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Sodium, element number 11, is one of the most common elements in the crust of the earth.  Except for school laboratory demonstrations, few people have ever seen elemental (metallic) sodium because it is so reactive and actually has very limited consumer uses (that would be about zero consumer uses).

We have hinted at the concept of periodicity previously, like last week when we saw how similar the chemical behavior of helium and neon are.  The similarities betwixt hydrogen and lithium are much less marked than those betwixt lithium and sodium, mostly due to the extreme low mass of hydrogen, making quantum effects more pronounced.  Thus, sodium is the second alkali metal after lithium even though hydrogen is in the same column in the periodic table.

In other words, the two first row elements, hydrogen and helium, are aberrant because of their low masses AND because they have only the K electron shell in the ground state AND as a corollary, only the 1s orbital that is filled with only two electrons.  Starting with the second row, the L shell begins to be filled and it contains, in addition to the 1s orbital, a 2s and three 2p orbitals.  Row three elements, sodium being the first of which, also contain in addition to those orbitals, a 3s and three 3p orbitals, making them more like the second row than the second is to the first row.

Sodium is common in the cosmos, being formed easily in massive stars by fusing two 12C nuclei into a 23Na nucleus and releasing a proton.  This process releases energy, although not as much as when 20Ne and an alpha particle are formed.

The reason that I say that you can not get away from it is that sodium is so common that almost everything is contaminated with at least trace amounts of it.  Normally this is not a problem since it is nontoxic, but for specialized applications sodium must be rigorously excluded, and that is just not easy much of the time.  The average human has about 56 grams of sodium in her or his body, that in chemical terms that is a LOT, two moles or 12.04 x 1022 ions.  So anything you touch becomes contaminated by sodium, because it is present in large amounts in sweat.

There is only one stable isotope of sodium, 23Na.  This is not unusual, because we have seen in the past that nuclei with odd numbers of protons very often have very few, if any, stable nuclei in comparison with nuclei with and even Z.  Other isotopes exist, the longest lived one 24Na, with a half life of around 15 hours.  It has no uses except for one, and that is very obscure.  It turns out that 23Na has a small but significant neutron absorption of 0.53 barns, being transformed to 24Na.  That decays to the stable 24Mg by emission of an electron and two gamma photons.  When people are exposed to thermal neutrons in an accident, such as containment failure in a nuclear reactor, the amount of 24Na is formed at a known quantity, and since both the half life of the isotope and the time of exposure are known, the total dose of neutrons can be determined by some rather simple instrumental methods, allowing medical types to determine how aggressively to treat the radiation poisoning.

Sodium is the term given to the element from some of its compounds known since antiquity, probably going all the way back to the Arabic suda, crude sodium carbonate.  The elemental symbol is taken from the post classical Latin (i. e., "made up Latin") natrium, derived from the Egyptian natron, another crude sodium carbonate, that was used in the mummification process.  I must mention that the symbol was conceived by one Jöns Jakob Berzelius, the brilliant Swedish chemist of times past and to who I can trace my Ph.D. legacy.  I must contact my major professor and get the entire genealogy.

It is of little use to look at molecular orbital diagrams for sodium since in the metallic state it is a metallic solid in which electrons are not bound to any given set of atoms, but delocalized throughout the entire mass of material, of at least within single crystals in the mass.  However, in the gas phase sodium form diatomic molecules, like lithium and hydrogen, with a single bond.

Sodium is really a pretty thing.  Freshly cut, it is very silvery but tarnishes rapidly and becomes dull and sort of yellowish.  In my professional life I have used lots of metallic sodium, and it is one of my favorite things.  Put a tiny piece in water, and it floats (its density is about 3% less than water), and then melts (mp is about 98 degrees C) and floats even higher (the liquid density is only about 0.93 that of water).  But it does more than float!  It reacts with water to form hydrogen gas and sodium hydroxide, which we shall discuss in detail next time, and the hydrogen bubbles lift it even higher.  Small pieces just dance around for a while, sometimes becoming alit when the hydrogen ignites, with a brilliant yellow flame.  In a bit I shall give you a safe home experiment to try.

Put a big enough piece in water and it does what was just described, but more times than not, it will explode into a million pieces.  This is a rather hazardous demonstration, but sure gets the attention of students.  It is best to keep them back, because both metallic sodium and sodium hydroxide are quite corrosive to tissue.  I am not sure why only the larger pieces explode, but I have an hypothesis.  I think that the heat transfer from the reacting sodium to the cooling water is fast enough in small pieces that the interior of them never reaches 883 degrees C, but that in larger pieces it can, causing the whole thing to explode.  Those of you who have more experience please comment.

