Now there’s a heading that’s really funny to those familiar with roman numerals and the periodic table. Nothing like an incredibly obscure joke to get us started on an in-depth exploration of the periodic table! Don’t worry, we’re starting off with a bang today.
Yes, today we’re going to look at Group I of the periodic table, also known as the alkali metals. Long beloved by chemistry students and irresponsible chemistry instructors for their highly reactive properties, one of the more distinctive features of this group is that all the elements explode when they come in contact with water. The explosions grow stronger as you move down the group from lithium (atomic number 3) to cesium (atomic number 55). For those of you who never got to play with sodium and water, or have mused about what would happen if you introduced cesium to water, this video is both entertaining and informative. And for the record, that “francium bomb” video floating around the internet is a hoax. If francium were actually available, I’m sure someone would try to weaponize the stuff, but as it happens, francium is both extremely rare and extremely unstable .
To get into why all the Group I metals are so reactive, we’re going to backtrack a little. Group I refers to the elements in the first column of the periodic table - lithium, sodium, potassium, cesium, rubidium and francium. As they’re all in the same group, we can assume that all these elements undergo chemical reactions in the same fashion - they are more likely to undergo some kinds of reactions, and less likely to undergo others. Hydrogen is placed on top of lithium in most table arrangements, but it isn’t actually part of the group. Group 1 is distinguished at a fundamental level by consisting of elements whose non-charged atoms have only one atom in the outermost valence shell. Remember those from last week? But wait, hydrogen only has one electron in its outermost shell. What makes hydrogen so different? Well, in the case of hydrogen, the outermost shell of electrons is the only shell of electrons. Unlike any other element with an unfilled outermost shell, there is no shell of electrons behind that partly empty shell. This leads the positively charged nucleus to have a much stronger draw on the lone electron - hydrogen might still lose its outermost shell to become a hydrogen ion, but it might also share that lone electron with another atom in a covalent bond (remember those from last week?). Also, the first shell to fill in any atom only holds two electrons - hydrogen could go either way. For comparison, the next shell to fill holds eight electrons.The alkali metals have a full shell of electrons behind the lone electron, and can deal with that unfilled shell by either losing one electron or gaining seven. Losing one electron is easier, so in a situation where there are atoms which attract electrons, an alkali metal is going to shed that one electron. And that shedding of an electron is the cause of alkali metal reactivity.
Ok, but how does that start a fire in water? Time for a quick tangent about the chemistry of water. Water is a bizarre substance. By virtue of its atomic weight alone, it should only exist at room temperature as a gas, not as a liquid - none of the other elements in the surrounding periods exist as liquids at room temperature. Some elements, like carbon or aluminum, bond many atoms together, to form solids. Other elements form relatively small molecules with other elements, and exist as gases. However, water experiences something called hydrogen bonding. Last week, when I mentioned covalent bonds, I described it as when atoms essentially merge their outermost shells and combine electrons. Well, it’s not a perfectly even merge. Frequently, one of the atoms involved is going to exert a stronger pull on the electrons, creating a molecule that is more positive at one end, and more negative at the other end. In the case of water molecules, the effect is particularly strong - the end of the molecule with the two hydrogen atoms has a relatively strong positive charge, the end with the oxygen atom has a relatively strong negative charge (water molecules have a sort of boomerang shape to them). Neither charge is anywhere near as strong as a true ionic charge (i.e., hydrogen with zero electrons instead of one), but the charge is strong enough to attract the positive end of one water molecule to the negative end of another, and on and on. The bonds are weak, but they allow water to remain liquid at room temperature.
Normally at room temperature, water bubbles along in its happy liquid state. Changes in temperature may either negate (by heating) or strengthen (by freezing) the hydrogen bonds, leading to changes in the phase state of water, but the molecules are stable, with each oxygen strongly attached to its two hydrogens. Alkali metals, however, are so reactive that they are capable of breaking water molecules. An oxygen atom is more strongly attracted to an alkali metal atom than it is to a hydrogen atom - if the oxygen atom is already bonded to hydrogen atoms (giving it a full outermost shell), it will break the bond with one of those hydrogen atoms, and make a new bond with one of the alkali metal atoms. Breaking any chemical bond releases energy in the form of heat. The hydrogen atoms released from the water molecule combine into hydrogen gas, which is highly flammable. The bigger the alkali metal atom, the stronger the reaction. I think you can see where this is going.
Ok, impossible to test, but a sound theory.
Cool as the reaction is, it's unusual (outside of chemistry classes). Normally, alkali metals exist as positively charged ions. They commonly combine with negatively charged ions - one such combination you might be familiar with is NaCl, aka table salt. Some of the alkali metals, in stable ionic form (i.e. Na+), play a role in biology - lithium functions as a neurotransmitter (and is administered in the treatment of certain forms of depression), while potassium and sodium are both key to proper cellular functioning. There aren’t as many non-explosive applications for applications for cesium and rubidium, but they are found in compounds in certain technical applications.
The one place you don’t generally find the alkali metals is in a form that most would recognize as a metal. Oh, it is possible to isolate alkali metals into a solid form - the result is a soft, silvery grey substance.
However, that substance reacts spontaneously with water, and in some cases, air. Chunks of solid alkali metals must be stored in a non-water liquid, like kerosene. Hence, you’re not going to see it much out of the lab. Alkali metal compounds and ions, on the other hand, are all around you, and even in you. When they aren’t exploding, Group I is a pretty useful collection of elements.
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