Words: 3227
1. 1 CHEMISTRY 2213a ORGANIC CHEMISTRY FOR THE LIFE SCIENCES – organic chemistry is the study of life at the molecular level; to many it is the key to understanding life “The language of chemistry- an international language, a language without dialects, a language for all of time, and a language that explains where we came from, what we are, and where the physical world will allow us to go” (Nobelist Arthur Kornberg, a biochemist, 2000) – but its study has been challenging for students for centuries “Organic chemistry nowadays almost drives me mad.
To me it appears like a primeval tropical forest full of the most remarkable things, a dreadful endless jungle into which one does not dare to enter for there seems to be no way out” (Freidrich Wohler, 1835 – he was the first to synthesize an organic molecule, urea, from inorganic materials in 1828) – but YOU can enjoy it if you make the effort to understand it. It is NOT “rocket science” 1. 2 1. COVALENT BONDING & SHAPES OF MOLECULES [text 1. 1-1. 7] ( ?? Chemistry Dept, University of Western Ontario, 2011) – largely a review of essential material from year-1 chem A.
Electronic Structure of Atoms – The bonding behaviour of atoms depends entirely on electron configuration, as revealed by an atom’s position in the Periodic Table. – The “Organic Chemist’s” periodic table. Key to numbers: – upper left = atomic number = number of electrons – lower left = number of outer shell electrons (= group number) – upper right = valence = number of unpaired electrons in valence shell – lower right = electronegativity value [a full periodic table with atomic numbers (only) is provided on tests] 1. 3 intro chem showed us how quantum numbers fix the identity of electrons in atoms, ie, 1. Principal Quantum Number, symbol n n may have any positive integral value ; 0, e. g. 1, 2, 3, 4… As n increases, the energy of the electron and its distance from the nucleus increase. Within each principal level there are sublevels denoted by 2. Orbital Quantum Number, symbol R R may have integral values from 0 to (n ! 1); thus if n=2, R can only have values 0 and 1. If n=1, only R=0 is allowed. The general shape of the electron cloud is determined by R.
We have ‘names’ for electrons with different values of R : R = 0, R = 2, ‘s electrons’ ‘d electrons’ R = 1, R = 3, ‘p electrons’ ‘f electrons’ 3. Magnetic Quantum Number, symbol m R m R may have integral values from ! R through zero to + R. thus if R = 1, a p electron, m R can be -1, 0, or +1. The magnetic quantum number sets the number of suborbitals for each value of R. 4. Spin Quantum Number, symbol m s Regardless of the values of the other quantum numbers, m s may be only +? or -? . when two electrons come together, their spins may be either the same (both + or both ! or different (one + and one ! ). We 1. 4 refer to these cases as, respectively, parallel spins and opposite spins, and they are usually represented by arrows pointing up or down ? ? (parallel, not allowed) ? ? (opposite, allowed). Ground State Electron Configuration – the lowest energy state for electrons in an atom can be determined by placing electrons in pairs into atomic orbitals, filling orbitals in succession from lowest energy to highest, ie, But, the 4s and the 3d are very close in energy – this approach works for the first several elements, such as 17Cl, whose configuration is .
It gives correct electron configurations for all the atoms normally encountered in organic chem (see B, Table 1. 3) – it is easy therefore to write out the valence electron structure, called the Lewis structure, for the first 18 atoms of the periodic table. Their valence shell electron configuration (Brown, Table 1. 4) can be shown as 1. 5 – NB: for elements in groups IIA, IIIA and IVA, the two paired electrons become unpaired before forming bonds so the valence of these elements becomes equal to the total number of valence electrons, eg C=4, Mg=2, etc B.
Ionic, Covalent and Polar Bonding Electronegativity reflects an atom’s ability to “attract a shared pair of electrons to itself”. – electronegativity values given on p 1. 2, but knowing the general trend is often enough, ie, j jj – electronegativity differences account for the characteristics of 3 main types of bond: 1) Ionic Bonds: result from the complete transfer of an electron from the less electronegative to the more electronegative atom. eg Na + Cl- 1. 6 occurs when electronegativity difference between bonded atoms > 1. 9 i. e. > 3 groups apart in the periodic table (this is not a hard and fast rule: if the electronegativity difference is large the bond will be very polar or ionic) 2) Polar Bonds: result from the unequal sharing of electrons between unlike atoms with electronegativity differences 0. 5, i. e. < 3 groups apart in the periodic table eg, where *+ = small +ve charge and *2 = small -ve charge but, j NB: rank H with C for electronegativity values molecules with polar bonds are usually also polar molecules, but sometimes the dipoles of polar bonds cancel each other out, so the molecule itself is non-polar (has no net dipole moment), eg, H 2O is polar; CO 2 is non-polar – this means that the 3D structure of a molecule needs to be known before its overall polarity can be recognized; structure covered in more detail later 3) Covalent Bonds: result from equal sharing of electrons between atoms – occurs when electronegativity difference < 0. 5 1. 7 eg, C-C, C-H, Cl-Cl C. Bond Cleavage All bonds can be broken (by heat, light, x-rays, etc. ) into separate fragments.
