Part 27: The Chem in SpaceChem - Round 8
The Chem in SpaceChem, Round 816: Halogen Sorting
Input molecules: ‘Chlorine’
Output molecules: ‘Bromine’
Net reaction: None
’Chlorine’, Chlorine radical, Cl•
These values are for Cl2, as data for Cl• is not available.
Melting point: -101 °C, -150 °F
Boiling point: -34 °C, -29 °F
Molecular mass: 70.91u (It’s half that amount, 35.45 for single Cl•)
Density: 3.21 g/L
Chlorine in its normal diatomic form Cl2 is a toxic green-yellow gas (chloros means ‘green’ in Greek). It is produced by electrolysis of brine and as an element it’s mostly used as a disinfectant, for instance in swimming pools, and as bleach, although bleach often contains other chlorine compounds as well.
It readily forms compounds, the most common one is table salt, NaCl. In salts, chlorine is often in the Cl- form, which is called the chloride ion. Chloride ions are quite the opposite of chlorine: they are stable non-toxic particles. Chlorine can form many other salts, with a broad spectrum of uses. It also forms hydrochloric acid, which is a strong acid that’s common in the lab and a main component of gastric acid. Other than that, it can be used to make many organochloride compounds, with again a broad spectrum of uses. For instance, polyvinylchloride is a common plastic, chloroform is an important organic solvent and dichlorodiphenyltrichloroethane (DDT) is an insecticide.
Now, what we actually have here is a chlorine radical, a single uncharged chlorine atom. These atoms have an unpaired electron, making them extra reactive. They can be produced by heating chlorine or bombarding it with UV-radiation of a certain frequency. They are also formed in the stratosphere by decomposition of chlorofluorocarbons (CFCs), where they catalyze ozone decomposition, destroying the ozone layer. The problem is that most other compounds have an even number of electrons, so the unpaired electron just keeps hanging around, causing more and more havoc. The reaction only stops when two radicals react with each other, and the two unpaired electrons pair up to form a Cl-Cl bond. This is the reason that CFCs have been banned in most countries.
’Bromine’, Bromine radical, Br•
Once again, the values for Br2.
Melting point: -7.4 °C, 19 °F
Boiling point: 59 °C, 138 °F
Molecular mass: 159.8u (It’s half that amount, 79.90 for single Br•)
Density: 3.10 kg/L
Bromine is one of the two elements that are liquid at room temperature (the other one is mercury). It is a fuming red-brown liquid, with a strong smell similar to chlorine.
Actually, bromine is in a lot of ways similar to chlorine. It is somewhat less reactive, but it can be used as a disinfectant and as a bleach. It can still form a lot of compounds, such as NaBr which looks like NaCl but has been used as a sedative and anticonvulsant in the past, and AgBr, which is the photosensitive compound in common kinds of photographic film.
Bromine also forms organobromides, which can be used as pesticides, dyes and flame retardants. Bromine has no normal function within the human body, though.
The bromine radical we have here works similarly to the chlorine radical. While certain bromine compounds can cause ozone depletion, bromine radicals are actually used in a good way in flame retardants. Within a fire, there are always some oxygen and hydrogen radicals as intermediates. Those keep the fire going. When bromine encounters an oxygen radical, it will react to form the less reactive bromine radical, stopping the oxygen radical chain reaction.
Feasibility of the reaction
Feasibility: high.
Well, we’re just purifying chlorine and bromine from a mixture, right? That’s easy enough. A ‘box or radicals’ isn’t stable so I’d suggest purifying Cl2 and Br2 instead. Either do a distillation or something based on difference in solubility.
Reaction energy: Very slightly endothermic
Purifying substances from a mixture means increasing the amount of ‘order’ in the substance, which is a decrease of entropy (‘chaos’). A decrease of entropy always costs some energy.
