Lecture 9: Lewis Structures, Covalent Bonds, and Resonance

Lecture 9: Lewis Structures, Covalent Bonds, and Resonance Summary Covalent bonds are created via the sharing of electrons rather than through electrostatic attraction between cations and anions, as in ionic bonds. The shared electrons in covalent bonds achieve a lower and more stable energy state by interacting with two nuclei instead of one. Lewis Dot Structures represent the two electrons in a covalent bond with lines, as shown for H2 to the right. H-H H-Cl Cl-Cl bonds diagram removed due to copyright restrictions. To represent more complex molecules using Lewis Dot Structures, the following set of rules come in handy: 1. Connect atoms, central often less electronegative 2. Determine total number of valence electrons 3. Place bonding pair of electrons between adjacent atoms 4. Starting with terminal atoms, add electrons to each one to form octet (2 for H) 5. If electrons are left over, place on the central atom 6. If central atom hasn’t reached octet, use lone pairs from terminal atoms to form multiple bonds to the central atom to achieve octet It is possible to have more than one Lewis structure that doesn’t valuate the octet rule. For these instances, we calculate the formal charge. The sum of the formal charge must equal the overall charge on the molecule or ion. To calculate formal charge, follow these steps: 1. Nonbonding electrons are assigned to the atom on which they are located. 2. Bonding electrons are divided equally between the bonded atoms. 3. Then, the formal charge can be computed for each atom as follows:   bonding e−0 s f ormal charge = valence e− − nonbonding e−0 s + 2 In addition to stability, formal charge tells us about the structure’s electronegativity—atoms with negative formal charge should be more electronegative, and those with positive formal charge less electronegative. It’s also possible to have two different Lewis structures with the same formal charge. This indicates that the actual molecule is a combination of its most stable resonance structures, as with the ozone example in class, where we have on average 1.5 bonds on either side of the center oxygen. We also discussed the concept of polarity—a measure of how unequally electrons are shared between two atoms in a bond. In nonpolar covalent bonds, the electrons are shared equally, as in H2. 47 In polar covalent bonds, the sharing is unequal. Electronegativity (χ), the tendency of an atom to attract a shared pair of electrons toward itself, can be used as a measure of polarity in a bond. In a completely nonpolar bond, the two atoms will have identical electronegativities, meaning Δχ = 0. In a strong ionic bond, there is a large difference in electronegativities. In CsF, using Pauling’s scale, Δχ = 3.98 − .79 = 3.19. Finally, we named some ways that the octet rule can be broken. Some elements, like boron, can have fewer than 8 electrons. Others, like sulfur, can have an expanded octet. Why this matters Speaking of the CO2 molecule and its formal charges and stable structure, that’s the focus of this Why This Matters. Let’s get this done up front: after a century and a half of experiments and data collection, the scientific community has established with enormous consensus that carbon dioxide emissions are the primary cause of climate change. Other theories such as solar cycles or volcanic activity do not correlate with observed temperature changes, and there is now only one consistent explanation for rising global temperatures and our changing climate: anthropogenic greenhouse gases, led by carbon dioxide. By the way, if anyone (say, I don’t know, a politician?) were to tell you something like, “There is tremendous disagreement as to whether the earth is round. The debate over the earth’s shape should be encouraged — in classrooms, public forums, and the halls of Congress,” you would call such a remark either obliviousness or a lie. This is the same kind of lie told by many people about the facts of climate change. A great deal of media coverage focuses on how to define such “alternative facts,” but categorizing the remark is beside the point. Regardless of what we call it, this type of statement is a threat to the very foundations of science, to reason-based argument, and therefore to society. Politicians may live in a “post-fact” world, but scientists do not, and cannot. There is no “post-fact science.” Ok, back to our molecule. Why is this particular molecule so important for climate change and global warming? CO2 is called a “greenhouse gas” because of its ability to absorb infrared radiation, which in turn warms the planet. Take a look back at the plot of intensity of solar radiation hitting the planet as a function of frequency, from Why This Matters in lecture 4. Yellow is what strikes the upper atmosphere and red is what hits earth’s surface, so the difference between the yellow and red curves is where some of that solar energy is being absorbed. Small molecules in the atmosphere, like ozone, water, and carbon dioxide, are what absorb that incoming solar energy. You can see where CO2 is absorbing, out towards the right in the infrared part of the spectrum, while ozone is over on the left in the UV, and water is everywhere. It’s the strength of IR absorption that makes CO2 such an effective greenhouse gas. About 25% of the solar radiation