Eventually we'll get to some elements you can't readily set on fire, but that day is not today.
Meet Group II! Like the elements we met last week, each of these elements burn a characteristic color when introduced to flame, and are sometimes used in fireworks manufacture. In descending order, beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra) make up the alkaline earth metals. Group II is, in some ways, very similar to Group I. Group II elements are characterized by having two electrons in the valence shell (maximum eight). While they are somewhat less reactive than Group I elements, having two electrons out of a possible eight still makes Group II elements likely to shed those two electrons in a reaction.
Remember how Group I elements react so violently with water they sometimes set it on fire? Well, the alkaline earth reactions occur with the same mechanism (metal in question donates its two electrons to water, water molecule splits apart yielding hydrogen gas and quite a bit of heat energy), but less vigor. Compare this video to the fire-in-water video (you may need to keep pressing play). That fizzing is the hydrogen gas being released from the reaction, but at no point is sufficient heat released to set the gas on fire, even in the case of Ba (Ra is radioactive and reacts a little bit differently). So, the alkaline earth metals are a bit more stable than the alkali metals, but still too reactive to exist naturally in pure form (again, like Group I). These metals often exist in oxide forms - a given alkali earth atom shares its two electrons with an oxygen atom, to form a stable compound with an extremely high melting point. The melting points are sufficiently high that the oxides remain stable in fire - when the properties of these elements were first being investigated, this was referred to as remaining in “earth” form. Hence the “earth” in “alkaline earth metal”.
So, what do the alkaline earth metals do when not fizzing in water? Be, Mg and Ca are all fairly widely distributed in naturally occurring compounds. Beryllium occurs in several minerals, including emeralds (aka beryls). Beryllium is the least reactive of the alkali earth metals - it does not readily form oxides or dissolve in water. So, Be is sometimes used in its pure form - as such, Be is the lightest solid metal (albeit a brittle one), and has applications in aerospace and nuclear engineering. Be is alloyed with other metals to produce strong, light, unreactive compounds that appear everywhere from golf clubs to dentistry tools. Beryllium has a sweet taste, but don’t go putting it in your tea, as the stuff is toxic in pure form.
I could do this entire post on Mg and Ca, but as I’m already writing a graduate thesis on those two elements, I’ll try to resist that urge and keep it brief. Magnesium, despite never being found in pure form in nature is the eighth-most abundant element in the universe, and the seventh-most abundant element in the crust of the Earth. Magnesium features prominently in such common minerals as dolomite, biotite, hornblende and chlorite. Magnesium salts are a component of seawater, and Mg is an essential nutrient for plant and animal life. At some point in the future, I will do a post on ecosystem acidification, and the exciting tree diseases connected to Mg deficiency. If you’ve ever ingested milk of magnesia or used Epsom salts, you’ve experienced the power of Mg. Similar to Be, Mg is alloyed to create an array of strong, light metals.
Hornblende: a good source of Mg, and according to my students, impossible to identify on a lab quiz.
Calcium is the fifth-most abundant element in the Earth’s crust, and features in feldspar, gypsum and limestone to name a few very prominent rocks and minerals. Calcium is more reactive than Mg, and less used in alloys. It is, however, widely used in fertilizers, as Ca is another essential plant nutrient. It’s also extremely important to human health, and widely taken as a supplementary vitamin.
This does contain Ca, but you're probably better off with a glass of milk.
The elements get a little more arcane as we move down the group. Strontium has some uses in metal refining and electronics manufacture. It also has a radioactive isotope called Sr-90 that gives the whole element a bad reputation. Strontium behaves similarly to Ca, and when ingested by humans, makes its way into the bones in the same way that Ca does. The effect is particularly pronounced people with very low Ca-intake; Sr is an imperfect substitute for Ca, but the body is willing to try. This only becomes really problematic when the Sr in question is radioactive. Sr-90 isotopes are one of the most readily incoporated radioactive isotopes out there, and thus a serious health concern. Barium is another poisonous alkali earth metal, although certain barium compounds are non-toxic, and are commonly used in x-rays of the intestinal tract. Barium compounds do not absorb x-rays, and if ingested prior to the procedure, have a highlighting effect on the tract. Other Ba compounds are used in the manufacture of pigments, fireworks (Ba is responsible for that great green glow), glass manufacture and resin manufacture. Certain naturally occurring Ba compounds glow in the dark, and were of great interest to early chemists.
And then there is radium (http://education.jlab.org/itselemental/ele088.html), alternately called Marie Curie’s doctoral thesis. These days, there aren’t as many uses for the stuff, but before the dangers of radioactivity were well understood, Ra was used in early forms of chemotherapy, and less potentially beneficial forms of radiation therapy. The unearthly glow of Ra was prized, and used for painting watches (seriously). Just how radioactive is Ra? Marie Curie did her pioneering work with Ra from 1897 to 1903. Her notebooks are still too radioactive to handle. The intense amounts of radiation emitted by Ra mean that its reactions are different from those of the rest of Group II, but it only has two valence electrons, and gets tacked on at the end for neatness’ sake.
So that’s the short version of Group II. I am, of course, leaving a lot out, but that’s inevitable in science. There’s always more to learn...
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