#24: Contact Detonation

Using my previously isolated iodine, along with some diluted ammonia (NH3), I synthesized a tiny bit of nitrogen triiodide (NI3). Due to the major size difference between the nitrogen atom and iodine atoms in the compound, nitrogen triiodide is extremely unstable. As a result, a tap with a stick will cause it to detonate. 

To see the demonstration performed on a much larger scale...

https://www.youtube.com/watch?v=H7oZafSywow

#23: Isolating Elemental Iodine

Experimentation is an important part of science. Unfortunately, it comes at a cost. That being, it costs money. This is especially true of chemistry experiments in which chemicals can be quite expensive. So while I would love to sit around throwing random chemicals together all day, it is simply not possible if I want to maintain a college fund. When it comes to my projects, I must be wise so as not to waste what I have. I also have to be resourceful with where and how I obtain my chemicals. A good way to do this is to buy them from places like pharmacies or hardware stores. Although they can be hard to find, many chemicals can be obtained much cheaper from commercial products such as cleaners. However, they are typically sold in a diluted form where they are often mixed with other substances. Fortunately, with a little bit of work, they can usually be purified.

Iodine (the 53rd element) is a very interesting chemical. On the periodic table, it lies in the same column as chlorine which makes it quite reactive. Despite being somewhat dangerous, it has many uses in the medical industry. While I would love to explore its chemistry, it costs nearly $30 per ounce. Plus, you often need to get a license just to buy it. Luckily for me, I can get it through the black market, aka CVS pharmacy. Yes, iodine is commonly found in antiseptics where it is usually called "tincture of iodine." These products contain a small percentage of pure iodine along with potassium iodide (KI) dissolved in water and/or alcohol. For me, this form of iodine is only useful if I want to keep myself from getting an infection. To make it suitable for chemistry experiments, I would have to purify it.

As I said before, the amount of pure iodine in antiseptics is very small. To obtain a useful quantity of the chemical, I knew I would need to isolate it from the larger portion of potassium iodide. To do that, I added an equal amount hydrochloric acid (HCl). The acid removed the potassium ion from the potassium iodide and replaced it with a hydrogen ion (HI). To separate the iodine from the hydrogen, I then mixed in a solution of dilute hydrogen peroxide (H2O2) which I also got from the drug store. The hydrogen peroxide molecules oxidized the hydrogen atoms whereupon water (H2O) was created. As for the iodine, it was now free. 

Here are the two reactions that occurred:
1. KI + HCl → KCl + HI
2. 2HI + H2O2 → 2H2O + I2

Because the iodine by itself was no longer soluble in water, it precipitated out in solid form. After all of it had gathered at the bottom of my beaker, I poured it over a filter and washed it with water. This left me with a black sludge of impure iodine.

To further refine my iodine, I relied on one of its fascinating properties. While most substances melt when they are heated, iodine sublimates. That means it goes from a solid state directly to a gaseous state. To take advantage of this, I dropped my globs of iodine in a beaker and placed it over my burner. As it began to heat, I covered the beaker with a watch glass to contain any gases that might escape. It wasn't long before a dark purple cloud of sublimated iodine appeared inside the beaker. As the gas thickened, I poured some cold water on top of the watch glass. This would cause the gas near the top of the beaker to cool and condense. Unlike hot water vapor, which becomes liquid when cooled, iodine gas becomes a solid. In doing so, it rapidly forms into crystals. In my case, the sharp crystals that appeared were astounding. When the iodine at the bottom of the beaker had disappeared, I turned off the heat. Once everything had cooled, I removed the watch glass which had been covered in a silvery layer of iodine. After scraping all of the crystals into a glass vial, I had my purified product. 

Here is the heated beaker containing sublimated iodine gas which cooled near the top to form tiny crystals.

My impure iodine slurry.

Here's the crust of pure iodine that formed beneath the watch glass.

Wet iodine before sublimation.

#22: Another Liquid Metal

Mercury is an awesome element. In addition to being a liquid at below room temperature, it has an extremely high density and surface tension. In fact, a solid steel ball placed in it will actually float. Unfortunately, mercury is also quite toxic. This, coupled with its high cost, means that I probably won't be getting my hands on it. However, it doesn't mean that there isn't an alternative.