that hits the earth (so 25% of the integrated red in the plot) gets reflected back out to the universe. But what is reflected is not what is absorbed: instead, most of the energy that is reflected is in the longer wavelength, infrared part of the spectrum. Since CO2 absorbs efficiently in that part, it acts like a blanket, keeping some of that reflected energy from escaping and redirecting it back to the atmosphere and surface. This blanket keeps us at a cozy 15◦ C instead of -18◦ C, which is what the average temperature on earth would be if all of that reflected solar energy were able to escape. As an aside, the term “greenhouse gas” is a bit of a confusing name that has unfortunately stuck. The effect of absorbing long wavelength radiation on the planet’s temperature has nothing to do with the effects of warming in a greenhouse used to grow plants, the latter relying mostly on preventing air convection from carrying heat away. Here’s the connection back to our Lewis structures and formal charge. That efficient absorption of CO2 in the infrared I mentioned, that is directly related to how the atoms in the molecule vibrate. The photon is absorbed because it gets transferred into vibrational energy, so the specific way atoms jiggle around is extremely important to what photons it can absorb and how efficiently it does so. If the CO2 molecule were to have a triple bond and a single bond, which is a valid Lewis structure but with higher formal charge, then its patterns of vibrations would change dramatically compared to the more stable double-bond case (with formal charge=0), which in turn would make it much less efficient at absorbing all the infrared energy! Why this employs Today’s lecture has been about the covalent bond, a convenient way to represent it with Lewis structures, and what we can learn about molecules from these structures. It’s hard to come across a job in any engineering or science field that doesn’t require at least some understanding of the covalent bond since it’s so ubiquitous. This also makes it hard to zero in on a specific job market for Why This Employs. That’s why I’m going to focus way down to a single material – carbon – and the jobs you could find related to its covalent bonds. In a few lectures we’ll be learning about molecular orbitals and as an example you’ll see just how flexible carbon can be in its bonding environment. It’s one of the most flexible elements in the whole PT in that it likes single, double, and triple bonds as well as anything in between (read: resonance structures and delocalization). Let’s pick one of the many forms carbon takes to focus on here: the aromatic ring. The term aromatic does have its origins in the fact that aromatic molecules can have a sweet fragrance to them, but in fact benzene—the classic aromatic molecule— has no such smell. Instead, the way the term is used in chemistry is that the electrons in a ring of carbon have lowered their energy by delocalizing (yes, resonance structures!). Most often this occurs for a ring that has alternating single and double bonds, which then form an average over its resonance structures. Benzene is the classic example of this, as it is the simplest aromatic ring one can make. And you’ve probably heard of benzene before, but what you may not know is just how many ways benzene can be modified to make or do something useful. Check out this chart of only a very small fraction of benzene derivatives, where the ring has some addition(s) to give it a specific property or functionality. This massive variation of the chemistry of a single molecule is possible largely because of that covalent bond that carbon makes. 49 Figure of benzene derivatives by Compound Interest 2014 - www.compoundchem.com. License: CC BY-NC-ND. This content is excluded from our Creative Commons license. For more information, see https://ocw.mit.edu/fairuse. Many materials we encounter on a daily basis are manufactured from derivatives of benzene. Polystyrene, for example, is a polymer formed by connecting styrene molecules together, and billions of kilograms of polystyrene are produced each year. So, certainly jobs in companies that involve manufacturing, recycling, or utilizing plastics like polystyrene are related (I’m thinking a senior chemist position at BASF, a chemical manufacturing company, for example). But the uses of benzene derivatives go far beyond plastics. For example, industries use benzene to make resins, nylon and many different synthetic fibers, as well as lubricants, rubbers, dyes, detergents, drugs, and pes3-D image of acetaminophen molecule ticides. Take medicine as another example: here’s the removed due to copyright restrictions. acetaminophen (Tylenol) molecule – it’s a benzene derivative! So the pharmaceuticals industry is also a great place to take your knowledge of benzene (I’m thinking Johnson and Johnson, Pfizer, Novartis, and so on). These companies invest billions and billions of dollars every year to fund research on the development and testing of new medicines, many of which have at their core a resonant Lewis structure. Example Problems 1. Draw the resonance for CHO2 − and assign formal charges to each atom. 2. True or false: The following picture shows two resonance structures of a molecule Solutions 1. 2. False. The two structures aren’t resonant structures because the atoms move around.