Coppery keys to chemistry

Here we are again, confronting another interesting chemistry topic. It’s back to the periodic table today, as we continue our march to the Nobel group. Before we go any further, I feel like I need to restate again what an incredible thing the periodic table is. As much as I joke about how dry the table can be, the fact that we have such a visual, user-friendly and logical method of ordering elements by physical and chemical characteristics might just be one of the most incredible accomplishments of chemistry. No matter how complex the chemistry, everything that happens is based on a relatively small number of laws, some general tendencies of matter, molecules and charges, and the characteristics of elements. The periodic table lays all of those characteristics out in a way that can be almost instantly grasped. No endless lines of text, no subdivided columns, cells and subscripts of subscripts. Can you imagine what a purely text periodic table would look like? Probably nothing good.

Tried to find an image, unable to find someone crazy enough to make an all-text table

 So, today, we’re going to be looking at that upper right section of the marvelous table. One of the fun things about moving towards the right side of the table is the increasingly cool names that the element groupings acquire - none of this “metalloid” business here (descriptive enough, but rather bland). We’re in to the vaguely alchemical-sounding Greek-derived roots here, and today’s sterling example of nifty names in science is the chalcogens.

 

Chalcogen, for those of you not as up on your ancient Greek, translates (roughly) to “copper former”. In addition to the chemical characteristics we’re about to explore, the chalcogens all tend to be found in copper ores. Keep the “-gens” ending in mind, as we’ll be encountering it again next week.

 

They're in here somewhere...

So, besides hanging out with copper, what do the chalcogens have in common? First of all, they are not metals. Additionally, oxygen, sulfur, and selenium all have 4 electrons (out of a maximum of 6) in the outermost valence shell (as does tellurium, but it is counted as a metalloid). Remembering back to what we’ve seen of how atoms achieve that coveted full outermost valence shell, a mostly-full valence shell is much more likely to gain additional electrons than it is to shed electrons.

Chalcogens can either gain two electrons and become an anion with a charge of -2, or else form covalent bonds (typically with other non-metals). In fact, covalent bonds involving chalcogens are probably the single most important bit of chemistry on the face of the Earth. Remember back when we took a look at hydrogen bonding, that weird form of bonding that was only possible because the covalent bonds between oxygen and hydrogen atoms in water molecules contain some charge? Those semi-charged bonds are a direct consequence of the fact that water, as a chalcogen, is more likely to gain an electron than to lose it. Even when in a covalent bond, electrons are going to be drawn to the chalcogen. And that is how we wind up with the partially charged covalent bonds that make hydrogen bonds possible. 

 

Thank you, electron-gaining tendency of chalcogens.

Oxygen is easily the most recognizable element within the chalcogens - the name recognition may have something to do with the fact that we breathe oxygen. Without turning this into a post about human physiology, I’ll say that we, along with quite a few life forms on Earth, use oxygen at a cellular level in order to convert sugar molecules into energy through aerobic respiration. Oxygen is quite reactive, and capable of combining with just about anything. Oxygen is also abundant - it makes up 21% of the atmosphere. This was not always the case. Way back in the Archaean, billions of years ago, there was relatively little oxygen present in the atmosphere, and those organisms in existance practiced anaerobic (no-oxygen) respiration. However, oxygen is by-product of photosynthesis. The rise of photosynthetic organisms raised the oxygen content of the atmosphere, which, in turn, made it possible for more organisms to use the more efficient aerobic respiration, leading to life as we know it.

Not pictured: life as we know it.

Enough about oxygen for now. How about sulfur? Well, it’s also fairly common. Sulfur is the tenth-most abundant element in the universe, and crops up with some regularity in both mineral compounds and fossil fuels. The later can be an issue, as burning fossil fuels releases to the atmosphere, among other things, sulfur. Once there, the sulfur forms acid rain, with corresponding consequences for any ecosystems receiving that low pH precipitation.

How did sulfur wind up in fossil fuels in the first place? Aren’t those largely hydrocarbons? Well, yes, mostly. Fossil fuels are the fossilized remains of plants and animals, which, like most life on Earth, are mostly carbon, hydrogen, oxygen and nitrogen. Those four elements can be expanded out by the addition of sulfur and phosphorus to a group of six elements, which make up the vast bulk of all organic molecules. So, those fossils in the fuel may well have contained sulfur prior to being fossilized.