These fragments can be either: 1) Charged ions (+ve cations or -ve anions) in which the shared electron pair is transferred to the more electronegative atom, eg, or, 2) Radicals (neutral species with one unpaired electron) in which one electron of the shared bond goes to each atom, eg, D. Lewis Structures jj Drawing Lewis Structures 1. Text, pp. 12-13; a method widely used, or 2. A more rational method (which assumes atoms carry their own electrons to molecular bonds), taught in yr-1 i) analyse the molecular formula to determine which atoms are bonded to which ii) write out the central atom in its normal Lewis structure 1. 8 ii) attach each atom bonded to the central one, one-by-one, where each attached atom begins in its normal Lewis structure and you a) attach each atom by a single bond using unpaired electrons in each of the bonded atoms b) if necessary, break apart an electron pair on the central atom to form bonds to the attached atom c) attach remaining unpaired electrons on the bonded atoms to form double or triple bonds iv) add one electron for each -ve charge, subtract one for each +ve charge, always adding or removing electrons that will result in filled valence shells v) assign formal charges to the atoms which bear them vi) repeat the process for each atom in a molecule vii) remember, atoms in the first full row of the periodic table (eg, C,N,O) cannot have more than an octet of electrons, second row elements can have more, eg, P can have 10, and S can have 12. Some examples. CH 3ONH 2 CH 2CHCN H 2PO 42 CH 2ClCONH 3+ 1. 9 E. Formal Charge – an atom in a molecule is assigned a formal charge if the number of electrons “belonging to” that atom differs from the number around it in its neutral, atomic state. *(“belonging to” = all non-bonded electron pairs + 1 electron for each covalent bond in the valence shell) eg. 8 O species = 1s 2 2s 2 2p 4 = 6 valence electrons structure charge on O # electrons “belonging” to O 6 6 atomic O H2O H3O+ CH3OF. Structural Formulas none none +1 7 -organic chemists use various short-hand techniques to show the location of electrons and bonds in molecules.
The different drawing conventions, beginning from the Lewis dot structure to the various representational schemes, are shown below for ethanol, CH 3CH 2OH. full Lewis dot structure (all valence electrons shown): 1. 10 Lewis structure, with covalent bonds as lines: covalent bonds omitted, but assumed; either with or without non-bonding pairs: (in wide use for simple molecules) C???C bonds shown only by a line, H atoms bonded to C not shown: Such line diagrams only show bonding sequences present in molecules and are not designed, or intended, to show actual 3-dimensional structure. The line diagrams are most often used in larger, ring molecules, eg, the natural flavour vanillin, C 8H 😯 3 G.
Molecular Shapes – the shapes of organic molecules are well-predicted by VSEPR theory (year-1), and the basic shapes are in B&P, Table 1. 8 or, 1. 11 – remember, lone pairs of electrons exert a greater repulsive force than pairs in bonds (and so reduce bond angles); double and triple bonds act in VSEPR predictions like single bonds, eg, NH 3 CH 2O – by applying VSEPR shapes to all bonded sites in a molecule, the correct shape of nearly every organic molecule can be predicted and drawn in three dimensions – the most common way of depicting 3-dimensional structures is the “dot-line-wedge” symbolism, where i) a line, , represents a bond in the plane of the paper i) a dotted line, , represents a bond directed behind the plane of the paper iii) a wedged line, , represents a bond directed in front of the plane of the paper Thus, a 3D representation of ethanol is depicted as The best 3-dimensional representation of ethanol is given by molecular models (or artistically drawn ball & stick models), which depict ethanol as below left. If one wishes to get an idea of the molecule in which its electronic field is indicated by spheres 1. 12 representing the outermost filled electronic orbitals, a “spacefilling” representation, such as below right, can be used (C atoms = gray, H atoms = white; O atoms = red). – we can use our knowledge of polar bonds and molecular shape to decide whether molecules with several polar bonds are polar or non-polar, eg, CH 2Cl2 CCl4
Looking Ahead: if you study well enough to be able to do the past test questions, with understanding not memorization, at the rate of 4 min per question, you should be well-prepared for tests Past Test Q: 1) Assuming filled valence shells for each atom, what is the overall charge on the molecule shown? A) -2 B) -1 C) 0 D) +1 E) +2 1. 13 Illustrative MCAT Q: 1) What determines the polarity of a covalent bond? A) difference in atomic size B) difference in electronegativity C) difference in total number of protons D) difference in total number of valence electrons H. Resonance Structures The nitrate ion, NO 32, can be drawn in 3 equivalent Lewis structures, termed contributing (or resonance) structures, ie, – note that such equivalent contributing structures differ from one another only in the position of the electrons.
