You can review content from Crash Course Organic Chemistry with the Crash Course App, available now for Android and iOS devices.

 Hi! I’m Deboki Chakravarti and welcome to Crash Course Organic Chemistry!

 Eating a balanced diet is important because there are some organic compounds that our bodies need, but can't make from scratch – like vitamins!

 So we get them (or precursors of them) in our diet. Let’s look at vitamin A, for example, which is found in meat, fish, eggs and dairy products. If you don't have enough vitamin A in your body, you can get all sorts of health problems including night blindness, infertility, and frequent infections. One solution to fight vitamin A deficiency around the world is a type of rice that has been genetically engineered to produce beta-carotene – a chemical precursor for vitamin A. Beta-carotene is also the orangey pigment in carrots, so this rice has an orangey color and has acquired the name Golden Rice. But why does beta-carotene look orange, and why don't other chemicals? The key is its conjugated electron system– and that's what we're exploring today.

 [Theme Music]

 In episode 36, we talked about benzene rings and conjugation: when a molecule has alternating single and double carbon bonds. Conjugated alkenes are stable because of all those overlapping pi bonds. One way to think about this is to look at the possible resonance structures. Resonance stabilizes molecules! We can prove this stability by looking at the thermodynamics, too. Alkenes can be converted into alkanes by hydrogenation, where we add hydrogen across the double bond. This is done with hydrogen gas and a catalyst– usually palladium on carbon, or platinum. If we hydrogenate an alkene with one double-bond, like but-1-ene, the enthalpy change is negative 127 kilo-joules per mole. But if we hydrogenate an alkene with two double bonds, conjugation can begin to play a part. Now, depending on the size of the molecule,two double bonds can be alternating, or not.

 So let’s start by looking at the hydrogenation of penta-1,4-diene, where the double bonds aren’t alternating. The enthalpy change is negative 252 kilo-joules per mole, which is pretty close to double what we got for but-1-ene. But if we hydrogenate 1,3-butadiene, with alternating double bonds, the measured enthalpy change is only negative 237 kilo-joules per mole. That's 15 less than for penta-1,4-diene!

This shows, experimentally, how conjugated double bonds of 1,3-butadiene stabilize the molecule. So we can see how conjugation stabilizes a molecule through resonance structures and measuring enthalpy. But we can also zoom in further to look at the physics of the covalent bonds. Remember, a covalent bond is when two atoms share a pair of electrons. Metaphorically, we can think of a covalent bond as a marriage. Like… a perfectly normal marriage between an all-powerful mutant witch and a sentient AI in a vibranium synthezoid body. Opposite charges attract, so there are attractive forces between the negatively charged electrons and the positively charged nuclei in the atoms. (These are the forces bringing our happy couple together.) But like charges repel each other, so electrons repel other electrons, and atomic nuclei repel other nuclei. (These are the forces, like the never-ending stresses of being a superhero, that can pull our couple apart.) Like any relationship, the forces that go into a covalent bond are carefully balanced for it to work!

 So let’s look at what happens when two lone p orbitals combine to make the molecular orbital of a carbon-carbon double bond. p orbitals are sort of hourglass-shaped. They’ve got a node in the middle, where no electrons can exist because there’s a change of phase. In organic chemistry we don’t need to worry about the nitty-gritty of what a phase is. It has to do with quantum mechanics. We just need to know that there are two phases, which are usually shown with shading or different colors in diagrams. When two p atomic orbitals overlap, we form two pi molecular orbitals. But there are two different types of pi molecular orbital, depending on how the phases of each p atomic orbital are oriented. When the lobes of the two p orbitals are the same phase, they overlap constructively and we form a bond between the two atoms. The attractive forces between the electrons and the nuclei outweigh the other repulsive forces.We call this a bonding orbital, and it results in a net lowering of energy. If the lobes of the two p orbitals are in different phases, the nuclei may be close together in space, but there aren't any attractive interactions and there's one of those zero-electron-density nodes between them.The repulsive forces win out over the attractive forces, and the result is even less stable than two non-bonded atoms. So we call this an antibonding orbital.

 Imagine the parents of the groom have been through an acrimonious divorce and refuse to talk to each other, but are forced to sit on the same table at dinner. You can practically feel the tension and it's higher in energy. Orbitals are a “one in, one out” kind of situation. One atomic orbital equals one molecular orbital, two atomic orbitals combining equals two molecular orbitals – like we just saw – and so on! Each orbital of any kind can hold two electrons, and the lowest-energy orbitals fill up first. So if we start taking our atomic orbital electrons and slotting them into molecular orbitals, we’ll usually run out before we fill the highest-energy orbitals. You can think of any marriage as a balance between good moments and bad ones – so in the case of an atomic bond, it's a balance between bonding orbitals and antibonding orbitals.

 The low-energy bonding orbitals fill up first, adding to the pile of good moments. But if too many electrons end up in high-energy antibonding orbitals, the bond or marriage becomes unstable with bad moments. Let's say we're looking at a conjugated molecule, like 1,3-butadiene. Each carbon atom has sp2 hybridization plus a p orbital. The basic principle stays the same: butadiene has 4 p atomic orbitals, and therefore 4 pi molecular orbitals. But all these molecular orbitals are a little different – and we can picture those differences by imagining a wedding cake!

 In the case of butadiene, with 4 pi molecular orbitals, we have 4 tiers. Each tier is a molecular orbital. The bottom tier of our cake is decorated in an ordered, consistent pattern. The designs – or p atomic orbital phases– are all aligned the same way. This forms a low-energy bonding molecular orbital, and this bottom tier is the most low-energy and stable. It has to be, because the rest of the cake is on top of it! The next tier is decorated with two designs– two up, two down. Two of the p atomic orbital phases are one way, and two are another way.