Alternately, sulfur came into the picture a little bit later. Sulfur belongs to the same group as oxygen, which would imply that sulfur behaves at least somewhat similarly to oxygen. And, in fact, there is an entire class of bacteria that live in non-oxygenated environments, and use sulfur to fulfill much the same role as oxygen in respiration. The process of converting organic matter into fossil fuel commonly involves these sulfur-oxidizing bacteria, leading to sulfur being incorporated in fossil fuels, and released upon burning.

It's round, it's pink, it breathes sulfur. Cool!

How about selenium? Well, selenium is the most metallic of the chalcogens, to the extent that it is occasionally included in the metalloids, and can be used as a semi-conductor. Selenium occurs in minerals (including that aforementioned association with copper ores), and is a nutrient to plants and animals, in small doses. Selenium is toxic at high doses - in certain soils, selenium is highly available, and accumulates in plants. When cattle eat those plants, they develop selenium poisoning - a fun fact admittedly not overly connected to the chalcogenic properties of selenium, but of note to any readers thinking of starting a cattle farm. Test your soils for selenium!

The chalcogens are a small group, but what they lack in numbers, they make up for in how many roles they play in chemistry. It figures that, with elements going up to number 118 (and rising), you're going to get a certain number of overachievers like oxygen. The nonmetals are a real cluster of such overachievers, as you might guess by the fact that we had to separate seven elements over two weeks, when the 38 transition metals all fit neatly into one post. It may be the natural bias towards organic chemistry of an organic life form, but I would argue that much of chemistry is contained within this small group of elements, together with next week's subject, the halogens. Stay tuned!

Bomb(ardier)s away!

The astute of you may have noticed that we’re running a little bit behind here, post-wise. So, I’m going to switch things around a little bit, and give us another break from the periodic table. It’s still Chemistry Wednesday, but we’re going to take a look at another piece of applied chemistry, this time, chemical ecology. Or chemical entomology, if you want to call it that. Or exploding beetles of doom, if you want to be awesome and call it that.

Definitely exploding beetles of doom.

Yes, the bombardier beetle, an organism so bizarre that its very existence is a fundamental part of the intelligent design argument - the beetle displays such an off-the-wall defense mechanism that it couldn’t possibly have evolved. Well, actually, it could have, and it did, and we’ll see why and how in a moment. But first, the chemistry part. And, in the interests of full disclosure, I must say that the guy who discovered all of this was one of my favorite professors in college. Hats off, Dr. Eisner.

Well, actually, first a crash course in the beetles. Bombardier beetles belong to the Carabidae family. These are ground beetles - they’re quick and nimble scuttlers, but not particularly strong or fast flyers. When Carabidae beetles get in trouble, they need a response other than taking wing, and boy, do bombardiers ever have a response.

Imagine that you are an ant. There, on the ground ahead of you, is a large, potentially succulent bombardier beetle. It’s got small, tightly furled wings - it’s going to be a ponderous process of getting airborne, and you’re a quick little thing. So, you run at the beetle. Bang! You’ve suddenly been hit with an extremely well-targeted, 100 degree Centigrade, chemically irritating fluid. What just happened?

The ant's eye view of what just happened.

Say hello to benzoquinones! And some other things. Benzoquinone, specifically 1,4 benzoquinone, is an extremely strong irritant, one that can cause blisters, blindness and respiratory damage in humans (and most other life forms). Clearly, being able to squirt this stuff at an attacker is a good survival strategy. It’s like taking the painful, itchy experience of encountering a poison sumac, cubing the pain, and subjecting anything that gets near the tree, not just touches it. Bombardier beetles aren’t overly fond of benzoquinone, mind. Part of the precision inherent in spraying is to allow the beetles to avoid accidentally spraying themselves.

Bombardiers avoid any negative effects of carrying this stuff around inside of them by, well, not carrying it around inside them. Instead, bombardiers carry precursors to benzoquinones, and combine these precursors when threatened, in order to synthesize the irritant spray. Bombardiers have a specialized arrangement of glands at the rear of the body - two glands empty into a reaction chamber, which then empties into a firing mechanism. The reaction chamber, as you might expect, is made of rather resistant tissue.

With a name like "explosion chamber", it would almost have to be.