In each of the resonance structures the double bond (and the -ve formal charges) reside on different atoms. – molecules and/or ions whose electron distributions are best represented by two or more resonance structures have greater stability than expected from average bond energy values alone. – resonance structures should always be written if more than one energetically equivalent valence bond structure is possible. eg, the acetate ion, CH 3COO G 1. 14 – in molecules/ions for which resonance structures can be drawn, no one Lewis dot structure is correct by itself. In practice each is considered to be ‘contributing’ to the actual electronic structure and the o symbol is used exclusively, in chemistry) to indicate this. The best known example of contributing (or resonance) structures is benzene, C 6H 6 – molecular stabilization is greatest in molecules whose resonance structures are energetically equivalent, as in the above examples, but molecules/ions are even stabilized (but to a lesser extent) by non-equivalent contributing structures, eg, OCN 2 O=C=N O???C / N j in general, the existence of resonance is indicative of increased stability. This is especially true in the delocalization of negative charge over several atoms, as in CO 32- j resonance is responsible for the stability and geometry of peptide (= amide) linkages in proteins, ie,
In the above examples, electron movements in Lewis structures needed to convert one resonance structure to another are shown 1. 15 by “curved arrows”; each curved arrow begins at an electron pair and ends exactly at the new location of the pair – finally, note that contributing structures must have exactly the same nuclear positions and bonding totals. In resonance structures, atoms do not change positions, only electrons. I. Arrow Formalism in Organic Chemistry i) Curved arrows. – used to show movement of a pair of electrons either to go from one resonance structure to another or for a chemical reaction. eg, ii) Fishhooks. – used to show movement of single electrons.
Used for reactions which form, change or destroy radicals. eg, iii) Straight arrows (regular or fishhook) – used to show reactions or equilibria eg, iv) Double headed arrows. – used specifically to indicate resonance . eg, 1. 16 J. Electron Orbitals and Bonding The molecular geometry of bonded atoms suggests that there must be some change in an atom’s atomic orbitals before it bonds with other atoms to form a molecule. – consider the C atom in methane, CH 4 – we must assume that, before bonding to H atoms, the one 2s and three 2p atomic orbitals are mixed and rearranged among themselves to give a new set of four equivalent “hybrid” atomic orbitals arranged at the tetrahedral angles of 109. 5??.
This reorganization of atomic orbitals, which occurs in most atoms prior to bonding, is termed hybridization. – the electron pairs in the hybrid orbitals may be bonding (as in methane) or both bonding and non bonding (as in NH 3 and H 2O). The hybridization of atomic orbitals in methane is one possibility. Other types of hybridization also occur, as an atom will undergo hybridization in a fashion that allows the most stable molecular structure to form. There are five major hybrid types of atomic orbitals involved in organic molecules. They are: But only the first three are needed for C, N and O compounds (nearly all of organic chemistry) 1. 17 1. sp hybrid atomic orbitals one s plus one p, to give two sp hybrid orbitals.
Each has the form – the two sp hybrid atomic orbitals are directed 180?? apart. The two unused atomic p orbitals retain their original shape and are at right angles to the two sp hybrid orbitals, ie, – each hybrid orbital can hold an electron pair, bonding or nonbonding. The p orbitals can be used to form pi-bonds with adjacent atoms. The resulting geometry of molecules containing a central sp-hybridized atom is linear. This geometry is that expected by VSEPR theory and hybridization theory. e. g. Cl???Be???Cl, H???C /C???H , O=C=O 1. 18 2. sp 2 hybrid atomic orbitals one s plus two p, to give three sp 2 hybrid orbitals there are three of these, 120?? apart.
