Exceeding Nature

Original

The scheme belongs to a recent chemistry paper entitled "Nonmetal Catalyzed Hydrogenation Of Carbonyl Compounds" which I think shows significant advancement in chemistry. For the non-chemist, I'll unpack the title.

You may not be interested in hydrogenation, but hydrogenation is interested in you: it feeds you. The metal-catalyzed hydrogenation of vegetable oils is big business. You may have gotten away from trans-fats, but are you free of cis-fats?  How about saturated fats? The food industry uses hydrogen and metals like nickel to hydrogenate food stuffs. And then there is the "hydrogenation" of nitrogen to make fertilizer.

What these guys in London did is remarkable because they used hydrogen (H₂) to make alcohols (top right) from ketones (top left). And they used only C, H, O, B, and F atoms, spatially arranged as shown. No metals.

Nature has little use for H₂, the simplest of molecules. Relatively little free H₂ exists on earth. There is a class of enzymes called hydrogenases, but guess what? They use metals to activate H₂. So this work goes above and beyond Nature itself.

More notes on "The Disappearing Spoon"

[continued from previous post]

Part I   "Orientation: Column By Column, Row by Row"

1. Geography Is Destiny: H, He, B, Be, Sb
__________________________

Page 21, bottom of page:

Egyptian women were applying a different form of antimony as mascara, both to decorate their faces and to give themselves witchlike powers to cast the evil eye on enemies.

They used stibnite in which you can still see the Latin origin of antimony's chemical symbol, Sb. Stibnite gave the blueish black look which is still alluring, though antimony has been removed from reformulated modern eyeliner. The alchemist's symbol for antimony is:

which sort of resembles an upside down version of the female symbol.
________________________

Kean writes at length about Gilbert N. Lewis, as have I. My take on him is here and here.
_______________________

Now on to some substantive descriptive chemistry: Page 24, bottom:

As we move horizontally across the periodic table, each element has one more electron than it neighbor to the left. Sodium, element eleven, normally has eleven electrons; magnesium, element 12, has twelve electrons; and so on. As elements swell in size, they not only sort electrons into energy levels, they also store those electrons in different shaped bunks, called shells.

Early German quantum mechanics called this Aufbau or building up. Kean describes how electrons build shells -- s, p, d, and f orbitals -- in a logical way. His descriptions of p-orbitals as a "misshapen lung" and "d-orbitals" as balloon animals is amusing, but I would explain it differently. They more resemble blobs with 0, 1, 2, and 3 nodes as described here.

What Aufbau builds on is how electrons self-organize around an increasingly charged nucleus in moving from hydrogen to higher and higher elements. Start with the simplest atom having one proton and one electron. The very first electron goes into a spherical shaped 1s orbital surrounding the proton. Now if we add another proton to that picture to get to the next element (helium), we must add a second electron. It too goes into the same 1s-orbital and two electrons are happy as clams--perfectly-paired. The pairing of electrons is one of the most sublime aspects of electronic theory and is one which I struggle to understand.

Now move on to element 3, lithium: the third electron cannot occupy the same orbital space as its first two, so it must go into a higher energy orbital, the so-call 2s orbital. The 2s orbital is not exactly just a larger s-orbital; it actually interleaves with the 1s orbital as I drew attention to here:

The fourth electron in element 4, beryllium, perfectly pairs with the third one and fills the 2s orbital. Now, the fifth electron in boron could go into what's called a 3s orbital depicted above, i.e., electrons could just keep building higher and higher energy shells of spherical symmetry, but something else happens. A different type of node appears which breaks the spherical symmetry, creating what's called a p-orbital:

The 2p-orbitals are lower in energy than the 3s orbitals and that's why the next 6 electrons fill those first. There are 6 spaces because the electrons pair and go into 3 different p-orbitals -- one for each Cartesian dimension, x, y, and z.

[more soon]

Gallium Arsenide is Germane to Solar Cells

Gallium arsenide, a simple combination of two elements, interconverts light and electricity; GaAs lasers turn electricity into light and GaAs solar panels convert light back into electricity. There are alternative combinations of elements for these tasks, but each has its limits. What strikes me is how gallium and arsenic bookend germanium:

I need a name for "binary combination of elements which brackets and mimics another element." The term isoelectronic is close but doesn't cut it for me. There is a mathematical symmetry about GaAs in view of Ge and it goes like this: (31 + 33)/2 = 32 or, in chemical logic symbols: (Ga + As)/2 = Ge.