So our second tier is stable – another bonding molecular orbital – but it's a little bit less foundational than our bottom layer. The third tier is also decorated with two designs – two up, two down – but they’re more split up. So we're getting a little more high-energyand precarious here. This layer is an antibonding molecular orbital. Finally, the top tier of our cake is the most unstable layer, and the highest in energy. All the atomic orbitals are out of phase with each other, and this is also an antibonding molecular orbital. So this tasty representation of butadiene captures the basic idea of molecular orbital energy levels. But we can state them as non-cake-based rules,too.

 Each energy level adds an extra node to the molecular orbitals. Also, the orbitals are symmetrical – so where there’s a single node, it has to go right through the middle, but where there’s more than one, the nodes are spaced out evenly. Like we said earlier, all these molecular orbitals exist, but not all of them are full of electrons. And that "fullness" is a big factor in the stability of our wedding cake – or our molecule.

 To talk about which molecular orbitals are full or empty, we use the terms Highest Occupied Molecular Orbital, or HOMO, and the Lowest Unoccupied Molecular Orbital, or LUMO. 1,3-Butadiene has four p atomic orbitals. So these four electrons fill up the first two tiers of our cake, which are low-energy bonding orbitals. The HOMO is a bonding orbital and, as we’d expect, butadiene is nice and stable. Stepping out of metaphors and into the lab, conjugated molecules allow us to use a new imaging technique: ultraviolet spectroscopy. We’ve talked about spectroscopy before in this series. Looking at the electromagnetic spectrum, ultraviolet falls between about 200 and 400 nanometers – from the “blue” end of visible light to X-rays. When you throw electromagnetic energy at a molecule, it either absorbs that energy, or not. Absorptions are often interesting because they give us a clue to the structure of the molecule. So let’s look at what happens when a conjugated molecule (like butadiene) is exposed to UV light.

In the ground state, the two lowest-energy, bonding molecular orbitals of butadiene are occupied. But when butadiene is exposed to certain wavelengths of UV, that energy is absorbed and electrons jump up from lower energy occupied orbitals to higher energy ones. For example, a pi electron can jump from the HOMO to the LUMO. Because the electron goes from a bonding pi molecular orbital to an antibonding pi molecular orbital, this kind of jump is called a pi to pi star excitation. For butadiene, we know the exact wavelength of UV light that leads to pi to pi star excitation: 217 nanometers. But other conjugated systems also have characteristic UV absorptions. The wavelength changes depending on the gap between the HOMO and LUMO, which depends on the conjugated electron system. Specifically, the energy difference between the HOMO and LUMO decreases as the conjugation increases. The spectrum looks a little bit like an IR spectrum, but in IR spectroscopy, peaks at different wavelengths indicate different sorts of bonding arrangements. UV spectroscopy tells us specifically about molecules with conjugated pi systems – and the spectra are often only a single peak. The amount of UV light absorbed is the sample’s molar absorptivity. It's represented by the symbol epsilon, and this equation. Molar absorptivity is a characteristic of a particular pi electron system. Assuming we can measure or look up the other variables, we can use this equation to figure out the concentration of our sample really quick.

For example, let's go back to golden rice! Say we have a sample and we'd like to measure how much beta-carotene it has using UV spectroscopy, because beta-carotene is a conjugated molecule. So we mush up the rice, use a specific volume of organic solvent to extract the beta-carotene, and pop this liquid sample into our spectrophotometer. We can look up that beta-carotene has a molar absorptivity of 138,000. From our experiment, we know the path-length of the cell is 1.0 centimeters. And we measure the UV absorbance as 0.37. So, the concentration is… 2.7x10-6 molL. Pretty neat!

Now, we could end the episode here with UV spectroscopy. But beta-carotene doesn't just absorb UV light! Its conjugated electron system can also absorb light in the visible spectrum, which explains why it – and a handful of other organic chemicals – are colorful. As you can see with a prism or in the sky when sunlight hits atmospheric ice crystals just right: white light is made up of light wavelengths that span all the colors of the rainbow. When an organic molecule like beta-carotene looks orange to us, that means orange wavelengths of light are reaching our eyes. So blue light – the complimentary color – is being absorbed by the molecule. We’re effectively seeing white light with the blue and violet wavelengths taken out! In fact, our eyes rely on lots of chemical reactions that involve conjugated molecules.

Our eyes have two types of light-sensitive cells: cones and rods. Cones help us perceive colors, while rods help us see in dim light. In the photoreceptor cells of the eye, a component of vitamin A called 11-cis-retinal, along with a large protein molecule, is converted into a light-sensitive compound called rhodopsin – which is why vitamin A deficiency can cause night blindness! When light hits the rod cells, isomerization occurs at the double bond between the 11th and 12th carbons. This change from Z isomer to E isomer causes a nerve impulse to travel through the optic nerve to the brain where, amazingly, our brain translates it into vision!

In this episode, we've learned that: Conjugation de-localizes electrons and stabilizes molecules, p atomic orbitals overlap to form pi molecular orbitals of different energy levels, and UV spectroscopy can help us determine the concentration of conjugated molecules. Next time we’ll talk more about molecular orbitals when we discuss pericyclic reactions!

Thanks for watching this episode of Crash Course Organic Chemistry. If you want to help keep all Crash Course free for everybody, forever, you can join our community on Patreon.