One of the two glands contains hydroquinones (a chemical that can be an irritant in sufficiently high doses, and is involved in regulating melanin) and hydrogen peroxide, which appears in everything from topical antiseptics to rocket fuel. The other chamber contains a mixture of enzymes. When the two glands are emptied into the reaction chamber, enzymes break hydrogen peroxide down into hydrogen and oxygen. The oxygen then converts the hydroquinones to quinones by removing two hydrogen atoms from the hydroquinone molecule.

Not pictured: heat

Ok, that explains the synthesis of a toxic spray, but not the explosive burning part. Ah, but it does. Remember that bit where hydrogen peroxide was split in half, generating oxygen? As we’ll see next week, oxygen is a very reactive element. That’s where the explosion comes from. The heat is the result of the reaction being exothermic. Exothermic reactions release heat. Sometimes, the heat is barely perceptible, and sometimes it heats the resultant solution to the boiling point. Really, it’s just a matter of magnitude. And for the record, the reaction is controlled, so there’s no risk of the beetle exploding - the reaction chamber can withstand the force of the explosion, particularly since said explosion is directed outward.

Now, we’ve taken a look at how the chemistry works. How on Earth did all of this evolve? Well, as tends to be the case in evolution, bombardier beetle spray arose through a sequence of relatively minor tweaks to existing mechanisms. Let’s run this backwards.

There is a beetle in the Carabidae family that synthesizes the same irritating mix of quinones, but rather than shooting them out at high temperature, exudes them as a high-temperature foam which bubbles off the resistant carapace of the beetle (the foaming action should be familiar to anyone who has ever used hydrogen peroxide to clean out a cut). The reaction is still exothermic and explosive, but it does not take place within a reaction chamber that empties into a firing mechanism. You might see a flaw with this method, however. By exuding the irritant as a foam, the beetle can only deter predators that have made it onto the body - a really determined predatory might be able to hang on. Beetles which could spray the exudation would have a higher chance of surviving encounters with a predator. It’s a relatively minor evolutionary matter for the reaction chamber to become more rigid, and for an aiming mechanism to develop.

 

It's the Windows 98 of bombardiers. Almost there.

All right, but how do we get to the hot, toxic exudation in the first place? Well, any number of insects secrete hydroquinones as defensive compounds. Most insects also produce small amounts of hydrogen peroxide as a byproduct of metabolic reactions. Over time, a beetle which initially secreted hydroquinone via a gland and duct arrangement may have begun to combine hydrogen peroxide with hydroquinone, generating both a more toxic secretion, and a secretion with more temperature oomph. It’s easy to see the evolutionary progression from a beetle which exudes a hydroquinone solution to a beetle which combines hydroquinone with hydrogen peroxide before exuding it.

Hydroquinone, in turn, is a tweak on a more primitive defensive compound, quinone. Quinone was not originally a defensive compound - it had other, structural roles to play in insects, but eventually began to be used as defensive secretions, as they have an unpleasant taste. What was that original structural role? That would be protective pigmentation - various quinone derivatives are widely used in order to provide some protection from solar radiation. You might be familiar with one of those derivatives - melanin.

So, however exotic the chemistry of the bombardier beetle may be, just remember, it’s just an evolutionary hop, skip and jump away from human skin tone. That’s wicked cool.

X-treme botany

I feel like maybe I’ve been giving botany a bad rap, what with all of these examinations of plant families. I’ve been focusing on the so-called commercially important plant families, and of those, the ones that are commercially important mostly due to being delicious. This might give the impression that all the incredible genetic and functional diversity of plants exists only to give us more things to eat. Well, today we’re going to switch it up, and look at some seriously hardcore plants. Plants that live in incredibly hostile environments. Plants so tough that when you talk about how extreme they are, you have to spell it with just an x.

I got your fertile soils right here...

There are a lot of ways that an environment can be extreme, but today, we’ll focus on chemical extremes:  pH (acid and alkaline), and salinity. Believe me, this will keep us busy for quite a while.

Let’s start with pH. Extreme pH can be defined as anything falling outside the range of 5.5 to 6.5, which is where the bulk of plants do best. Plenty of commonplace graden plants like azaleas, dogwoods and blueberries thrive at a more acid pH, often in the 5.0 - 5.5 range. Once soil pH dips below 4.5, the soil is considered “strongly acid”, and most plants begin to suffer from the acidity (for further detail, wander back to this post), mostly in the form of nutrient deficiencies. Simplifying matters greatly, let’s leave it at “nutrient uptake is impaired at low pH”.

  And how!

