50 | The History and Use of Our Earth’s Chemical Elements Solutions and powders of several lithium salts are very toxic to the human nervous system,
50 | The History and Use of Our Earth’s Chemical Elements Solutions and powders of several lithium salts are very toxic to the human nervous system,
thus requiring close observation by a physician when used as antidepressant drugs. SODIUM
SYMBOL:Na PERIOD:3 GROUP:1(IA) ATOMICNO:11 ATOMICMASS:22.9898amu VALENCE:1 OXIDATIONSTATE:+1
NATURALSTATE:Solid ORIGINOFNAME: TheLatinnameforthesymbolfor“sodium”(Na)isnatrium,andthe name“sodium”inLatinis sodanum,whichwasknownasanancientheadacheremedy andwascalled“soda”inEnglish.
ISOTOPES:Sodiumhas14isotopes.Theonlystableisotopeofsodiumhasanaverage atomicweightof23( 23 Na)andmakesupabout100%ofalltheisotopesoftheele- mentsodiumfoundonEarth.Alltheother13isotopes(from 19 Nato 31 Na)areradioac- tivewithrelativelyshorthalf-livesandthusareunstable.
ELECTRONCONFIGURATION EnergyLevels/Shells/Electrons Orbitals/Electrons
s2,p6
3-M=1
s1
Properties Sodium is a soft, wax-like silver metal that oxidizes in air. Its density is 0.9674 g/cm 3 , and
therefore it floats on water as it reacts with the water releasing hydrogen. It has a rather low melting point (97.6°C) and a boiling point of 883°C. Sodium is an excellent conductor of heat and electricity. It looks much like aluminum but is much softer and can be cut with a knife like butter. Its oxidation state is +1.
Characteristics On the periodic table sodium is located between lithium and potassium. A fresh cut into
sodium looks silvery but turns gray as sodium oxidizes rapidly in air, forming sodium oxide on its surface.
Sodium is extremely reactive. It reacts explosively in water as it releases hydrogen from the water with enough heat to ignite the hydrogen. The resulting compound of this reaction is sodium hydroxide (2Na + 2H 2 O → 2NaOH + H 2 ↑). Due to its extremely electropositive reactivity, there are few uses for the pure metallic form of sodium. Because of its reactivity, hundreds of sodium compounds are found on the Earth’s surface.
Guide to the Elements | 51 An unusual characteristic of several alkali metals is that a mixture of two or more has a
lower melting point than the melting point of the separate metals. This is referred to as a eutectic system of metallic alloys. For instance, sodium has a melting point of 97.6°C, and potassium’s melting point is 63.25°C, but when the two are mixed, the eutectic melting point (turning into a liquid phase) of the combined Na-K system is below zero degrees Celsius (–10°C). If cesium metal (melting point of 38.89°C) is added to the Na and K mixture, the melting point of this eutectic alloy (Na-K-Cs) is the lowest of any eutectic alloy at –78°C.
AbundanceandSource Sodium is the sixth most abundant of the Earth’s elements. Since it is a highly electroposi-
tive metal and so reactive with nonmetals, it is not found in its pure elemental form on Earth. Rather, it is found in numerous compounds in relatively abundant quantities. About 2.83% of the Earth’s crust consists of sodium in compounds.
Sodium is produced by an electrolytic process, similar to the other alkali earth metals. (See figure 4.1). The difference is the electrolyte, which is molten sodium chloride (NaCl, com- mon table salt). A high temperature is required to melt the salt, allowing the sodium cations to collect at the cathode as liquid metallic sodium, while the chlorine anions are liberated as
chlorine gas at the anode: 2NaCl (salt) + electrolysis → Cl 2 ↑ (gas) + 2Na (sodium metal). The commercial electrolytic process is referred to as a Downs cell, and at temperatures over 800°C, the liquid sodium metal is drained off as it is produced at the cathode. After chlorine, sodium is the most abundant element found in solution in seawater.
History Before the year 1700, chemists were unable to distinguish the differences among the vari-
ous alkali metals. Sodium was often confused with potassium, which was artificially produced by slowly pouring water over wood ashes and then drying the resulting alkaline crystal depos- its. Some natural alkali metals were also found at the edges of dried lakebeds and mines and even on exposed surfaces of the Earth.
In the early 1700s, Henri-Louis Duhamel du Monceau (1700–1782) was the first to real- ize that many minerals exhibited similar alkaline (basic) characteristics. He studied samples of salts both derived artificially and found in nature, including saltpeter (potassium nitrate used in gunpowder), table salt, Glauber’s salt, sea salt, and borax.
In 1807 Sir Humphry Davy (1778–1829) devised an electrolysis apparatus that used elec- trodes immersed in a bath of melted sodium hydroxide. When he passed an electric current through the system, metallic sodium formed at the negative (cathode) electrode. He first per- formed this experiment with molten potassium carbonate to liberate the metal potassium, and
he soon followed up with the sodium experiment. Today, sodium and some of the other alkali metals are still produced by electrolysis. The types of electrolytes may vary using a mixture of sodium chloride and calcium chloride and then further purifying the sodium metal.
CommonUses Because sodium is such a reactive element and is not found in its elemental form, it is
responsible for the formation of many compounds on the Earth’s surface. Sodium oxide (Na 2 O), also known as sodium monoxide, is the most abundant and caustic salt of sodium in the Earth’s crust, but sodium chloride (NaCl) is probably the most common and useful. Other