Compounds containing sodium and potassium are very common. When they react with water, they form strong bases, called alkalines. The word ‘alkali’ comes from the Arabic expression for wood ash that has been heated in the absence of oxygen to remove impurities. So the inspiration for the name of the entire alkali metal group is this simple waste product at the end of a wood fire—ashes.
Alkali Metals React Easily
Some metals can and very often do appear in nature as pure metals. And, there are other metals that can be fairly easily isolated from their oxides by heating them in a coal furnace. But group 1 metals like lithium, sodium, and potassium are more stubborn when it comes to taking their metallic form.
Removal of that outermost electron gives an alkali metal an octet, but also creates a positively charged ion that strongly prefers to hide in compounds with other elements. Chemists measure how difficult it is to remove an electron in terms of how much electrochemical energy must be put into a system to pull the outermost electron away from the neutral metal atom.
And for the alkali metals, it’s very easy: this group accounts for six of the eight lowest first-ionization energies of all elements. What this means is that nonmetals like oxygen, for example, can very easily react with alkali metals to form ionic metal-oxide materials—and the ionic bonds that they form are very strong. They are so strong that even the usual methods of smelting could never coax their metallic forms from their oxides or from the touch of aluminosilicate minerals in which they appear in great abundance.
Alkali metals also take the form of ions when they are inside living things. Sodium and potassium, for example, are essential nutrients for humans and plants alike. But again, that sodium and potassium is not actually metallic. Rather it takes the form of dissolved ions that the body uses to perform needed biological functions.
So alkali metals are very common in the ground and even in living things, but their proximity to the noble-gas configuration makes them particularly content to give up an electron and form ions which conceals them from easy observation. So how do sodium and potassium get into the ashes of a wood fire? Fire is a chemical reaction that adds oxygen to fuel, releasing useful heat and light in the process. But the byproduct of this process is a fuel that has been oxidized.
Now, plants like trees contain lots of complex materials, like DNA, cellulose, and proteins. These compounds are mostly made of nonmetals like carbon, oxygen, sulfur, phosphorous, hydrogen, and nitrogen. But like us, plants also rely on certain metals, like potassium and sodium ions, to help manage their biochemistry. When the nonmetals in a burning log react with oxygen, they tend to form gases.
But metals behave differently. Because of their position on the far left of the periodic table, they have very low electronegativities, and when they combine with oxygen, they form ionic compounds. These ionic compounds remain solids. So, the act of burning a log helps to isolate metal oxides, which remain solid, from nonmetal oxides, which float away into the atmosphere. Only the metal oxides collect in the ashes at the base of the fire, where they can be collected.
This article comes directly from content in the video series Understanding the Periodic Table. Watch it now, on Wondrium.
Group 1 Metals Get Trendy
In larger alkali metals, screening of the nucleus from valance electrons by interior electrons makes it ‘easier’ for valence electrons to be released to any reagent that will take it. This trend is born out if we look at the first ionization potential of each alkali metal. With each step down the table, we add a new layer of screening electrons, making the valence electron easier and easier to pull away.
The most common way teachers demonstrate this trend is through the reaction of alkali metals with water. This produces alkali metal hydroxides, with hydrogen gas as a byproduct. Lithium, sodium, and potassium create progressively more vigorous reactions because each loses an electron from a progressively larger energy level in the reaction.
The vigorousness with which heat and hydrogen gas are produced increases, with lithium producing little more than a fizzle, sodium dancing and popping more aggressively, and potassium often bursting into flames. Rubidium and cesium react even more aggressively. Their explosions may sometimes appear less spectacular, but this is only because the metal itself is ripped apart in the reaction so far that no significant hydrogen can accumulate!
How Do We Obtain These Metals?
So, we have established that all alkali metals are too reactive to occur as pure metals in nature. However, lithium, sodium, and potassium metals can be easily produced by electrolysis and stored under mineral oils that protect them from oxygen and water. Removing the oil or creating a freshly exposed metal surface by cutting the sample causes the exposed metal to quickly tarnish as it reacts with oxygen and moisture in the air itself.
Rubidium and cesium, however, require even more careful handling. Samples of these elements are often stored in sealed ampules to be certain that no oxygen or moisture whatsoever can reach the sample. Cesium is sometimes even stored as cesium azide, a compound of cesium and nitrogen that can be heated to release the metal when it is needed, avoiding the tricky task of storing the perilously reactive metal itself.
Common Questions about the Fascinating Qualities of Alkali Metals
The term ’alkali’ is an Arabic expression that refers to wood ash that has been heated in the absence of oxygen to get rid of impurities.
Alkali metals such as sodium, potassium, and lithium are stubborn in taking their metallic form since the removal of the outermost electron leads to the creation of a positively charged ion that strongly prefers to hide in other compounds. And since taking away the outermost electron in these metals is easier than others, they normally create bonds with other elements such as oxygen.
Alkali metals such as lithium, sodium, and potassium are stored under mineral oils to avoid them interacting with oxygen or water. On the other hand, rubidium and cesium have to be stored in sealed ampules just to be safe. Sometimes cesium is stored as cesium azide, which can easily be heated to obtain the cesium when needed.