Most plants, but not all. Some plants do very well at extremely low pH - river birch (Betula nigra) is known to enjoy pHs as low as 2, while pitch pine (Pinus rigida) in addition to tolerating a fire regime tolerates soils down to a pH of 3.4. How do they do it?

Well, remember from the last time we took a look at acid soils, there are two main problems to deal with. One, nutrient availability is low. Two, aluminum (Al) availability is high (and aluminum is toxic). These problems have a common solution - mycorrhizae! 

 From left to right, the art view and the technical view

Mycorrhizae are a family of fungi that grow in a symbiotic association with many plant species. Mycorrhizae do all kinds of cool things, including making nutrients more accessible. The explanation is simple - associating with a mycorrhizal network gives plants access to a huge amount of soil, without having to put out all of those roots. When nutrient uptake is limited or nutrient concentrations are low, having access to more soils means more nutrient uptake. Ok, so that should solve the low nutrient problem, but shouldn’t it also increase the Al problem? Not quite. Acid tolerant plants have been shown to block out Al with mycorrhizae - rather than be brought into the plant, Al accumulates harmlessly on the surface of the plant-mycorrhizae juncture.

So, one end of the pH spectrum can be managed, with a little help from some fungal friends. What about the other end of the spectrum, the high-pH, alkaline soils? As that image of nutrient availability shows, limitations kick in at higher pH as well - iron (Fe), for example, is often less accessible. So, some successful alkaline-tolerant plants have evolved efficient nutrient usage to limit any limitations. Mycorrhizae aren’t just for acid soils, either - plants growing in alkaline soils have been found to benefit from the increased nutrient access made possible by such a symbiotic fungal relationship.

Just like acid soils are plagued by both low nutrient availability and high Al concentrations, so alkaline soils have more than limited access to Fe to contend with. Here, the problem isn’t Al, but rather, salt. Extremely alkaline soils (pH of above 9) are the result of accumulated sodium carbonate - while salt is not the main limiting factor in such soils, plants must content with the salt to thrive. Remember, as we saw in our examination of dune plants, salt is not a friend to all living things. Plants in alkaline soils have been shown to excrete salts accidentally taken up, or to store water within the leaf in such a way as to dilute salts.

This makes a great segue to saline soils. Saline soils are (duh) saline, but not necessarily due to high concentrations of sodium. High concentrations of calcium, magnesium, potassium, chloride, and bicarbonate, among other things, can make a soil saline. Saline soils can be directly toxic to plants - chloride and sodium are both capable of damaging plant cells and tissues. Nutrient deficiencies are also an issue - high concentrations of the saline ions can make it more difficult to bring up actual nutrients, and sodium actively impairs nutrient uptake. Then there is the water issue.

What a wonderful place to be a plant.

There are a number of controls on how water moves in an environment, one of these controls being the concentration gradient. In brief, water moves down a gradient of solute concentration - water with lower concentrations of solutes moves towards regions of higher solute concentration, and eventually reaches a uniform solute concentration. This becomes a problem when in order to survive, a plant needs to maintain a lower solute concentration in water within the root - water within the root is attracted to the highly alkaline water outside of the root. Additionally, when the plant goes to take up water, it is forcing water to move in the opposite direction of the gradient. So, how do salt-tolerant plants cope with this?

Actually, there a couple ways to do this, including a really complicated arrangement of cell membranes in the roots, but to my mind, the most elegant solution is to create an osmotic gradient within the plant. Certain plants sequester either organic compounds or ions from the saline soils within certain parts of the plant, at greater concentrations than the ions within the soils. This has the effect of making the concentration of solutes higher within the plant than without, and driving water into the plant. It takes less energy for the plant to use ions taken up from the soil solution, and this also has the effect of harmlessly sequestering salts brought into the plant. Pretty cool, huh?

This is only scratching the surface of the extreme environments plants laugh at. In future posts, we'll look at plants which thrive in freezing temperatures, drought conditions, and unbelievably windy mountaintops. Not to mention the bacteria out there that regard acid mine drainage, petroleum derivatives and toxic chemical spills as a delicious breakfast buffet. Excited yet? 

The fantastic four

Wow, can you believe that we’re nearly through the periodic table? I believe it was way back in March when I got the crazy idea to work through the entire thing. Fear not, Chemistry Wednesdays will continue, but the topics will get just a little bit more unpredictable. But, before we reach that point, we’ve got three more groups of elements to go through. We’ve looked at the metals, and the semi-metals. This can only mean one thing is coming up - the non-metals!

