Polar Covalent Bonds: Understanding Electronegativity, Bond Polarity, and Dipole Moments -, Study notes of Organic Chemistry

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Polar Covalent Bonds: Acids and Bases Based on McMurry’s Organic Chemistry, 6th edition, Chapter 2 2 2.1 Polar Covalent Bonds: Electronegativity  Covalent bonds can have ionic character  These are polar covalent bonds  Bonding electrons attracted more strongly by one atom than by the other  Electron distribution between atoms in not symmetrical 5 Bond Polarity and Electronegativity  Metals on left side of periodic table attract electrons weakly: lower electronegativities  Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly: higher electronegativities  Electronegativity of C = 2.5 6 Bond Polarity and Inductive Effect  Nonpolar Covalent Bonds: atoms with similar electronegativities  Polar Covalent Bonds: Difference in EN of atoms < 2  Ionic Bonds: Difference in electronegativities > 2 (approximately).  Other factors (solvation, lattice energy, etc) are important in ionic character. 7 Bond Polarity and Inductive Effect  Bonding electrons are pulled toward the more electronegative atom in the bond  C acquires partial positive charge, +  Electronegative atom acquires partial negative charge, -  Inductive effect: shifting of electrons in a bond in response to the electronegativities of nearby atoms 10 Polar Covalent Bonds: Dipole Moments  Dipole moment - Net molecular polarity, due to difference in summed charges   - magnitude of charge Q at end of molecular dipole times distance r between charges   = Q  r, in debyes (D)  1 D = 3.336  1030 coulomb meter 11 Dipole Moments in Water and Ammonia  Large dipole moments  Electronegativities of O and N > H  Both O and N have lone-pair electrons oriented away from all nuclei TABLE 2.1 Dipole Moments of Some Compounds Dipole moment Dipole moment Compound (D) Compound (D) NaCl 9.0 NH; 1.47 O CH. 0 ee , aS 3.46 ccl, 0 O- CH3CH3 0 Nitromethane O CHCl 1.87 0 H,0 1.85 Benzene CH;0H 1.70 BF; 0 + H,C=N=N- 1.50 Diazomethane ©2004 Thomson - Brooks/Cole 12 15 2.3 Formal Charges  Sometimes it is necessary to have structures with formal charges on individual atoms  We compare the bonding of the atom in the molecule to the valence electron structure  If the atom has one more electron in the molecule, it is shown with a “-” charge  If the atom has one less electron, it is shown with a “+” charge a formal Charges Number of Number of Formal charge =| valence electrons | — | valence electrons in free atom in bound atom Number of Half of Number of =| valence |-—j| bonding | — | nonbonding electrons electrons electrons 16  og, Nitromethane: H c() \ VA _-— Formal positive charge 4 v Nw __— Formal negative charge H H :O:7 Nitromethane ©2004 Thomson - Brooks/Cole 17 20 2.4 Resonance  Some molecules have structures that cannot be shown with a single Lewis representation  In these cases we draw Lewis structures that contribute to the final structure but which differ in the position of the  bond(s) or lone pair(s)  Such a structure is delocalized and is represented by resonance forms 21 2.4 Resonance  The resonance forms are connected by a double-headed arrow 22 Resonance Hybrids  A structure with resonance forms does not alternate between the forms  Instead, it is a hybrid of the two resonance forms, so the structure is called a resonance hybrid  For example, benzene (C6H6) has two resonance forms with alternating double and single bonds  In the resonance hybrid, the actual structure, all of the C-C bonds are equivalent, midway between double and single bonds 25 2.5 Rules for Resonance Forms  Individual resonance forms are imaginary - the real structure is a hybrid (only by knowing the contributors can you visualize the actual structure)  Resonance forms differ only in the placement of their  or nonbonding electrons  Different resonance forms of a substance don’t have to be equivalent  Resonance forms must be valid Lewis structures: the octet rule usually applies  The resonance hybrid is more stable than any individual resonance form would be 26 Curved Arrows and