\OH Carboxylic Acid if C / \ . Ester R OR T’ C / \ R NR'R" Amide Nitrile R—CEN CHAPTER 1: Structure and Properties 5 Atomic Orbitals An atomic orbital (A0) is a region of space about the nucleus in which there is a high probability of finding an electron. For organic molecules, the atomic orbitals of most interest are the s orbital and the p orbitals. The s orbital is a sphere around the nucleus. as shown below. A p orbital has 2 lobes touching on opposite sides of the nucleus. The three p orbitals are labeled px. py. and p1 because they are oriented along the x—, y—, and z-axes. respectively . In a p orbital there is no chance of find— ing an electron at the nucleus—the nucleus is called a node. Three principles are used to distribute electrons in orbitals. l. “Aufbau” or building-up principle. Orbitals are filled in order of increasing energy: 1.9.25, 2]), 3s, 3p, 45, 3d, 4p, 53, 4d, 5p, 65, 4f. 5d. 6p, etc. 2. Pauli exclusion principle. No more than two electrons can occupy an orbital and then only if they have opposite spins. 3. Hund’s rule. When filling orbitals of equal energy, place one elec- tron in each orbital (using parallel spins) before pairing electrons. (Substances with unpaired electrons are paramagnetic—they are attracted to a magnetic field.) Hybridization and Bonding A carbon atom must provide four equal-energy orbitals in order to form four equivalent bonds, as in methane, CH4. It is assumed that the four equivalent orbitals are formed by blending the 25 and the three 2p AO’s 6 ORGANK30HEMBTRY to give four new hybrid orbitals, called sp3 orbitals. The larger lobe, the “head." having most of the electron density. overlaps with an orbital of its bonding mate to form the bond. The smaller lobe. the “tail," is often omitted when depicting hybrid orbitals. However. at times the “tail” plays an important role in an organic reaction. The s and p orbitals of carbon can hybridize in ways other than sp3, as shown below. Repulsion between pairs of electrons causes these hybrid orbitals to have the maximum bond angles. The sp2 and sp hybrid orbitals induce geometries about the CS as shown below. 2 54‘ SP3 Sp2 Sp (hybrid orbials shown in white above) Head—to-head overlap of AO’s gives a sigma (5) bond. The bond angles (angles between s-bonds) at sp3 carbons are 109.5”, leading to a tetra— hedral geometry. The bond angles at sp2 carbons are 120°. leading to a trigonal planar geometry. and the bond angles at sp carbons are 180°, leading to a linear geometry. The imaginary line joining the nuclei of the bonding atoms is the bond axis, whose length is the bond length. Two parallel p orbitals overlap side—by—side to form a pi (TE) bond. The bond axis lies in a nodal plane (plane of zero electronic density). Single bonds are 0 bonds. A double bond is one 6 and one 1: bond. A triple bond is one 6 and two It bonds. 0 O C) CC 55 Overlap of 2 p orbitals creates an: bond CHAPTER 1: Structure and Properties 7 Electronegativity and Polarity The electronegativity of an atom is its tendency to attract bonding elec- trons toward itself. The higher the electronegativity, the more strongly the atom attracts and holds electrons. A nonpolar covalent bond exists between atoms having a very small or zero difference in electronegativ- ity. A few relative electronegativities are F (4.0) > O (3.5) > C1, N (3.0) > Br (2.8) > S, C. I (2.5) > H (2.1) A bond formed by atoms of dissimilar electronegativities is called polar due to partial charge separation. The more electronegative ele— ment of a covalent bond is relatively negative in charge, while the less electronegative element is relatively positive. The symbols 8+ and 8— represent partial charges in polar bonds. These partial charges should not be confused with ionic charges. Polar bonds are indicated by +—>; the arrow points toward the more electronegative atom. ThelVectdt‘ stint of alllfihdifiidual bond moments gives the net] dipole moment of the molecule H 20 has polar bonds. Since the molecule has :a bent shape; theE “ dipoles of the bonds do not cancel?) and the molecule has a net dipole moment Resonance and Delocalized rr Electrons Resonance theory describes species for which a single structure does not adequately describe the species’ properties. As an example, consider the cation on the next page (called the ally] cation): 8 ORGANIC CHEMISTRY A comparison of the calculated and observed bond lengths shows that the 2 C—C bonds are the same length. Neither resonance structure alone can explain this similarity in bond length. When resonance struc- tures form a resonance hybrid, we obtain a structure consistent with the observed bond length. The resonance hybrid has some double-bond character between the central carbon and both outside carbons. This state of affairs is described by the non-Lewis structure in which dotted lines stand for the partial bonds in which there are delocalized 1t elec- trons in an extended It bond created from overlap of p orbitals on each atom. The symbol <—> denotes resonance. nor equilibrimn. ti *3 r H\C,’C\C,H H\C,C\\C,H chf’cgcfi | ICE 9| I 599 I .563 H H H H H H Hybrid The hybrid is more stable than any single resonance structure. The more nearly equal in energy the contributing structures, the greater the reso- nance energy. When contributing structures have dissimilar energies, the hybrid looks most like the lowest—energy structure. Contributing structures (a) differ only in positions of electrons (atomic nuclei must have the same positions): (b) must have the same number of paired elec— trons; and (c) must not place more than 8 electrons on any second peri- od atom. Relative energies of contributing structures are assessed by the following rules. 1. Structures with the greatest number of covalent bonds are most sta- ble. However, for second—period elements (C. O, N) the octet rule must be observed. . With few exceptions, structures with the least charge separation are most stable. 3. If all structures have formal charge. the most stable (lowest energy) one has — on the more electronegative atom and + on the more elec- tropositive atom. 4. Structures with like formal charges on adjacent atoms have very high energies. 5. Resonance structures with electron-deficient. positively charged atoms have very high energy, and are usually ignored. [\J CHAPTER 1: Structure and Properties 9 : You Need to Know iLewrs Doti Structures Formal Charges Hybridization ' “ Molecular Geometry Resonance “ Solved Problems Problem 1.1 Find the formal charge on each element of ArBFS. and find the net charge on the species. :Ar— 8— F: 'F: Atom Group Unshared 1/2 Shared Formal Charge Electrons Electrons F 7 6 l 0 B 3 O 4 —l Ar 8 6 1 +1 0 = net charge Problem 1.2 (a) NO:+ is linear, (b) N02“ is bent. Explain in terms of the hybrid orbitals used by N. (a) NO; (structure). N has two 0 bonds, no unshared pairs of electrons and therefore needs two hybrid orbitals. N uses Sp hybrid orbitals and the 0’ bonds are linear. The geometry is controlled by the arrangement of the sigma bonds. (b) NO; (structure). N has two 5 bonds, one unshared pair of elec— trons, and therefore, needs three hybrid orbitals. N uses spg hybrid orbitals and the bond angle is about 120”. ...
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