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17: Magnets
Input molecules: Iron, Oxygen
Output molecule: Magnetite
Net reaction: 3 Fe + 2 O2 --> Fe3O4
Iron, Fe
Melting point: 1538 °C, 2800 °F
Boiling point: 2862 °C, 5182 °F
Atomic mass: 55.85u
Density: 7.87 kg/L
I haven’t talked about a pure metal yet, have I? Well, iron is certainly a metal. All metals have a similar structure: a bunch of ordered positive metal ions, with loose electrons jumping freely from ion to ion. The free movement of electrons allow metals to conduct electricity with ease, and their fleeting association with the ions cause the bonds that keep the metal together. Those bonds are simply known as metal bonds. It is important to realize that like salts, metals do not form distinctive molecules but crystals. The bonds are not as strong though, so metals are typically much more malleable.
Pure iron can bend rather easy, too. It will become much harder by adding just a few per cent of carbon and possibly some other metals, forming an alloy known as steel. Iron is a very common element and it makes up most of the Earth’s core. Its natural magnetism causes the Earth’s magnetic field which deflects dangerous solar radiation away from our planet. It’s not a coincidence that iron is so common. As I explained for an earlier Fusion puzzle, nuclear reactions in stars come to a natural stopping point at iron.
Iron has been used by mankind since ancient times. The ‘official’ iron age starts around 1200 BCE, but there’s archeological evidence that people did some basic iron work since at least 3000 BCE. Iron is cheap, common and strong, making it ideal for construction and other purposes. However, when it comes in contact with water or oxygen, it rusts (oxidizes) easily, forming several iron oxide and iron hydroxide salts. Fine enough iron powder will even burn. Stainless steel is often used to overcome these problems.
In biology, iron plays a number of roles. It is a part of the heme-group in hemoglobin, the oxygen carrying molecule which gives blood the red color. Interestingly, Antoine Lavoisier, who is known as the Father of Modern Chemistry, among other reasons because he was the first chemist to do quantitative measurements, used iron in an experiment to prove that water is not an element, but a compound made of two gases.
Oxygen, O2
Melting point: -219 °C, -362 °F
Boiling point: -183 °C, -297 °F
Molecular mass: 32.00u
Density: 1.43 g/L
Oxygen gas is a dangerously reactive compound that is deadly to all life.
Or at least, it once was. While oxygen is a common element, oxygen gas isn’t very common in the universe. Oxygen likes to react with things, forming metal oxides, carbon dioxide, water, you name it. Young Earth didn’t have any oxygen gas in its atmosphere. Early life was anaerobic, but oxygen gas was released as a waste product in certain photochemical reactions. This oxygen first reacted with all non-noble metals on the surface. After that was done (we’re talking geological ages here, a billion years or so), the gas started to build up in the atmosphere, and dissolved in water. This was problematic for all lifeforms, they couldn’t use it and it was toxic to them.
Some species of prokaryote managed to adapt: it found a way to ‘fix’ that nasty oxygen in other compounds in a safe way. Coincidentally, this reaction released a lot of energy that the micro-organism could use. This proved to be such an advantage that oxygen-using lifeforms quickly took over the planet. Some of these prokaryotes merged with larger cells, forming the organelle (cell-organ) we now know as the mitochondrion. Mitochondria still are the eukaryote’s “power plant” and they still have their own DNA. Anaerobic bacteria still exist in places where no oxygen can reach, and where there’s another source of energy, such as near volcanic vents deep in the ocean.
Oxygen gas constitutes about 21% of the volume of air, and that’s only because photosynthesis keeps producing it. It is expected that if aliens are looking at our planet, it’s the large amount of oxygen gas that will make them think there’s a very unusual process happening on Earth, maybe even some form of life.
Oxygen was first discovered by Antoine Lavoisier. Before him, it was believed that combustible materials contained some substance called ‘phlogiston’, which would leave the material in the form of smoke when it burned. Air could become saturated with phlogiston, making the flame go out. It made a lot of sense when you didn’t do actual measurements. Lavoisier did quantitative experiments and found that the mass of a material would increase as it burned, and that it would equal the decrease of mass of the air. He proposed the law of conservation of mass and proved that during combustion, a material combined with a gas in the air (oxygen).
