118AMan1 - CHEM 118A Week 1.Structure and Bonding I Organic...

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Unformatted text preview: CHEM 118A Week 1 ....... ..Structure and Bonding I Organic molecules have very well defined structures. These structures are governed by the principles behind covalent and ionic bonding. Because most of the bonds in organic compounds are covalent, we will focus on this mode of attraction between atoms. We will see that the use of covalent bonds gives rise to molecules where atoms are attached in a definite order and to specific other atoms. Even though you may at some time have seen printed formulas of organic molecules that look like C12H22011, the "formula" of sucrose (common sugar), be aware that each of the 45 atoms in a molecule of sucrose has a very specific location and set of bonds with one or more other atoms of the molecule. It is absolutely essential to know the order that the atoms are attached in an organic molecule. Otherwise, you will be completely unable to solve any kind of problem that pertains to how the molecule behaves. The first session will cover covalent bonding and the common ways that molecules are drawn in order to depict the order in which their atoms are attached, called the atomic connectivity. Electronic Structure and Bonding We begin by reviewing covalent bonding, especially with respect to the atoms we will commonly encounter in organic molecules. You should learn the positions of about 20 atoms in the periodic table, of which about a dozen (in bold) are the most commonly encountered: ABBREVIATED PERIODIC TABLE Group: 1 2 3 4 5 6 7 8 H He Li Be B C N 0 F Ne Na Mg A1 Si P S Cl Ar K 7 Br I Know the number of electrons in the outer shell of each of these elements. Or you are likely to get your clock cleaned in this class. Next, bonding between atoms happens when they achieve filled outer shells by either gaining, losing, or sharing electrons with other atoms. The kind of bonding chosen by a pair of atoms is made on the basis of energy: the type of bonding that releases the most energy will be the most favorable and will be the kind that occurs. A. Ionic Bonding One atom loses electrons to another. They become ions and attract each other ‘ electrostatically. Example: lithium fluoride (LiF) Li + F —> Li+ F— Lithium loses an electron to make Li+, with a filled outer shell (like He). Fluorine picks up the electron to become F’, with a filled outer shell (like Ne). Why does this happen? Let's look at the energetics of the process (on the next page). Li —> Li+ + e— AB = +124 kcal/mole (energy input needed) F + e— —> F— AE = — 80 kcal/mole (energy is released) Sum: Li + F ——> Li+ F" AB = + 44kcal/mole Oops. A net energy input appears to be required to make this pair of changes occur. T hat's no good. BUT, we're not done yet. Once formed, the ions are oppositely charged, and they attract each other electrostatically. This attraction releases energy—a LOT of energy: Bringing Li+ and F” together (to the actual distance between them in crystalline LiF) releases electrostatic energy of attraction AE = —184 kcal/mole Final net energy change AE = —140 kcal/mole So, overall, the reaction of Li and F to give ionic LiF proceeds with release of 140 kcal/mole. It is a very exothermic (energetically favorable) process. What about reactions between other atoms? If we do the same energy analysis for the bonds between a large number of atoms, we find that ionic bonding is favorable only when we combine a very electropositive element (like Li) with a very electronegative element (like F). Looking at the periodic table above, we see that many of the elements important to organic chemistry are in the middle groups. They have intermediate electronegativity values. For them, the energetics of gaining or losing enough electrons to give ions with filled outer shells are too unfavorable to be offset by the energy of electrostatic attraction. They will not form ionic bonds. Instead, the bonds between them will involve sharing of electrons and will be covalent. B. Covalent Bonding Two atoms share electrons with one another. The shared electrons are mostly located in the space between the atoms and attract both atomic nuclei simultaneously, forming the covalent bond. Most of the time, the atoms achieve filled outer shells by means of electron sharing. Example: molecular hydrogen (H2) H- + H- —> H:H (Lewis structure) or H—H (straight line or "Kekulé" structure) Each H brings in its single electron (in its ls orbital). By sharing, each achieves the 2e— closed shell configuration of He. The two types of representations shown above Will be discussed shortly. Example: methane (CH4) (the simplest organic molecule) . H I? + 4 H- —' H:C_:H or H-C—H H H Each H brings in one electron and shares one of the four valence electrons that belong to the C. So all the atoms end up with closed shells: 2 e— for each H and 8 e— for the C. Further information on bOnding will follow. Let us look first at methods for depicting structures of organic molecules. Drawing Structures of Organic Molecules There are four common ways of drawing molecules to describe connectivity. You must be completely comfortable with all of them, because they are used interchangeably. A. Lewis structures This form of structure serves several purposes. In addition to describing onnectivity, it also tells you the location of all valence electrons, both the ones in bonds and the ones in )ne pairs. lfa molecular species has an overall charge, or is neutral but has separation of charge due to nusual electron distribution, the Lewis structure assigns the charges to specific atoms. The Lewis "ructure is not complete unless such charges are shown in the proper places. Example: l—propanol, a molecule with the formula C3 H80, has the Lewis structure H : : : z H H H H B. Straight-line (Kekulé) structures Straight—line structures are a simplified version of Lewis .tructures. Each pair of electrons in abond is replaced by a line. Lone pairs are usually still depicted by lots, but are occasionally shown by a line, or may be omitted altogether if they are not "doing something" Luch as participating in a subsequent chemical reaction. It is recommended that you include lone pairs in zll structures you draw for at least the first couple of weeks of the course. .When you are asked on an :xam to draw a structure, always assume that lone pairs (and charges, if any) must be