Resonance Forms  We can imagine that electrons move in pairs to convert from one resonance form to another  A curved arrow shows that a pair of electrons moves from the atom or bond at the tail of the arrow to the atom or bond at the head of the arrow Curved Arrows and 4 Resonance Forms The red curved arrow indicates that a lone pair of electrons moves from the top oxygen The new resonance structure atom to become part of an N=O double bond. has a double bond here... H O27 H :O: e O Simultaneously, two electrons from the CO and has a lone pair N=O double bond move onto the bottom of electrons here. oxygen atom to become a lone pair. © Thomson - Brooks Cole 27 Resonance in the acetone enolate 7 a i /\ /% H HH H Acetone © Thomson - Brooks Cole This resonance form has the This resonance form has the negative charge on carbon. negative charge on oxygen. (:0: 70: 9 I n | tron; = —— Bs, <> Be /\ | /\ | H H H H H H Acetone anion (two resonance forms) 30 31 2,4-Pentanedione  The anion derived from 2,4-pentanedione  Lone pair of electrons and a formal negative charge on the central carbon atom, next to a C=O bond on the left and on the right  Three resonance structures result + 2,4-Pentanedione enolate oe Gp) 7% cL Uc — OFC — C C ea ne" CH, H.C~ “e CH, H,C~ ~ ~CH, H H © Thomson - Brooks Cole 32 Practice Prob. 2.3: Draw at tree resonance forms: Unpaired electron Pentadienyl radical 35 Solution: Three-atom grouping —_—————.4 [iceman ornare ia 1 H C C._. JH H CL. C H H H H H H H ©2004 Thomson - Brooks/Cole Three-atom grouping ——K; "HH i tf H C.. « ~C H H Cc C H ©2004 Thomson - Brooks/Cole 36  | | Cc Gn. is ‘ow “eo on exodd Thomson! prooksicd le 37 40 The Reaction of HCl with H2O  When HCl gas dissolves in water, a Brønsted acid–base reaction occurs  HCl donates a proton to water molecule, yielding hydronium ion (H3O+) and Cl  The reverse is also a Brønsted acid–base reaction of the conjugate acid and conjugate base The Reaction of HCI with 4 _—™&™ H ——~ H H ci~ oO” = Cl ee No + conjugate acid \ conjugate I base a - H]* Acid Base Conjugate Conjugate acid base ©2004 Thomson - Brooks/Cole 41 H—A + :B == A:~ + H—B* Acid Base Conjugate Conjugate base acid ‘©2004 Thomson - Brooks/Cole O Oo | I Nor be + -7$-H == No SO + 58 /\ i | H H H F H Acid Base Conjugate Conjugate base acid ‘aie + os — 0-H + aoe H H H Acid Base Conjugate Conjugate base acid ©2004 Thomson - Brooks/Cole 42  a K, - the Acidity Constant HA + H,O == A- + H.O+ [H,;07][A7] [HA] K, = K,,|H20] = 46 2.8 Acid and Base Strength  The ability of a Brønsted acid to donate a proton to is sometimes referred to as the strength of the acid.  The strength of the acid can only be measured with respect to the Brønsted base that receives the proton  Water is used as a common base for the purpose of creating a scale of Brønsted acid strength 47 pKa – the Acid Strength Scale  pKa = -log Ka (in the same way that pH = -log [H+]  The free energy in an equilibrium is related to –log of Keq (G = -RT ln Keq = - 2.303RT log Keq)  A larger value of pKa indicates a stronger acid and is proportional to the energy difference between products and reactants 50 2.9 Predicting Acid–Base Reactions from pKa Values  pKa values are related as logarithms to equilibrium constants  The difference in two pKa values is the log of the ratio of equilibrium constants, and can be used to calculate the extent of transfer Predicting Acid-Base Reactions from pK, Values ‘Bo@ SA - + “0-H — se Le + 0-H “\ “ “1 oe | H H H H H Acetic acid Hydroxide ion Acetate ion Water (pK, = 4.76) (pK,= 15.74) © Thomson - Brooks Cole 51 Predicting Acid-Base Reactions 4 from pK, Values O O | J CH,COH + HO- =~ H,O + CH,CO- Stronger Stronger Weaker Weaker acid base acid base ©2004 Thomson - Brooks/Cole 52 Organic Acids H 0 H :0: Anion is stabilized by having See ay a, ae negative charge on a highly ~N -X electronegative atom. H H H H 10 1% tie | I | Anion is stabilized by H. CL... HH _y He Onin x Hy ox having negative charge C —_—-, C O: —_— Cc QO: on a highly electronegative gen “ /\ fi atom and by resonance. H H H H H H © Thomson - Brooks Cole oO :O: Anion is stabilized I I | by