1 1.1 1. Lewis Structures. 1.1 What is Organic Chemistry? Why Study It? Organic chemistry is the study of the chemical behavior of compounds containing the element C. C is unique among all the elements for its ability to form strong bonds to itself (catenate). Thus, by bonding to itself, C can form a multitude of chemically and thermally stable chain, ring, and branched compounds. This ability of C to catenate is what makes it possible in for millions of carbon-containing compounds to exist. In this course we'll be studying the structure, reactivity, and synthesis of organic compounds. For example, we'll look at the following reaction. We'll ask, what are the three-dimensional shapes of the starting material and the product? What parts of the starting material are likely to be transformed under given conditions? What are the unstable intermediates in this reaction, if any? How many different products can be obtained? And, what is the nature of the bonds that are broken and the bonds that are formed? 1.2 Electronic Structure of Atoms and Molecules. Valence Structures. - every atom has a characteristic number of electrons, given by its number in the periodic table. H has 1, C has 6, N has 7, O has 8, Cl has 17 electrons. Every electron in an atom resides in a region of space called an orbital, which has its own characteristic energy. Orbitals are divided into groups called shells. It takes two electrons to fill shell 1, which has only one s orbital. It takes eight electrons to fill shell 2, which has one s and three p orbitals. The order of energy of orbitals is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s .... (The 3d orbital needs not to be filled for shell 3 to be considered filled.)2 Some atomic orbitals Following the aufbau principle, the electrons in an atom go into these orbitals in order of increasing energy. So the following atoms have the electronic structures shown – Figure 1.12. The electrons with which we concern ourselves are the outer-shell, or valence, electrons. Valence electrons are the electrons in the outermost shell. For neutral atoms the number of valence electrons can be obtained by looking at the element's column in the periodic table: H has 1, C has 4, N has 5, O and S have 6, F and the other halogens have 7. F– or Ne has no or eight valence electrons, depending on how you count it, because their outer shell is filled. 1.3 Lewis Structures Atoms link together to form molecules because they "like" to achieve a filled shell configuration. They achieve this by giving up or taking electrons (ionic bonds) or by sharing electrons (covalent bonds). In organic chemistry, we can have molecules that consist of as few as two atoms (CO) or those that consist of millions of atoms (DNA). The tremendous diversity of structures is what makes life possible. In organic compounds, most bonds are covalent and the compounds are usually written as Lewis Structures Rules to construct the electronic structure of organic molecules. Consider amide ion NH2–, ammonium ion NH4+, methanol CH3OH, hydrogen cyanide HCN, and methyl acetate, CH3CO2CH3. (1) Count the total number of valence electrons. Add one for every negative charge and subtract one for every positive charge. So, e.g., NH2– has 83 valence electrons (5 + 2x1 + 1) and NH4+ has 8 also (5 + 4x1 – 1). Methanol has 14, HCN has 10, and methyl acetate has 30 valence electrons (3x4 + 6x1 + 2x6). (2) Attach the atoms in the correct order using single bonds, being careful to give no atom more than four bonds (duet for H). Each bond uses two valence electrons. There are often many ways of combining the same atoms (this is called isomerism), and so you must have more information than just the empirical formula to do this. This may be supplied by an extended formula or by other information that we will learn to interpret later in the course. For NH2– and NH4+, there is only one way to attach the atoms to each other (Draw). These use up 4 electrons and 8 electrons, respectively, leaving 4 and 0 unused electrons, respectively. For methanol, the way the formula is written suggests that there are three H's attached to C and one attached to O. For HCN, there are two ways to arrange the atoms (H–C–N or C–N–H), but we can tell by the way the formula is written that the atoms should be attached H–C–N. We use up four electrons, leaving six. For methyl acetate, there are several ways to arrange the atoms, but we can tell by the way the formula is written that the atoms should be attached as shown. (The designation CH3, the methyl group, indicates C bound to three H atoms and one other atom.) This uses up 20 valence electrons, leaving 10 more. When you are constructing the single bond framework of an organic compound, it is useful to remember that H and the halogen atoms are normally found on the periphery of an organic structure, and not in the middle, because they have only one free valence. The CH3 group has only one valence, too. (3) Place the extra valence electrons on the heteroatoms (atoms that are different from C and H) or between pairs of neighboring electron-deficient4 atoms, being careful to give no more than a total of 8 electrons to each atom. (Each bond to an atom counts as two electrons.) Usually C makes a total of four bonds, N three, O two, and H and the halogens one. For NH4+, no extra electrons need to be placed. For NH2–, four electrons go on N. For methanol, the four electrons go on O. For HCN, the six electrons go on N. For methyl acetate, the five pairs of electrons can be distributed among the two O atoms to give the structure shown. (Alternatively, two pairs of electrons may be placed on each O and one pair between the electron-deficient C and upper O.) (4) Are any atoms electron-deficient by two or more electrons? If so, take a lone pair of electrons on a next-door neighbor and turn it into a bond between the electron-poor atom and its neighbor. If you have a choice of neighbors from which to borrow a pair of
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