 

Helpfully marked in orange

Yes, the non-metals, a group of elements largely defined by what they are not. Also, a somewhat misleading name, as all non-metal elements are not metallic,but not all non-metallic elements are referred to as non-metals. Confused yet?

All right, remember in the good old days of the alkali earth metals, when everything was simple, and each group of elements simply consisted of one of those straight-down-the-table groups like Group I and Group II? Then along came the transition metals and screwed all of that up, by being very similar to one another across groups. As we move towards the right side of the table, groups become more distinctive relative to one another once again. The last two groups on the table are usually examined as separate categories, as is (sometimes) the third-to-last. We’ll look at those next week, as I don’t fancy covering carbon chemistry and oxygen chemistry in one post. Today, we’re going to look at a kind of catch-all group - those non-metallic elements that do not belong to one of the three distinctive final groups on the table. Don’t worry! There are only four of them!

Most of the elements in this group are clustered together. Most, but not all. Yes, today is the today when we finally look at hydrogen which, in spite of hanging out on top of Group I winds up being discussed along with carbon, oxygen (sometimes) and nitrogen (that little tetralogy of elements is the basis for an entire branch of chemistry, which we will get to in a moment).

Ok, so besides being excluded from the metal club, what exactly distinguishes non-metals from the pack? Well, they lack the properties of metals - ductility, ability to form alloys, tendency to efficiently conduct heat and electricity, etc. Non-metals also have more electrons in the valence shell, usually from four to eight electrons (remember the oddities of the d and f orbitals, which keep them from ever being proper valence shells). The non-metals we’re looking at today usually have four or five valence electrons (hydrogen, with only one electron, being the exception). With half-filled (or more) valence shells, the non-metals are more likely to bond in such a way as to fill the valence shell. Contrast this with the tendency of metals to shed valence electrons.

Non-metals, then, tend to either be the negative side of an ionic bond, or bond covalently. The elements we’re looking at today are all more likely to bond covalently - we’ll get to the ionic bonds next week. This tendency to form covalent bonds gives rise to a rather important bit of chemistry - organic chemistry.

Organic chemistry is defined as the branch of chemistry concerning carbon compounds. After carbon, the main elements involved in organic chemistry are hydrogen, oxygen and nitrogen, following the easy-to-remember acronym “CHON”. Phosphorus and sulfur are two common additional guests, and, not coincidentally, also in the non-metals group (although, again, oxygen and sulfur are sometimes separated out).

Organic chemistry is sometimes called the chemistry of life - it deals with the compounds that form all living tissue, from DNA bases to skin and sinew. Certain characteristics of the nonmetals, particularly carbon, are responsible for the primacy of organic chemistry in biology, so let’s focus on carbon for a little bit.

 

Thanks, organic chemistry!

Carbon (abbreviated C) is interesting, in that it is capable of covalently bonding with four other atoms at once, due to its 4 valence electrons (in an 8-max shell). Silicon, germanium and tin are also (theoretically) able to do this, but, to varying extents, all three of those elements are somewhat likely to give up some or all of their valence electrons, due in part to the presence of a d-orbital. Carbon is the only 4-valence electron element with no tendency to give up electrons, and thus almost always forms covalent bonds. Carbon is also the sixth-most abundant element in the universe. Combine these two things, and the fact that life on Earth tends to be carbon-based starts to make sense. Covalent bonds, after all, are possible between elements of all electron configurations (in contrast to metallic bonding), form easily at the standard range of temperatures on Earth (in contrast to a form of bonding we’ll look at in a few weeks), and don’t dissolve in water (in contrast to ionic bonding). That’s reassuring, isn’t it?

Nitrogen (N) has interesting bonding abilities, as well. Nitrogen has one more valence electron than carbon, and needs three additional electrons to form a full shell. Now, plenty of N atoms achieve this by bonding to three additional electrons, but some N atoms achieve this by forming a triple covalent bond to another N atom. In two triply bonded atoms, three pairs of electrons are shared, forming a very strong bond.Theoretically, any atom with the right range of electrons in the valence shell can form double or triple bonds (C, ever the overachiever, is capable of forming a triple bond with another C or N atom, and a single covalent bond with another atom, putting a certain amount of pizzazz into organic chemistry), but N is especially likely to form triple bonds. Two triply bonded N atoms form a stable molecule, as both have full valence shells. In its stable, 2-atom molecule, N exists as a gas. And not just any gas. Gaseous N makes up 78% of Earth’s atmosphere. Nitrogen is the fifth most abundant element in the universe, and plays all kinds of interesting roles in ecology. Which, one of these days, I will get into.