Magnetite, Lodestone, Iron(II),(III) Oxide, Fe3O4
Melting point: 1597 °C, 2907 °F
Boiling point: I don’t know, but I suspect it decomposes first.
Molecular mass: 231.5u
Density: 5.17 kg/L
We’ve seen the mineral FeO (Wüstite) in a previous update. The black mineral Magnetite is a lot more common. It might not look like it, but it actually contains three types of ions. Each oxygen ion has a charge of -2, which means that for a neutral compound, the 3 iron ions need a total charge of +8. In fact, two of the iron ions have a +3 charge, and the third one is charged +2, leading to the Iron(II),(III) Oxide name.
Magnetite has been known as a material that ‘attracts iron’ since ancient times. Its name is derived from Magnesia, the name of a region and ancient city in Greece, where the mineral naturally occurs. Lodestones have been used in magnetic compasses for a long time. Magnetite has also been found in the brains or beaks of some animals, such as homing pigeons, where it acts like a built-in compass. It is used in magnetic storage devices such as hard disks and cassette tapes.
Coincidentally, last Sunday a new study was published on using magnetite in electronics. It can be used in transistors as a very fast (picoseconds) switch, which could speed up electronic computing by a lot. [1]
Uses of magnetite that don’t depend on its magnetism are as a black dye and as a sorbent to remove poisonous arsenic from drinking water.
Feasibility of the reaction
Feasibility: high.
This reaction is a simple oxidation of iron, and it does occur in nature. However, it seems that at Earth surface conditions, the formation of hematite (Fe2O3 with the recognizable red color) is favored. I’m not exactly sure what conditions are needed to make magnetite, but it can’t be too difficult. Natural magnetite has formed from iron and oxygen too, after all.
Reaction energy: Quite exothermic, about -1121 kJ/mol
As expected for iron oxidation. We’re basically burning the iron (in a specific way, so we get the right oxide), and a lot of heat is released in that process (even if it goes slowly, without flames).
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Challenge 8: Dessication Station
Input molecules: Na+, OH-, H2O
Output molecules: NaOH, H2O
Net reaction: Na+(aq), OH-(aq) --> NaOH
(aq) stands for aqua, it just means ‘dissolved in water’.
Basically, in this reaction we’re removing the ions from water, forming solid NaOH. I talked about water in this post, so I’ll skip that one now. Ions do not occur by themselves, so I’ll talk about sodium hydroxide and sodium hydroxide solution.
Sodium hydroxide, caustic soda, lye, NaOH
Melting point: 318 °C, 604 °F
Boiling point: 1388 °C, 2530 °F
Molecular mass: 40.00u
Density: 2.13 kg/L
As a solid, NaOH is a white salt commonly sold in pellets. It attracts water. The main use of solid NaOH is actually to make solutions of a specific concentration. NaOH is a base. A base is any substance that can produce OH- or remove H+. If a substance removes H+ from a water molecule, of course OH- is left, so those definitions match. In water, NaOH completely dissolves into Na+ and OH- ions, making it a strong base.
One advantage of a NaOH solution is that the sodium atom is quite boring and doesn’t do much. It just forms a table salt solution when it reacts with hydrochloric acid. NaOH is ‘just’ a strong base without unexpected side-effects. It’s used a lot in the lab. It’s also used in the paper industry in making pulp from wood, as a cleaning agent and as a drain cleaner. Concentrated NaOH solution is dangerous, as it causes chemical burns when it touches the skin.
Feasibility of the reaction
Feasibility: high.
Perfectly possible. Just boil off the water. Be careful though, because as you’re doing this, you will form a hot, very concentrated NaOH solution, which is dangerous.