Like gallium arsenide, germanium is a photovoltaic material. Google "germanium solar cell" and you will find cutting edge research involving blends of gallium arsenide with germanium. I'm glad there is on-going research into new materials because I am not sure we should be putting arsenic on every rooftop much like we're putting mercury in every lightbulb.

A similar "bookend relation" occurs a couple rows up in the Table between boron, carbon, and nitrogen. Look how boron and nitrogen bracket carbon:

Once again, (5 + 7)/2 = 6. And just like carbon, boron nitride (BN) has both graphite- and diamond-like structures. One type of BN is even harder than diamonds: link

I see a pattern here: the centrality of the carbon group, C, Si, Ge, etc. to the family of main group elements:

A Girl's Best Friend is the Blue Diamond*

The structure of graphene got me to thinking of that other form of carbon, viz., diamond:

3D movie: link

Why is diamond so tough, so adamant, so opposed to physical change? I think the answer is called "perfect covalency" but not in an electron sharing sense:

The paradox of the diamond is interesting. Its atoms are not arranged in a tight, closest-packing order. They lack the triangulation of sound architecture. In order for its remarkable rigidity to be understood, I assume that the electrons which surround this meager structure supply it with its resistance to deformation. This is one reason why I cannot assume that electron clouds can infiltrate one another like vapors or ghosts. link

Pure diamond is also colorless and transparent, so what gives fancy (colored) diamonds their colors? The answer is not simply: "there must be something blue inside." An impurity is involved, but not the usual colored metal atoms like iron or chromium found inside other gemstones. The impurities in blue and yellow diamonds are carbon's left- and right-hand periodic neighbors--boron and nitrogen--playing little tricks on the lattice electrons.

Take the perfect 3D lattice of carbon atoms pictured above. Now suppose that we could randomly go in and replace every millionth carbon atom with a boron atom without perturbing anything else. What we get is a boron-doped diamond lattice. Because boron has one less electron than carbon, the entire lattice structure of the diamond is riddled with electronic "holes."

Now it just so happens that reddish-orange light has just the right energy match to promote an electron on an adjacent carbon atom into a "hole" next door. That jump in turn creates a new "hole" and so the next neighbor carbon jumps at the chance to fill the new hole and so on and so forth throughout the entire diamond lattice. Really, a blue diamond is rather like a doped silicon p-type semi conductor. In fact, blue diamonds are semi-conductors--albeit rather expensive ones!

Because only reddish-orange light is absorbed, the remaining visible light appears bluish to our eye because the white light lacks its reddish-orange component: remember the color wheel and complimentary colors!

Likewise, water in a white bathtub appears bluish because it absorbs some of the reddish component of the incident white light. That's also why heat lamps are red too--they are more or less tuned to the wavelength (infrared) that water in food absorbs and converts to heat. Microwave ovens are even better at this.

So what makes yellow diamonds? The answer is slightly more complex. If nitrogen atoms, carbon's other nearest neighbor, are doped into the diamond lattice instead of boron, each nitrogen brings an extra electron into the lattice which is easily promoted to the existing conduction band of diamond by violet light--ergo yellow appearing diamonds. Yellow diamond is analogous to an n-type semiconductor. Here's a link explaining in more detail why blue diamonds are blue and why yellow diamonds are yellow: Link.

Meanwhile, here's a very pretty picture of the Hope Diamond:

____________

*I don't mean "blue diamond-shaped" Viagra either.

Chemistry Is Like Sex: Coupling Illustrated

G.N. Lewis (I wrote about him back here) was from Berkeley and thus a bit more liberal when defining acids and bases: He gave more general definitions of them than Brønsted and Lowry did. And while more open-minded, Lewis was a bit of a chauvinist when he argued that a base's precious electrons helped "complete the octet" of an acid when they coupled.

Consider the coupling of a simple base, ammonia, with a simple Lewis acid, borane. If you're a jaded chemist who has "seen it all" you might consider just skipping to this link dealing with borane and ammonia making borazane as a hydrogen fuel energy source.

We already "know" what ammonia looks like here--but what about borane?  I wrote a bit about boron the element back here.  Turns out that the word "boron" is etymologically linked to the Arabic tongue as well, via the word borax.