Compared to potassium, to be discussed another time, sodium is relatively nonreactive.  This is general trend for Group 1 (or, as I prefer Group 1A) elements to become more reactive as Z increases, just the opposite of Group 17 (my preferred Group VIIa) elements, the halogens, whose reactivity decreases with a larger Z.  You can open a half kilo sealed can of sodium in the open air and take your time to cut it into pieces of useful size (sodium is so soft that a pocket knife will go through it more easily than it will cold butter) and drop them into the mineral oil to prevent reaction with atmospheric oxygen and water without much hazard.  Trying to do the same thing with potassium almost always results in a major fire!

Let us get to the safe home experiment that you can do, but first some background.  Sodium has been used as a reference standard for light wavelengths for a long time, and the reasons are that it has an intense line in the yellow (actually, for quantum mechanical reasons, TWO lines very close together) and that sodium is universally available, so no expensive techniques are required to prepare the material.  Here is a picture of the spectrum, sort of normalized both for absolute intensity and sensitivity to the human eye.  In reality, the yellow lines are much more intense both in in absolute intensity and in perception, so they are the only ones that we see.

Insert picture.

The bottom line, to make a pun, is that the so called sodium D lines are intense, easily visualized, and constant.  Before lasers were commonly available,  the sodium D lines were used (and still are to a large extent) to decipher the optical properties of materials.  They are intense, cheap to produce, and easy to isolate.   Here is the experiment.

Sodium flame color

You need the following:

Cotton swaps, with paper connectors.  The ones on plastic stems will just melt.

Table salt, maybe a teaspoonful

Water, maybe an ounce

Colorless gas flame, like a gas range, space heater, gas grill, or turkey fryer

Dissolve the salt in the water and light up the flame.  Dip the swab into the salt water and put into the flame.  You will see a fine, bright yellow color.

Over time, we shall have a few more of these flame color experiments.

Elemental sodium has rather limited uses, as I alluded to earlier.  But those are important ones, at least on industrial scales.  Interestingly, even though sodium compounds, like salt, are known from prehistoric times, the element itself was not isolated until 1807.  The brilliant and short lived English chemist Sir Humphrey Davy was the first to make pure sodium, using the relatively new technique by which electricity was put to good application.

Davy took sodium hydroxide (actually at the time lye, a mixture of the hydroxide and carbonate of sodium) and melted it at high temperature.  Then he lowered electrodes connected to a stout battery into it, and at the negative terminal metallic sodium was produced.  Contemporaneous accounts indicate that Davy and his laboratory assistant danced and shouted around the room when the first elemental sodium was produced, much like my professor and I did when we (well, I) discovered the very first evidence of radical ion produced chemically induced dynamic nuclear polarization,CIDNP back in graduate school.

I keep talking about s orbitals and p orbitals, but have not done much to explain them, so here we go.  This is fundamental for understanding how bonding works, so let us take a stab at them.  It has to do both with geometry and number of electrons.  It turns out that s orbitals can hold only two electrons when fully filled.  The term "s" stands for "sharp", an old connexion betwixt spectroscopy and quantum theory.  It also turns out that s orbitals have a spherical symmetry.

The p orbitals (from the old spectroscopy term "principal") are not spherical and are pointed at different directions.  Each p orbital is directed either in the x, y, or z direction and the three of them can hold two electrons each, for a total of six electrons.  Here are some pictures of 2s and 2p orbitals.  They are pretty much to relative scale.  Note that the s orbitals are of much lower volume than the p orbitals.

This is the small, spherical 2s orbital

This the the larger, spatially directed 2px orbital.  The different colors represent the two different phases of the orbital.  Only orbitals with similar phases can form bonds.

This is the 2py orbital, identical to the 2px one except for direction

Finally, this is the 2pz orbital, identical to the other 2p orbitals except for direction

Remember back when I mentioned sigma and pi bonds in molecular orbital diagrams?  Now we are ready to define what those are.  In words, a sigma bond is a strong bond where electron density is built up DIRECTLY betwixt two nuclei.  This is possible only with two s orbitals or two px ones, viz.:


Sigma bonds are the strongest because negatively charged electron density is DIRECTLY betwixt the two positive nuclei.  The other possible bonds, pi bonds, occur when electron density is built up IN THE PLANE, but not directly betwixt of the two nuclei.  They are not as strong as sigma bonds because there is no negative electron density DIRECTLY betwixt the two positive nuclei, viz.:


This picture would have been better if the two p orbitals had different colored lobes to show that only orbital lobes of similar phase can bond.