Phosphorus (P) behaves somewhat similarly to N, but a) the bonds aren’t as strong, c) there is a d orbital involved, and c) P has a tendency to spontaneously combust in air. Let’s make a very long story short and simply say that P has a very strong affinity for oxygen, and rapidly reacts with any available oxygen to form certain compounds. As you may remember from Group I chemistry, rapid reactions sometimes generate, well, fire. Who says organic chemistry is boring? Phosphorus is an important nutrient, and an ingredient in explosives. How many elements can you say that about?

That leaves hydrogen, our weird little outlier. Now, the reason we didn’t stick hydrogen in with Group 1 is that, in spite of having only one electron, hydrogen still has a half-full valence shell. So while hydrogen can lose that one electron, it’s just as likely to acquire another electron via covalent bonding. Hydrogen may exist as a gas (two hydrogen atoms can bond covalently forming a stable molecule), or in combination with another element - in fact, hydrogen naturally combines with every element in the periodic table, except for the ones at the bottom (they don’t exist in nature) and the Nobel gases (wait a few weeks for the answer to that one). Hydrogen is the most abundant element in the universe, comprising something like 90% of all matter. Hydrogen pops up everywhere from organic chemistry to fertilizers, to the hearts of stars, where it burns in a process that, if only we could understand exactly how it works, could fuel the future of humanity. 

Yeah, I'm a playa

Wow. That was a lot of information for four elements. Non-metals are quite the grab bag, eh?

Things that should not be in the ocean, Part I

Technical difficulties and life in general have not been cooperative in this week's rigid posting schedule. Fear not, we're getting a double post today.

 This past weekend, I was involved in a semi-serious conversation about using genetically engineered algae to turn plastic garbage into biodiesel. It’s an elegant solution, in a way - using the ever-growing avalanche of humanity’s own waste to power further growth. The idea raises two points - one, the science (and feasibility) of that kind of biodiesel production, and two, the availability of the medium (plastic waste). We’ll take a look at the microbiology part of the equation in a few weeks, but today, let’s look at plastic waste. And not just any plastic waste, but what I consider to be its most dramatic manifestation. Because really, there’s no way to start the week off right like the Pacific Garbage Patch.

 Wait, that’s it? I bet you were hoping for something more impressive. After all, isn’t the garbage patch a Texas-sized agglomeration of debris that can be seen from space? Shouldn’t it look something like this?

 

Not quite. The Pacific garbage patch is real, and a real threat, but it doesn’t quite look like that. It isn’t a solid island, and it can’t be seen from space. It doesn’t have a proper fixed location, either. Let’s look at this one term at a time, starting with “Pacific”. 

 

Ocean currents are fun. They arise from a combination of tides, differences in the density of water, and wind patterns (which, in turn, ultimately derive from the uneven heating of Earth’s atmosphere). In this case, we’re mostly concerned with the wind-driven currents. Global winds create what are called gyres within oceans - massive spirals of circulating water, with areas of relative calm in the middle. 

There are five main gyres out there - the North Pacific, South Pacific, North Atlantic, South Atlantic and Indian Ocean gyres. Way back in 1988, the National Oceanic and Atmospheric Administration predicted that there was a high likelihood of plastic debris accumulating in the calm area at the center of the North Pacific gyre. These predictions were confirmed by the discovery of high concentrations of plastic debris exactly where expected. It’s entirely possible (and highly likely) that debris are accumulating within the other gyres, however, the massive amount of “ground” to cover somewhat hinders the amount of sampling possible.

Ok, let’s move onto the “garbage” part of the phrase. The garbage is all plastic or plastic-derived - nothing else would float in water. The sources vary - about 80% comes from land, and of that, 65% comes from garbage. Garbage might be accidentally or intentionally dropped into a stream, washed off of a landfill (plastic has a regrettable tendency to float), or spilled in the process of being transferred from the home to the landfill. Other plastics make their way from the land to the sea when the plastic pellets used in manufacturing plastic goods are spilled into waterways. Finally, the remaining 20% or so is dropped, washed or dumped from ships directly into the ocean - this could be anything from a coffee cup washing off a fishing boat to a container being washed off a cargo vessel.

