Week 2 Aromaticity Workbook PDF - Monash S2 2024

7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 Week 2: Aromaticity - workbook Site: Monash Moodle1 Printed by: Kaltham Alzaabi...

7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 Week 2: Aromaticity - workbook Site: Monash Moodle1 Printed by: Kaltham Alzaabi Unit: CHM1022 - Chemistry II - S2 2024 Date: Sunday, 28 July 2024, 2:02 PM Book: Week 2: Aromaticity - workbook https://learning.monash.edu/mod/book/tool/print/index.php?id=2779874 1/11 7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 Table of contents 1. Pre-workshop material 1.1. Benzene 1.2. Drawing benzene 1.3. Resonance structures 2. Summary 3. Preparation quiz 4. Online lectures https://learning.monash.edu/mod/book/tool/print/index.php?id=2779874 2/11 7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 1. Pre-workshop material History of benzene Michael Faraday first isolated the hydrocarbon that we refer to as benzene (Figure 1) in 1825 from coal tar distillation. The structure of benzene was subject to considerable controversy. The formula of benzene was soon established to be C6H6, but the arrangement of these atoms was widely debated. Figure 1: Benzene Please read Section 16.7 of Chemistry, Blackman et al. (4th ed.) Please read Chemistry: Atoms First (2019), Chapter 21.1 https://learning.monash.edu/mod/book/tool/print/index.php?id=2779874 3/11 7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 1.1. Benzene The C6H6 formula suggested a structure containing multiple alkenes (a polyene), however, the behaviour of benzene was remarkably different to other polyenes. Various structures of benzene were suggested, but none accounted for the observed unreactivity of benzene. It wasn’t until 1872 that the first sensible structure of benzene was postulated by August Kekulé after a dream where a snake had “seized hold of its own tail…” This dream inspired his cyclic arrangement with alternating single and double bonds. Kekulé proposed that the double bonds in benzene rapidly equilibrates structures shown (Figure 2) giving the average between the two structures. Figure 2: Kekulé’s structure of benzene Evidence Against the Kekule Structure of Benzene We know now that benzene does not equilibrate between the two Kekulé structures. In the 1900s, diffraction methods (analytical techniques that can determine structure) confirmed the equivalence of the carbon-carbon bonds. The π electrons in benzene are equally spread over the carbon atoms in the molecule. These π electrons are delocalised over all six of the atoms in the ring meaning they are not localised in specific double bonds between two precise carbon atoms. http://molview.org/?cid=241 Figure 3: You can rotate and zoom around the structure by using your mouse with the 3D image. The delocalisation of the π electrons in benzene is supported by experimental evidence and theoretical calculations. Electron diffraction studies conducted on benzene show it to be a planar hexagon with identical carbon-carbon bond lengths of 140 pm (which is between that of a carbon-carbon single bond (154.1 pm) and a carbon-carbon double bond (133.7 pm)). Thermochemical evidence suggested benzene should not be described as a triene (hydrocarbon with 3 carbon-carbon double bonds) with localised single and double carbon-carbon bonds. https://learning.monash.edu/mod/book/tool/print/index.php?id=2779874 4/11 7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 1.2. Drawing benzene Drawing the Structure of Benzene We have various ways of representing the structure of benzene, however, it is important to point out the limitations of our representations. Benzene is a hybrid of the two Kekulé structures. The Kekulé structures are incorrect by themselves; the true structure of benzene is somewhere in between the two, as shown in Figure 4. Each makes an equal contribution to the hybrid and thus the carbon-carbon bonds are neither double nor single, but something in between. These two contributing structures as called resonance structures. Figure 4: Benzene as a hybrid of two equivalent contributing resonance structures Some chemists use a circle to represent the equivalence of the carbon-carbon bonds but how many electrons does a circle represent? Although the structures we use to represent benzene are not entirely accurate they are the best representations we have, and all are acceptable (Figure 5). Figure 5: The representations of benzene https://learning.monash.edu/mod/book/tool/print/index.php?id=2779874 5/11 7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 1.3. Resonance structures While Lewis structures are a useful way for chemists to represent and communicate the structures of molecules, these representations are often limited. The distribution of delocalised electrons within a molecule can be better represented by drawing a number of contributing resonance structures, with the overall electron distribution being a resonance hybrid, or combination of each contributing structure. Resonance structures are depicted with a special double-headed arrow (Figure 6). Figure 6: The resonance arrow Drawing resonance structures The Lewis structure of acetonitrile is shown in Figure 7. The C-N triple bond is polarised towards the more electronegative nitrogen atom, resulting in a partial positive change on carbon and partial negative change on nitrogen. This is not well described by the simple Lewis structure below. Figure 7: Lewis structure of acetonitrile The polarisation of the C-N triple bond in acetonitrile can be described by two resonance structures (Figure 8). The alternative resonance structure of acetonitrile can be drawn by moving a pair of electrons (blue arrow) from the C-N pi-bond to the nitrogen atom. Figure 8: Resonance structures of acetonitrile In the case of the acetate anion (Figure 9), the Lewis structure shows the negative charge located on one oxygen, while the other oxygen has a CO double bond. In reality, we know that both oxygen atoms have a partial negative charge, and that both C-O bonds are the same length. This can be shown with two resonance structures, or with a single resonance hybrid (right). Because the negative charge is distributed across two oxygen atoms, the acetate anion is more stable that the ethoxide anion, where the charge is located on one oxygen atom. This stability of the acetate anion explains why acetic acid is much more acidic than ethanol (more on this in week 6). https://learning.monash.edu/mod/book/tool/print/index.php?id=2779874 6/11 7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 Figure 9: Resonance structures of acetate anion In some cases, more than two resonance structures can be drawn for a molecule. The negative charge in the hydrogen sulfate anion (Figure 10) is distributed across three different oxygen atoms. This means that each oxygen atom carries only 1/3 of a negative charge, making it especially stable. As a general rule, the more resonance structures that can be drawn for an ion, the more stable that ion is likely to be. This explains why sulfuric acid (H2SO4) is a very strong acid. Figure 10: Resonance structures of hydrogen sulfate https://learning.monash.edu/mod/book/tool/print/index.php?id=2779874 7/11 7/28/24, 2:04 PM Week 2: Aromaticity - workbook | MonashELMS1 2. Summary This week, we have learnt about aromatic hydrocarbons, and their important place in organic chemistry. We have covered the discovery of and the electronic structure of benzene, and how to draw its structure. We have started to discuss resonance structures and their importance. https://learning.monash.edu/mod/book/tool/print/index.php?id=2779874 8/11