Reaction energy: Endothermic, 44.51 kJ/mol
That value is the opposite of the dissolution energy, the energy which is released when Na-OH ionic bonds are broken and replaced by interactions between the ions and water molecules. In principle, you need to put in this energy to produce solid NaOH from the solution. In practice, you need more because you need to boil off the water. That’s actually the case for any reaction which involves a phase change, but in this case it’s particularly noticeable.
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Challenge 9: Oxygen Candle
Input molecules: K+, (ClO4)-.
Output molecules: KCl, O2
Net reaction: KClO4 --> KCl + 2 O2
I decided to write the reaction like this, as both KClO4 and KCl are ionic compounds, and it’s strange to write one as separate ions but not the other, when both seem to be in the solid state.
Potassium perchlorate, KClO4
Melting point: 400 °C, 752 °F (decomposes)
Molecular mass: 138.55u
Density: 2.52 kg/L
There’s some confusion about the melting/decomposition point. Wikipedia has a source which actually lists a melting point which is higher than the boiling point (610 and 400 °C). Another page lists a melting point of 525 °C. But most sites seems to agree that it just decomposes at 400 °C, so there’s no boiling point at all. It’s a bit of an average, as the melting point changes during the decomposition reaction.
Anyway, KClO4 is a strong oxidizer, which it can do by releasing oxygen. It is quite powerful, and it’s used in fireworks, ammunition percussion caps, rocket propellant and flash powder. Flash powder, which was commonly used before photo cameras got an electronic flash, is often a mix of potassium perchlorate and aluminium powder. The perchlorate quickly oxidizes the aluminium, causing a bright flash. Potassium perchlorate is quite similar to potassium chlorate, but it is a bit less likely to blow up in your face. It still gives a fast energetic reaction with sugar, there’s some nice clips of that on Youtube if you want to check it out.
Potassium Chloride, ‘muriate of potash’, Sylvite, KCl
Melting point: 770 °C, 1418 °F
Boiling point: 1420 °C, 2588 °F
Molecular mass: 74.55u
Density: 1.98 kg/L
Potassium chloride is a salt with a structure similar to sodium chloride, although it tastes bitter. It still has some use as a food additive, in which case it’s usually mixed with sodium chloride in order to lower the sodium content without changing the taste too much. Potassium ions are essential for the human body, so this doesn’t cause any harm. The main use of this salt is as a fertilizer, because plant growth depends on the amount of potassium available. It occurs naturally as the mineral sylvite, but only in very dry places, as it attracts water and dissolves easily.
By the way, like any potassium compound, KCl is slightly radioactive. About 1 in 10000 potassium atoms are a radioactive isotope. A large enough bag of KCl is enough to do some simple radiation experiments in the classroom, and potassium ions are the main source of radioactivity within the human body. In an average 70 kg person, about 4400 potassium nuclei decay per second.
Feasibility of the reaction
Feasibility: high.
This one is easy. Just heat up KClO4. The decomposition products will be KCl and O2. However, if your sample is contaminated with only a tiny amount of organic molecules, heating will cause a dangerously energetic reaction.
Reaction energy: Very slightly exothermic, -6.56 kJ/mol
That’s a theoretical value like usual. In this case, I happened to find a measured value as well, at -4.02 kJ/mol. Not a big difference.
I didn’t expect this reaction’s energy to be this close to neutral. I guess it can be explained by the fact that the energy is ‘in the oxygen’, and it doesn’t matter much if the oxygen is elemental or stored in the perchlorate.
In practice, for oxygen candles, which are basically emergency devices that will release a few hours worth of oxygen upon activation, mixtures of compounds are used. Potassium perchlorate is mixed with iron. When the reaction starts, some of the oxygen atoms will react with the iron, releasing heat. That way, the reaction will reach the decomposition temperature at which O2 is formed without the need for external heating. These oxygen candles have to be carefully designed, though. We don’t want a perchlorate-fueled iron fire to break loose.
[1] http://www.nature.com/nmat/journal/...df/nmat3718.pdf