shown for your Lnswer to be, complete. @1313: l—propanol has the Kekulé structure H H H heated. i i l C. Condensed formulas In a condensed formula, the order of attachment of the atoms is depicted or the most part along a single horizontal line without showing bonds. This kind of formula is often [86d in printed books, because it is compatible with single-line typesetting. However, it is the most likely o be read incorrectly by persons unfamiliar with the c0nventional patterns for interconverting condensed ormulas with the other, more descriptive ones. Example: l—propanol has the condensed formula CH3 CHZCHZOH. D. Bond-line (zig-zag) formulas In bond—line formulas, the framework of carbon atoms is eprescnted only by zig-zag line patterns. Each line end and each angle is understood to stand for a arbon atom and any hydrogen atoms attached to it: the C and H symbols are not written. Other atoms re still represented by their chemical symbols and are shown connected by lines to the appropriate arbon atom in the framework. Hydrogens on these other atoms are shown by the letter H. This is the ormula most frequently used by chemists, because it is the fastest to write. Example: l—propanol has the bond—line structure OH N n which the end of the zig—zag line (at the left) represents the CH3 group, and the two angles each epresent CH2 groups. i few more examples of molecular structures depicted in each of these four ways follow (next page). Butane, C4H10 HHHH a??? H: : z : : H H—IC—CII—C—C—H CH3CH2CH2CH3 N N —Methylpropanamide, C4H9NO . . H HH'O'HH HHiO'HH I n c 1—Chloro—3—methoxybutane, CSHI 1ClO I I H I I N H:g:c:C:I\;:§:H H—C—C—C—N—C—H' CH3CH2CONHCH3 /\”/--\ :O: H H I. I H: C : H H_C_H .- I II/ :0 H :0: H H H :(l): H H H:C:C:C:C:Cl: H—c—C—C—c4—(EII: CH3CH(OCH3)CH2CH2C1 M51; HHHH tlti Notice several patterns from these formulas. For example, the number of bonds an atom in groups 4 through 7 usually forms is related to its position in the periodic table by the relationship 8 — Group number = Number bonds So, halogens normally form 8 — 7 = 1 bond oxygen and the other elements in its group normally form 8 — 6 = 2 bonds nitrogen group elements normally form 8 — 5 = 3 bonds carbon groups elements normally form 8 — 4 = 4 bonds. When we say "normally" we refer to the situation where the element has zero charge. We will see many cases where these elements are charged, and in these cases the number of bonds will be different. For example, oxygen in water is neutral and has two bonds. But oxygen in hydronium ion is positively charged and has three bonds, while oxygen in hydroxide ion is negatively charged and has only one. H ..+ .. .. _ HzOzH H:O:H H:O: Hydronium ion Water Hydroxide ion Be aware that elements in the third row of the periodic table and below can "expand their octets" and form more bonds than second row elements. For our purposes, watch out for expansion of the octets in P and S. Elements in groups 1—3 normally form a number of bonds equal to their group number. Condensed formulas, the ones most likely to cause confusion, are typically constructed from left to The first non—hydrogen atom is followed immediately by all the atoms or groups immediately , I n to it (so you have to know how many bonds it makes!!!) Then each successive non—H atom is 5 as beng bonded both to the non—H atom to its left and to as many atoms to its right as are necessary it g‘ve it its proper number of bonds. Parentheses are used to eliminate ambiguity. Confused? Let's t it at the examples shown previously for clarification. We have labeled the non—H atoms for clarity. Butane: CH3CH2CH2CH3 a b c d Carbon (a) has four bonds, three going to the three hydrogens immediately following it in the formula, .1 the fourth to carbon (b). Carbon (b) has one bond to carbon (a), two to the hydrogens immediately filowing it in the formula, and the fourth to carbon (c). Carbon (0) is similar. Carbon (d) has one bond t) carbon (c), and the remaining three to the hydrogens to its right. CH3CH2CONHCH3 N—M thl 'd: e ypropanami e a b Ode f Carbons (a) and (b) are like those in butane. Beginning at carbon (c) is a grouping that can cause trouble. Carbon (0) has one bond with carbon (b). It needs three more. Students in cases like this often jump to the conclusion that the three bonds go to the next three atoms, but that leads to a nonsense answer, because then the O and the N do not get their proper numbers of bonds, and the place to attach carbon (f) is unclear. Another nonsense answer attaches all the non—hydrogen atoms in a row; that is, CH3—CH2—C—O—NH—CH3. This answer gives 0 and N the right number of bonds, but carbon (c) has only two, not four as it should. Let's look more closely at the problem. We need to add two more bonds to carbon (c), using only the atoms that are already there. We can get to the right answer by process of elimination. Carbons (a), (b), and (f) are OK as written in this structure. What can we change? We can assume that the N is attached to carbon (c) rather than 0. This removes a bond from the 0, but we can fix that by adding a second bond between carbon (c) and oxygen: making them doubly bonded. That's it! Look at the correct answer on the previous page! So, how do you tell if a condensed formula implies carbon double bonded to oxygen (or to any other atom)? Here's how: If a carbon in a condensed formula does not have four bonds by the time you get to the next non— hydrogen atom to its right, then it is possible the structure contains a multiple bond between the two. Each of the starred carbons in the structures below is an example of this situation: CH3COCH3 means CH3—|(|3—CH3 CH3COOH means CH3—!(lZ—-O—H (not CH3—C~O—CH3) (not CH3—C—OO—H) o O H CH3CHO means CH3—ICl—H CH3CONH2 means CH3—lCi—lll—H (not CH3—C—O—H) (not CH3—C—O—NH2) A multiple bond is not required in every case: look at the last example one page ago. There, the parenthesis tells you that the second carbon from the left is attached to both the O of the OCH3 group and the CH2 carbon next down the chain. It has four single bonds, and everybody is happy. ...
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This note was uploaded on 02/21/2009 for the course CHE 118A taught by Professor Patten during the Fall '08 term at UC Davis.

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118AMan1 - CHEM 118A Week 1.Structure and Bonding I Organic...

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