resonance and H Cc H H C_= ,H H Cc H by having negative ~~ “oT a, “ow “a7 — a ct So < charge on a highly /\ i% /\ | /\ | electronegative aA H i Y H H H H H H atom. homson - Brooks Cole 55  2 Poxylic Acids: O | CH, —C—OH Acetic acid ©2004 Thomson - Brooks/Cole Oo O | ll CH; —C—C—OH Pyruvic acid i a | HO—C—CH, ¢ CH,—C—OH CO.H Citric acid 56 Organic Acids © & & Some organic Ae Ok. acids C 4 H H Methyl alcohol pK,= 15.54 ©2004 Thomson - Brooks/Cole fs H H Acetic acid pK,= 4.76 57 Organic Bases | 1 Some organic H N. H 0 H Cc H awa is das ge “oe er “ i “sf H 4H H H H HH H Methylamine Methy] alcohol Acetone ©2004 Thomson - Brooks/Cole 60 61 2.11 Acids and Bases: The Lewis Definition  Lewis acids are electron pair acceptors; Lewis bases are electron pair donors  The Lewis definition leads to a general description of many reaction patterns, but there is no quantitatve scale of strengths as in the Brønsted definition of pKa 62 Lewis Acids and the Curved Arrow Formalism  The Lewis definition of acidity includes metal cations, such as Mg2+  They accept a pair of electrons when they form a bond to a base  Group 3A elements, such as BF3 and AlCl3, are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases  Transition-metal compounds, such as TiCl4, FeCl3, ZnCl2, and SnCl4, are Lewis acids eles Acid/Base Reaction: . : : Boron Dimethyl Lewis acid-base trifluoride ether complex © 2004 Thomson/Braoks Cole 65 aerome Lewis Acids: Some Lewis acids Some neutral proton donors: H,O HCl HBr HNO, 4H,SO, O | OH a“ CL H;C OH CH;CH,OH A carboxylic acid A phenol An alcohol Some cations: Lit Mg?* Br’ Some metal compounds: AIC BF; TiCl, FeCl; ZnCl, ©2004 Thomson - Brooks/Cole 66 67 Lewis Bases  Lewis bases can accept protons as well as other Lewis acids, therefore the definition encompasses that for Brønsted bases  Most oxygen- and nitrogen- containing organic compounds are Lewis bases because they have lone pairs of electrons  Some compounds can act as either acids or bases, depending on the reaction ie Midazole (Prob. 2.19): H nA N—H Imidazole — H AY H ©2004 Thomson - Brooks/Cole 70 71 2.12 Drawing Chemical Structures  Condensed structures: C-H and C-C and single bonds aren't shown but understood  If C has 3 H’s bonded to it, write CH3  If C has 2 H’s bonded to it, write CH2; and so on.  Horizontal bonds between carbons aren't shown in condensed structures—the CH3, CH2, and CH units are assumed to be connected horizontally by single bonds, but vertical bonds are added for clarity 4 2-methylbutane Structures i —cC —H Condensed structures H CH | i —C—C—H = CH;CH,CHCHs; or CH;CH»CH(CHs3). | | H H 2- i Z—-O— t—A—F Methylbutane © Thomson - Brooks Cole 72  TABLE 2.4 Kekulé and Skeletal Structures for Some Compounds Skeletal Compound Kekulé structure structure i H: H Isoprene, C;H, i H Cc Cc J ~ ce ae ~ H a ee H H Methylcyclohexane, C;H,, Phenol, C,H,O ©2004 Thomson - Brooks/Cole ‘ UO OH f 75 76 Practice Prob. 2.7: How many H’s on each carbon? GE Solution: OH 2H 2H O 3HiA OH os 1H 3H 2H ©2004 Thomson - Brooks/Cole Carvone, C,9H,4,0 77 80 2.13 Molecular Models  We often need to visualize the shape or connections of a molecule in three dimensions  Molecular models are three dimensional objects, on a human scale, that represent the aspects of interest of the molecule’s structure (computer models also are possible)  Drawings on paper and screens are limited in what they can present to you 81 2.13 Molecular Models  Framework models (ball- and-stick) are essential for seeing the relationships within and between molecules – you should own a set  Space-filling models are better for examining the crowding within a molecule Space-filling Framework 82 Summary  Organic molecules often have polar covalent bonds as a result of unsymmetrical electron sharing caused by differences in the electronegativity of atoms  The polarity of a molecule is measured by its dipole moment, .  (+) and () indicate formal charges on atoms in molecules to keep track of valence electrons around an atom  Some substances must be shown as a resonance hybrid of two or more resonance forms that differ by the location of electrons.