Sodium is always a crowd pleaser. As an alkali (Group I) metal, it reacts easily to lose an electron. In fact, it reacts so readily with other materials that it is never found in the pure state in nature. Most people are familiar with common sodium compounds, perhaps without always realizing it. Table salt (sodium chloride) is known to many of course. Then there’s mono-sodium glutamate, traditionally encountered in Chinese food, at least in America. Bleach is sodium hypochlorite, and Lye is a solution of sodium hydroxide in water. Reading the labels on medicines, shampoos, and foods will reveal that sodium is found in mixtures galore.
Sodium is a metal, as mentioned, and does have a metallic silvery color. Rarely is this observed though, because it quickly reacts with water from the air and gains a white-gray coating as a result. When pure sodium is stored, it is usually bottled with mineral oil, which helps to keep the water away. It can be cut while beneath the surface of the oil (sodium is a very soft metal, it can be cut easily with a spoon!), temporarily revealing its true color. Even the oil doesn’t keep all the moisture away though, and the dull coloring soon returns.
Sodium’s reactivity with water is the reason that most people remember it. When sodium and water come into contact, they react to form sodium hydroxide, hydrogen gas, and a lot of heat. If a small piece of sodium is used, the hydrogen gas will cause the sodium to dance across the surface of the water, atop a bed of hydrogen bubbles. A larger piece generates more hydrogen, more heat, and has the potential to make the sodium jump, as pockets of hydrogen ignite from the heat. The satisfying “pop” gets a lot of attention. Even bigger pieces of sodium will explode when dunked in water, and are quite dangerous to anyone nearby.
For demonstrations, I found a safer way to demonstrate sodium’s reaction with water at Flinn Scientific. A tall cylinder is filled halfway with water, and a few drops of phenolphthalein added. Then the cylinder is filled the rest of the way with mineral oil. A piece of sodium is dropped in. The sodium sins through the oil until it reaches the water. At the water’s surface, it reacts, causing the formation of sodium hydroxide (which is basic) in the water, turning it purple (because of the phenolphthalein) and hydrogen gas bubbles that float upwards, carrying the sodium with them. The hydrogen cannot ignite, because there is no oxygen present to support combustion. As the bubbles escape, the sodium sinks again, and repeats the cycle, dancing up and down, ever shrinking as it reacts little by little.
Sodium will even burn, at high enough temperatures. Once ignited, a sodium fire is very dangerous. It burns with an intense heat and very bright light. Most hazardous though, is the attempt to extinguish it with water. Unthinkingly or unknowingly, a person may dump water on the flames. The burning sodium then explodes, sending flaming hot bits of sodium in all directions. (Use an extinguisher that removes oxygen from the flame, or better yet, pile sand on the metal instead. It wouldn’t be helpful if your carbon dioxide extinguisher blew the burning sodium across the room.)
A few last bits of trivia about sodium:
It is atomic number 11 (11 protons, 12 neutrons, and 11 electrons – until it reacts, after which the sodium ion has 10 electrons), it has an atomic mass of 22.99 amu, and has the chemical symbol “Na”, which stands for Natrium, its classical name.
Think sodium is too tame? Upgrade to potassium.