Borane, BH₃, is a natural fit for ammonia's lone pair.  Consider its structure:

BH₃ looks a bit like NH₃ but completely lacks a lone pair.  BH₃ has only six surrounding electrons instead of eight and so is electronically unfulfilled. In a sense, it has a big hole in its middle. In the absence of an available lone pair, BH₃ readily dimerizes in a head-to-tail fashion with another sister BH₃ molecule to form B₂H₆. Here's an illustration of two BH₃'s getting it on together:

When NH₃ and BH₃ prepare to bond, a natural question is where should NH₃ put its lone pair?  BH₃ has what's called a "virtual orbital" (there's nothing virtuous about it).  A virtual orbital is just an empty electron orbital. Another name is a LUMO. Empty orbitals have metes and bounds, despite there being nothing there there. Here's a lurid depiction of borane's virtual orbital:

It's a bit hard to see in the depiction above but all three of borane's tripodal hydrogen limbs are squished flat into a planar configuration between the two swollen globes. The red and blue empty lobes are equivalent in the eyes of ammonia's incoming lone pair: Borane's empty orbital can be approached from above or below.  As the ammonia approaches one side of borane, one empty lobe enlarges to accept the lone pair while the other shrinks. Also, borane's little hydrogen limbs fold back away from the incoming lone pair to accommodate the fit.  The final coupling product looks like this:

BH₃, with the help of ammonia's lone pair, now has an octet of electrons. 

I Must Be Bored Today

The fifth element, boron, is essential for healthy plants. I once wrote a research proposal focused on that topic. Challenging the notion that boron's role in plants is structural in nature, I proposed instead that boron plays a dynamic role in plant tissue (cellulose) growth instead. The proposal went nowhere, maybe because I was wrong, but also perhaps because I was way out of my element. I'll never know the difference.

Whether we humans require boron is still in question: rats apparently do-though only in miniscule amounts. But boron is absolutely crucial for our silicon-based life support systems, in particular, for our computer chip-driven way of life. Boron's primary use however is boring and mundane: the bulk of it is used in borosilicate (pyrex) glass, commonly used in chemistry labs and in cookware; pretty much the rest of it is used in laundry detergents.

Now the word boron ends in "on" just like carbon, but this wasn't always so. Boron used to go by names like "bore" and "boracium" among others. The name "boron" didn't stick (in the English-speaking world at least) until early workers recognized its profound chemical similarity to carbon (btw, I found the ultimate authority on chemical name origins, so unless I think of something original to add, I'm just going to link to van der Krogt from now on. Notice that one can scroll up and down by atomic number or alphabetically. It's a totally cool website IMO *jealous*. And it figures that the guy is Dutch too).

Boron sits atop an imaginary line running diagonally down across the periodic chart which divides the metals from the non-metals:

Boron is one of the so-called metalloids, having chemical properties intermediate between metallic and non-metallic, just as the name suggests. What makes some elements metals and other non-metals is pretty well explained by band theory, but a simplistic view is to invoke electronegativity. Electronegativity is like a measure of electron selfishness: some elements cling so selfishly to their own electrons that they're unwilling to share them, not even with their identical neighbors-no conductance! So the diagonal line also demarcates a certain threshold electronegativity, the non-metals being more electronegative, epitomized by fluorine in the upper right (not shown).

Boron's chemistry is dominated by its electron deficiency, a consequence of its possessing fewer valence electrons than it has available electron valence orbitals.  Like carbon, boron builds borocentric molecules, typically with up to 4 other atoms surrounding it. Those four neighbors require that boron contribute four electrons to the magic octet, but boron only has three valence electrons to share (neutral boron actually has five electrons to match the +5 charge of its nucleus, but two of the five electrons are permanently locked away in a helium-like configuration: the filled so-called K-shell in the Kos reference). So when forming tetravalent compounds, boron always comes up one electron short and has to borrow one [LOL-boron builds more house than it can afford!].

Boron is used in the semi-conductor industry to introduce electronic "holes" in the atomic structure of chip material like silicon. When a boron atom assumes a position in a silicon structure normally occupied by a silicon atom, there is a bond missing one electron (or in other words, a hole). These holes facilitate the movement of positive charges or "holes" through materials. Other atoms can be doped into chips to do the opposite, i.e., introduce negative charges (i.e., electrons) and when used in conjunction, truly miraculous things happen which actually allow you to read what you're reading here.

This concludes another boring lesson in the chemical elements.