If you've read my previous posts, you might recognize the element gallium. Gallium has a relatively low melting point of 86 ℉. While that's low enough for it to melt in my hand, the metal solidifies rather quickly. Plus, gallium forms a pesky oxide layer which prevents it from flowing freely. Some alkali metals such as cesium also have low melting points. However, they are extremely reactive to say the least. There are a few other low-melting metals but they are either too dangerous or too expensive. Fortunately, pure metals aren't the only option.

A while back, I explored an alloy of gallium and aluminum. The alloy was interesting because it took on different characteristics than its two components. In general, alloys are useful because they have properties that aren't found in pure metals. For example, bronze was important because, unlike copper or tin, it was both strong and resistant to corrosion. Similarly, many alloys have lower melting points than the metals that form them. These alloys are often used as solders. But is there a low-melting alloy to replace mercury? The answer is yes.

The alloy is called Galinstan. Its name is derived from the three metals that are used to make it: gallium, indium, and stannum (the latin word for tin). Galinstan's melting point is about -2 ℉. It is also nontoxic, which makes it a fantastic replacement for mercury in many devices such as thermometers. The fact that it is safe also means it can be explored by people like me. So after buying the necessary materials online, I went ahead and made the alloy for myself. To do so, I melted down 7 grams of gallium (86 ℉), 2 grams of indium (314 ℉), and 1 gram of tin (450 ℉). When the liquid metal cooled, I removed any unalloyed bits which had solidified. After purifying the alloy, I had my liquid Galinstan metal.

If you noticed, all three of the metals that I used had relatively high melting points. So how did their combined melting point end up so much lower? To figure this out, we must understand what determines a substance's melting point. To do that, let's first look at the concept of temperature itself. Temperature is actually a measure of kinetic energy. On the molecular level, kinetic energy is the constant movement or "jigglyness" of the atoms in a substance. So when something is getting hot, it means that the atoms inside are moving around more. And when a substance's atoms are moving around more, that substance begins to move around more as well. Eventually, that substance begins to melt, or in the case of a liquid, it evaporates. The important thing to note here is that different types of substances require different amounts of energy in order to melt or evaporate. This amount of energy is often determined by the molecular structure of the substance. So in the case of Galinstan, its molecular structure simply requires minimal energy to melt. Thus, it is able to remain a liquid at low temperatures.

Here you can see the shiny, mercury-like globs of my Galinstan alloy. 

Here are samples of my gallium, indium, and tin (from left to right).

For more information about Galinstan and its properties...
http://en.wikipedia.org/wiki/Galinstan

#21: Synthesizing Salt (Part 2)

In my last experiment, I created a salt solution by combining hydrochloric acid and sodium hydroxide. Out of curiosity, I decided to let the solution evaporate with the hope that crystals would form. Sure enough, after a few days, they did.

After filtering the solution and letting it evaporate on a paper plate, it formed beautiful crystals of table salt.

#20: Synthesizing Salt (Part 1)

I recently conducted another simple, yet fascinating chemical reaction. In doing so, I relied on two chemicals: hydrochloric acid (HCl) and sodium hydroxide (NaOH). Dilute hydrochloric acid, often called muriatic acid, is commonly sold as a powerful cleaning agent. Even in low concentrations it has the ability to strip cement as well dissolve most metals. In higher concentrations it represents a scary, yet useful lab chemical. Similarly, sodium hydroxide, also known as lye, has a variety of uses. As you might know, lye is an important component of soap, although its concentration is usually quite low. In its pure form, sodium hydroxide can destroy most metals as well as glass. Essentially, both hydrochloric acid and sodium hydroxide are extremely corrosive substances. So what dangerous purpose would I be using them for? Well, you'll see soon enough.

To start, I did some calculations to figure out how much of each substance I would need. I then weighed out granules of white sodium hydroxide and added them to a glass jar. I used a jelly jar so as not to harm one of my precious beakers. Next, I measured out my hydrochloric acid using a graduated cylinder. With my camera ready and at a good distance away, I poured the acid into the jar. As the chemicals reacted, they began to froth and boil. However, after about ten seconds, the mixture calmed down. The reaction's exothermic nature had caused the jar to become quite hot. The resulting product was a clear-white solution with a syrupy consistency. The question was, what was it chemically?