It is possible to form a triple bond, like in nitrogen, by forming a sigma bond betwixt the two 2px atomic orbitals, one from each nitrogen atom, and two pi bonds betwixt the two 2py and the two 2pz orbitals from each nitrogen.  The only difference is that one pi bond is in the plane of the paper like the one in the picture and the other one has the electron density going behind the screen and in front of it.  That is sort of arbitrary.

The reason that p orbitals are bigger than s orbitals is that s orbitals have to accommodate only two negatively charged electrons, while p orbitals have to accommodate up to six.  Putting that much negative charge in an orbital necessarily makes those orbitals larger.

Sodium has the following electron structure:  1s22s22p63s1.  Neon, the most chemically stable element, has this structure:  1s22s22p6.  Thus, the chemistry of sodium is strongly influenced by the tendency for it to lose a single electron to attain the neon octet.  Except in the gas phase and as the metal, sodium is almost always encountered as sodium ion, with a single positive charge, Na1+.  Next time we shall get into more of the chemistry of sodium.

Elemental (metallic) sodium has a few industrial and laboratory uses, but not nearly as many uses as compounds of sodium.  Elemental sodium is used to some extent as a heat transfer medium, but suffers from the disadvantage that if the temperature of the piping falls below 98 degrees C, the sodium freezes and that completely gums up the works.  Thus, the eutectic mixture of 78% potassium and 22% sodium (NaK), with a melting point of -12.6 degrees C is often used instead.  NaK is some nasty stuff, reacting explosively with water and air, so lots of precautions have to be taken.  It is most often seen in experimental nuclear reactors because they are often shut down to change parameters or to remove specimens.  Sodium would freeze when the reactor stopped producing heat. Sodium is used in the stems of some high performance internal combustion engine valves to help keep the valve head cool.

It is also used as a starting material for several chemical processes, as a drying agent, as a reducing agent, and in street lights.  Street lights around astronomical observatories use low pressure sodium lamps since they are almost monochromatic.  The D lines can easily be filtered out and thus do not spoil the view of the night sky.  In other areas high pressure sodium lamps are used because they have a greater output and a broader spectral range.  I was once at a retreat in a community near the Palomar observatory and the egg yolk yellow streetlights were quite strange.

When I was in graduate school I used sodium to dry and deoxygenate solvents.  It only works with ethers and hydrocarbons because halogenated solvents and alcohols react with it.  I used a lot of tetrahydrofuran for synthetic work (along with diethyl ether), and to dry them I would put the solvent, after predrying with something like calcium chloride or sodium sulfate, in a still, add benzophenone and sodium metal, and let the solvent reflux under a stream of nitrogen.  After a while the solvent would turn deep purple (tetrahydrofuran) or blue (diethyl ether and benzene).  This is because sodium donates an electron to the benzophenone, forming diphenylketyl, a radical anion (one unpaired electron makes is a radical, and the negative charge makes it an anion).  The material is destroyed as soon as it forms by water and oxygen, so when the solvent turned purple or blue, it was dry and deoxygenated.

One of my former schoolteachers told me of the time that he came into ten pounds of metallic sodium and needed to get rid of some of it (even under kerosene it slowly forms the superoxide, and that is explosive).  He took about 90% of it to a lake and cut big chunks of it and used a slingshot to hurl it far out into the water where it would burn with a bright yellow flame (the same color as you get with the demo above) and then explode.  He did it at night and said that the results were spectacular!

Sodium will dissolve in liquid ammonia, forming a blue solution of "solvated electrons".  This is a powerful reducing agent, being able to reduce benzene to 1,4-cyclohexadiene.  Not many things will react with benzene to reduce it, so it is indeed a powerful reducing agent.  This is the basis of the Birch Reduction, an application of which has made possible the synthesis of human sex hormones from sterols found in wild yams from Mexico.

That will about do it for tonight.  Next time we shall look at a little of the chemistry of sodium and several of its important compounds and may look at some of its biological consequences.

Well, you have done it again!  You have wasted many more einsteins of perfectly good photons reading this alkaline piece.  And even though Mitt Romney realizes he has no chance of ever becoming president when he reads me say it, I always learn much more than I could possibly hope to teach by writing this series, so keep those comments, questions, corrections, and other feedback coming.  Tips and recs are also always quite welcome.  Remember, no science or technical issue is off topic here, so you do not have to confine your thoughts to sodium.

I shall stay around tonight as long as comments warrant, unless the opportunity arises for me to visit The Woman after The Little Girl goes down for the night.  If so, I will let you know.  I certainly will be back after The Woman's conk out time tonight, and shall return tomorrow evening around 9:00 Eastern for Review Time.

Warmest regards,

Doc, aka Dr. David W. Smith

Crossposted at

The Stars Hollow Gazette,

Docudharma, and


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