Cyalume reaction mechanism

Experiment 10 Spring 2020 Experiment 10 — Synthesis of Cyalume and Chemiluminescence _____________________________________________________________________________ Pre-lab preparation: Read the Wikipedia article on "glow sticks". If necessary, review recrystallization, i.e., (a) if you do remember, fine; (b) if you don't remember, then review it! (1) Briefly state the purpose of this experiment. (2) You may remember light sticks as one of those magical wonders of childhood. Ahhh... It's hard to imagine a better toy for a small child than a pressurized plastic tube filled with chemicals and broken glass. (a) Commercial light sticks contain a solution that surrounds a fragile glass tube containing a different solution. What are the key components of each solution? (b) What happens chemically when the glass is broken and the two solutions mix? (c) Why does pressure increase (just a bit) inside the tube? (3) (a) Put the following colors of light in order from lowest energy to highest energy: yellow, UV, orange, red, blue, IR. (b) Label one end of your series "lowest E" and the other end "highest E". (c) Label one end of your series "longest l" and the other end "shortest l". (4) Look at your TLC analysis of the Wittig product (Exp 9) — what color did you record for its fluorescence? (5) Write the mechanism of the reaction of the first equivalent of trichlorophenol with Et3N and oxalyl chloride. (6) Draw the structures of toluene and diethyl phthalate. Write down any physical data for these two solvents that are relevant to this experiment. (7) Why is it important to handle the reagents with gloves and keep them in your fume hood? (8) (a) Calculate the quantities of trichlorophenol, oxalyl chloride solution, and Et3N that you plan to use to make the cyalume. Keep in mind that we normally measure solids by mass and liquids by volume. (b) Calculate the theoretical yield of the cyalume product. As always, write concise, easy-to-read procedural notes for both parts, A and B, that you will use in lab. Reminder: These are notes on what you plan to do — These notes are not to be copied into your notebook! In your notebook, as you work, you'll of course write down what you actually did, along with relevant data and observations. There's a lot in this write-up — it provides an overview of excited states, absorption and emission of light, then explains how fireflies glow, in terms of reaction energetics, before moving on to the specifics of your reaction. Allow enough time to read this carefully before lab! 1 Experiment 10 Spring 2020 Light and color. Many molecules absorb light in the ultraviolet (UV) and visible regions of the electromagnetic spectrum. Molecules that absorb in the visible region appear colored under white light. When we look at a material under white light, the absorbing substance "removes" certain wavelengths. The color we see is determined by the wavelengths that are not absorbed. The unabsorbed light is reflected by opaque materials or passes through transparent things like solutions or colored glass. It's important to note that light emission is not involved here — in fact, most molecules do not emit light. The colors we see come from the wavelengths of white light shining on the substance that are not absorbed. A convenient approximation assigns each of the six basic colors to a 50-nm range of wavelengths (above). For example, purple light extends from about 400 to 450 nm, blue from about 450 to 500 nm, etc. That's easy to remember. (What most people perceive as yellow actually extends over a band of about 30 nm, and the transition from orange to red is around 620 or 630 nm, not 650 — that's fine-tuning.) What we see is the complement of the color absorbed (or the color in the center of a broad absorption band). Complementary pairs are purple–yellow, blue-orange, and green–red. So if a compound absorbs orange light, it will appear blue, etc. Light absorption, MOs, and electronic states. Absorption of light by a molecule causes an electron to "jump" from a lower energy molecular orbital (MO) to a higher energy empty MO. For organic molecules that means the electron must go from a bonding MO or from a lone pair orbital into an antibonding MO. There are lots of filled MOs and lots of empty MOs, so several electronic transitions can result from absorption of UV and visible light. The lowest energy transition — the one that requires the longest-l light — is the one from the highest occupied to the lowest unoccupied MO, or HOMO to LUMO (red in the diagram). This creates an excited state of the molecule that has 1 e– in each of these MOs, and that we refer to as S1. The other electronic transitions shown 2 Experiment 10 Spring 2020 require higher-E, shorter-l light, and these create higher-energy excited etc states, S2, S3, S4, etc. (The lowest-E electronic state — the ground state — is called S0.) LUMO In molecules, those higher-energy excited states quickly drop down to S1 without emitting light; the "excess energy" is lost as heat. E In molecules that emit light, the emission almost always comes from HOMO that lowest excited state, S1, falling to S0. So, while a molecule might absorb light at several different wavelengths (S0 —> S1 or —> S2 or — > S3, etc), it will only have one emission (S1 —> S0). etc What can excited states do? The S1 excited state of an organic molecule usually lives for only about 1 ns (10–9 sec). In some molecular orbitals (MOs) molecules this state can fluoresce — drop down to S0 by emitting a photon of light (remember, S1 —> S0 happens when the e– in the LUMO falls back into the HOMO). But in the grand scheme of things, fluorescence is not very common. Most molecules find other ways to lose that "excess" energy. In most cases when the excited state (S1) returns to the ground state (S0) it just turns all that energy into heat. What a waste. Some molecules can instead use the high energy of the S1 state to do chemistry that would have been impossible in the ground state. That's called photochemistry. Photochemical reactions often involve homolytic bond breaking to create free radicals, as well as a huge variety of isomerizations and rearrangements. Promotion of an e– by light can also lead to electron transfers. After light absorption, the high-energy e– can "hop" over a short distance to another molecule, and/or an e– from a nearby molecule can "hop" into the empty space (the "hole") that the excited e– left behind. Light-initiated e–-transfers are the basis for photosynthesis. You've probably also heard of phosphorescence, in which light is emitted over a much longer timescale — ms to sec or even longer (light emission in "glow-in-the-dark" stuff). Very few organic molecules phosphoresce.