Considering the destructive nature of the chemicals used, you might expect this new concoction to dissolve diamonds. Yet, the two substances that I had synthesized were far more sinister. Behold sodium chloride (NaCl) and dihydrogen monoxide (H2O)! Wait what? Table salt and water? Yep, here's the chemical equation to prove it: HCl + NaOH → NaCl + H2O. See how that works?

This simple reaction represents a fundamental area of chemistry. As is implied, hydrochloric acid is an acid. Sodium hydroxide on the other hand is a base. As I pointed out before, both can be extremely corrosive. However, they operate in completely different ways. In other words, they represent opposite ends of the spectrum. In fact, they sit on opposing sides of the pH scale. At the same time, acids and bases are sort of like chemical counterparts. When they come together, they neutralize each other. And, in the majority of cases, they are converted into salt and water. It is important to keep in mind that salt is a general term for any ionic compound that is formed by the reaction of an acid and a base. However, in this reaction, the resulting salt is the same one that you put on your french fries. In summary, I was able to take two chemicals, both of which can dissolve metal, and transform them into nothing but simple salt water.

In order to make sure the solution was truly neutral, I tested it with litmus paper and adjusted it accordingly.  

To learn more about the concept of acid/base neutralization...

http://en.wikipedia.org/wiki/Neutralization_(chemistry)

#19: Dissolving Color

If you have ever looked through a chemistry textbook, chances are you are familiar with a certain image: that of a nerdy kid wearing nerdy goggles who is examining beakers full of colorful liquid. As if glassware filled with water and food coloring is supposed to inspire you in your battle with stoichiometry. While the scene may be cliche, it illustrates something significant. It may surprise you to know that colors represent a big part of chemistry. If you remember, colors are just different wavelengths of light. You might also recall that different substances reflect some wavelengths but not others. For example, a red apple reflects red light while absorbing most other colors. This is because the molecular structure of the apple's skin allows the non-red wavelengths to pass into it. Essentially, a chemical's color is a literal reflection of its molecular construction. This makes color a useful tool when you consider that chemistry is the study of matter.

As usual, I performed an experiment that related to this concept. To start, I added some water to the bottom of a large flask. I then stirred in a small amount of potassium permanganate crystals (KMnO4). When the crystals dissolved, I had a dark purple solution. Next, I prepared a second solution of hydrogen peroxide (H2O2) and sulfuric acid (H2SO4). With my camera ready, I began pouring the mixture into the large flask. As the two liquids met, something astonishing happened. Within seconds, the purple color of the first solution had disappeared. The contents of the flask had gone from dark purple to crystal clear in an instant.

The speed at which the reaction took place gave me an idea. I wanted to see what would happen if I reversed the positions of two liquids. So after adding the hydrogen peroxide and sulfuric acid to the big flask, I made another solution of potassium permanganate in a separate beaker. Now came the test. I held the beaker above the flask and slowly poured out its colorful contents. As I did, something magical happened. Upon landing in the flask, the purple solution became invisible. Even after emptying the entire beaker, the liquid in the flask remained as clear as water. The effect was simple yet quite satisfying.

Surprisingly, both of these demonstrations have a simple explanation. When the two liquids met, a chemical reaction occurred. While I won't go into detail, the hydrogen peroxide and sulfuric acid caused the potassium permanganate to break apart and form smaller pieces. But how does this explain the color change? Well, as a mentioned before, different substances reflect different colors. In the case of the potassium permanganate, it was reflecting purple light. However, thanks to the chemical reaction, it was transformed into different substances. And unlike the potassium permanganate, these substances did not reflect purple light. In fact, they barely reflected any light, which is why they appeared colorless.

The demonstration's simple yet satisfying nature makes it one of my favorites.

#18: Beautiful Bismuth

I recently created some scientific splatter art with molten bismuth metal.

As the hot bismuth solidified, it formed iridescent oxide layers in a variety of colors.

The liquid metal also created these vein-like patterns as its surface became stretched. 

Here you can even see tiny crystals that began to form.

This piece was made using a round mold which ended up leaking.