Thus, three electron pairs can be accommodated in the hybrid orbitals, one p orbital is unaltered, and the resulting molecular geometry is planar triangular, e. g. CH 3+, H 2C=O 3. sp 3 hybrid atomic orbitals one s plus three p, to give four sp 3 hybrid orbitals. Already described above for methane. 1. 19 and for methane, CH 4 – with sp 3 hybrids, the electron pair arrangement is always tetrahedral, but the shape description (based on bonds) may differ if NB pairs are present, ie, 4 B pairs tetrahedral 3 B 1 NB pyramidal 2 B 2 NB bent – organic compounds containing C, N, and O as the central bonding elements usually contain only atoms hybridized in the sp, sp 2 and sp 3 forms – thus it is simple to relate hybridization and shape in organic compounds, ie, 1. 20 jjj
VSEPR # of Type electron pairs on central atom AX 2 AX 3 AX 4 2 3 4 electron pair geometry linear trigonal planar hybridization of central atom sp sp 2 tetrahedral sp 3 eg Q. Label the hybridization (and give the approximate bond angles) of each C, N and O atom in the following molecule. (build this molecule with your model set) K. Functional Groups The reactivity patterns of organic molecules are most simply understood by recognizing specific atomic groupings that confer characteristic reaction possibilities to all molecules of which they are a part. Such groupings are known as functional groups, and most of those covered in this course are tabulated below. it is important to understand that, no matter how large or complicated the organic molecule, the functional groups in that molecule will undergo the characteristic reactions listed. Therefore it is important that you can identify the functional groups present in a molecule. 1. 21 name alkane general formula R3C???H CnH2n+2 R2C=CR2 CnH2n RC/CR CnH2n-2 contain characteristic reaction combustion example alkene addition alkyne aromatic compounds addition substitution alcohols R???OH oxidation ethers R???O???R combustion amines R???NH2 basicity aldehydes oxidation ketones reduction carboxylic acids acidity esters amides hydrolysis hydrolysis [in all the functional groups listed above, R = any chain of carbon atoms, or an H atom] 1. 22 although alkanes are unreactive and convey no specific reactivity (except for combustibility) to compounds of which they are the structural backbone, they are usually included in functional group lists for completeness . But alkanes are not counted as functional groups in molecules. – C atoms are categorized as primary (1??), secondary (2??), tertiary (3??) or quaternary (4??) depending on the number of alkyl groups (C chains, R) bonded to the C atom, eg, a 1?? carbon is bonded to one other C, eg, R-CH 3 a 2?? carbon is bonded to two other Cs, eg, R-CH 2-R a 3?? carbon is bonded to three other Cs, eg, R 3-CH a 4?? carbon is bonded to four other Cs, eg, R 4-C – amines are also frequently categorized as primary (1??), secondary (2??), tertiary (3??) or quaternary (4??) depending on the number of alkyl groups (C chains, R) bonded to the N atom.
Thus 1?? amine has one R group, eg, CH 3NH 2 2?? amine has two R groups, eg, CH 3NHCH 2CH 3 3?? amine has three R groups, eg, (CH 3)3N 4?? amine has four R groups, eg, (CH 3)3NCH 2CH 3 – alcohols are categorized as primary (1??), secondary (2??) or tertiary (3??) depending on the type of C to which the OH is attached, ie, 1?? alcohol has an OH bonded to a 1?? C, eg, CH 3CH 2OH 2?? alcohol has an OH bonded to a 2?? C, eg, (CH 3)2CHOH 3?? alcohol has an OH bonded to a 3?? C,eg, (CH 3)3COH 1. 23 eg Q. Label all functional groups in the molecule capsaicin, responsible for the “hot” taste sensation in peppers and paprika. pgs covered: Chap. 1,1-36. Quick Quiz Ques: all text problems: 24,26,27,28,30,32,34,38,42,47,49,50,52,54, 56,58,61,62,63,66,68. Oct. test problems: 2007: 1-6; 2008: 1-6; 2009: 1-4; 2010: 1-5 Past Test Q: 1) In which of the following molecules is the atom marked with an asterisk (*) not sp 2-hybridized? 1. 2. 3. 4. 5. [A] 3 & 4 only [D] 2 & 4 only [B] 1 & 2 only [E] 3 & 5 only [C] 1 & 5 only 2) Which one of the following anions is not stabilized by resonance delocalization? A] [B] [C] 1. 24 [D] [E] MCAT Q: Refer to the structure of aspartame, an artificial sweetener, to answer questions 1 and 2. aspartame 1. The hybridization of the N8 nitrogen in aspartame is: A) sp B) sp 2 C) sp 3 D) sp 3d 2. The bond angle formed by H1, C2 and H3 in aspartame is: A) 180?? B) 120?? C) 109?? D) 90?? 3. A brightener that is used on wool and nylon fabrics is 7dimethylamino-4-methylcoumarin. What functional groups are present in this brightener? A) aromatic ring, alkene, ketone, ether and amide B) aromatic ring, cycloalkene, ketone, and ether C) aromatic ring, cycloalkene, ester and amine D) aromatic ring, cycloalkene, ester and amide