Once in the ocean, the plastics quickly degrade. Well, sort of. I’m going to swerve over into earth science here, and present a relevant (I think) concept. Weathering, in this context, describes the breakdown of rocks into smaller and smaller pieces (and eventually soils and mineral-rich water). There are two types of weathering, physical and chemical. Physical weathering is the physical breakdown of rocks - imagine a hammer smashing a rock into bits. The bits get smaller and smaller, but the chemical structure remains unchanged - quartz beach sand has the same chemical formula (SiO2) as a gigantic quartz crystal. Chemical weathering acts at the chemical level - the rock is dissolved or leached, and the resulting pieces have a chemical composition different from the rock pre-weathering. Plastic debris in the ocean can be understood to undergo physical weathering - the particle size is reduced, reduced, and reduced again, but the chemical structure remains the same. Even those compostable plastics that are coming into greater use don’t really dissolve in the oceans - compostable plastics are designed to dissolve under land conditions, not in oceans. You’re starting to see why “proper disposal” is so important with any kind of waste, right?

One more word about the garbage bit. The tendency of plastic waste in the ocean to become floating “microplastics” explains why you can’t see the patch from space. Much of the patch consists of a high concentration of microplastics in the upper portion of the water column.

 

Ok, so these are a little bigger than micro, but you get the point.

Now, a little more about the patch part of the phrase. We already know that the garbage patch is not actually a floating island of waste. Instead, it is an area with a high concentration of plastics relative to things that would normally be in the water column (plankton, for example) - typically within the patch, the concentration of plastic is greater than the concentration of plankton. The high concentration is a consequence of water movement being relatively restricted within the center of an oceanic gyre - plastic that finds its way in does not readily find its way out. The location and size of a given patch depend on the gyres. Now, is it possible for the location and size to shift?

Of course it is! Both the size and extent of the garbage patch change throughout the year, in combination with the shifting weather patterns. Given that the patch we’re focusing on is in the Pacific Ocean, there is also the (so-far uninvestigated) effect of the El Nino - La Nina cycle on the whole mess.

The patch is also, well, patchy. While concentrations of microplastics might be relatively constant throughout the entire estimated garbage patch, larger plastic debris are going to congregate in certain areas, as a result of more local tidal action. Reports of sailors in the North Pacific encountering floating piles of visible plastics are tied to these “sub-patches”. All of this makes it hard to pin down the exact location or extent of the garbage patch, and subsequently, hampers clean-up of the patch.

And do we ever need clean-up. All of this plastic in the ocean is not doing marine life any favors. There are the obvious problems, like suffocation or entanglement in larger pieces of ropy plastic - anyone who has ever cut up the plastic rings on a six-pack holder has been trying to avoid this outcome. There is also a more insidious threat, from the smaller plastics. A number of marine organisms feed on plankton, fish eggs, small fish, or other bits of tasty marine life floating on or near the surface. Now, what happens when the concentration of small pieces of plastic rises above the concentration of that upper-water column marine life?
 

Rather than insert a disturbing image, I suggest that the strong-stomached among you image search “albatross plastics autopsy”

Birds are particularly threatened. Take the increasingly unhappy case of the Laysan albatross, a truly remarkable bird. Like all albatross, these birds spend much of their lives flying over the open ocean - albatross are capable of both weathering harsh weather, and sleeping on the wing. Laysan albatross are particularly large representatives of the albatross family, with a wingspan that often reaches (or exceeds) 2 m. They only come ashore to breed, and that’s where the trouble starts. Laysan albatross prefer to feed their young on the fish eggs and small squid found in the upper water column. Much of the feeding range of Laysan parents happens to overlap with the North Pacific garbage patch. I won’t get into the (depressing) mechanics, but up to forty percent of Laysan albatross chicks born each year die from ingestion of plastics which their parents mistake for fish eggs. 

I shot the albatross

Clean-up of the garbage patch is difficult, maybe impossible. However, it’s very possible, even easy, to prevent the patch from getting any bigger. The damage is done, but we can prevent it from getting worse. Reducing, reusing and recycling plastics (in that order), along with taking pains to properly dispose of plastics (remember the biodegradable plastics problem) that can’t be reused or recycled can cut off the source for much of the plastic in the patch. Sure, there’s only so much we can do about containers being lost from cargo vessels, but that’s only 20% of the problem. The other 80% is up to you.