|Boiling Point: 1156°K, 883°C, 1621°F
Melting Point: 371°K, 98°C, 208°F
Electrons Energy Level: 2, 8, 1
Isotopes: 16 + 1 Stable
Heat of Vaporization: 96.96 kJ/mol
Heat of Fusion: 2.598 kJ/mol
Density: 0.971 g/cm3 @ 300°K
Specific Heat: 1.23 J/g°K
Atomic Radius: 2.23Å
Ionic Radius: 1.02Å
Electronegativity: 0.93 (Pauling); 1.01 (Allrod Rochow)
Vapor Pressure: 0.0000143 Pa @ 961°C
|Sodium is perhaps
the most characteristic alkali metal, reacting violently with water and rapidly with the
oxygen in air. It symbol (Na) comes from its Latin name, Natria, whereas its
English name is taken from soda which contains it.
Sodium was discovered and isolated in 1807 by Sir Humphry Davy. In its pure form it is silvery white and soft enough to cut with a knife. It is the sixth most abundant element in the earth's crust, occurring in large amounts in both (sea)water and soil in various mineral compounds, the most common of which is sodium chloride.
The metal melts below the boiling point of water (97oC). In its elemental state it has been used as a molten coolant in nuclear reactors and is currently under research for sodium/sulfur batteries. Its most common compounds for industrial use include sodium chloride, sodium hydroxide (lye), sodium carbonate (washing soda) and sodium sulfate.
Sodium is a chemical element which has the symbol Na (Latin: natrium), atomic number 11, atomic mass 22.9898 g/mol, oxidations number +1. Sodium is a soft, silvery, highly reactive element and is a member of the alkali metals within "group 1" (formerly known as group IA). It has only one stable isotope, 23Na. Sodium was first isolated by Sir Humphry Davy in 1807 by passing an electric current through molten Sodium Hydroxide, NaOH. Sodium quickly oxidizes in air so it must be stored in an inert environment such as kerosene. Sodium is present in great quantities in the earth's oceans as Sodium Chloride, NaCl. It is also a component of many earthly minerals, and it is an essential element for animal life.
1s2 2s2p6 3s1
Sodium (English, soda) has long been recognized in compounds, but was not isolated until 1807 by Sir Humphry Davy through the electrolysis of caustic soda (NaOH). In medieval Europe a compound of sodium with the Latin name of sodanum was used as a headache remedy. Sodium's symbol, Na, comes from the neo-Latin name for a common sodium compound named natrium, which comes from the Greek nítron, a natural mineral salt whose primary ingredient is hydrated sodium carbonate. The difference between the English name soda and the abbreviation Na comes from Berzelius' publication of his system of atomic symbols in Thomas Thomson's Annals of Philosophy.
The flame test for Sodium displays a brilliantly bright yellow emission due to the so called "Sodium D-lines" at 588.9950 and 589.5924 nanometers. As early as 1860, Kirchhoff and Bunsen noted the high sensitivity that a flame test for Sodium could give. They state in Annalen der Physik und der Chemie in the paper "Chemical Analysis by Observation of Spectra":
In a corner of our 60 cu.m. room farthest away from the apparatus, we exploded 3 mg. of sodium chlorate with milk sugar while observing the nonluminous flame before the slit. After a while, it glowed a bright yellow and showed a strong sodium line that disappeared only after 10 minutes. From the weight of the sodium salt and the volume of air in the room, we easily calculate that one part by weight of air could not contain more than 1/20 millionth weight of sodium.
Sodium is present in fair abundance in the sun and stars. The D lines of sodium are among the most prominent in the solar spectrum. Sodium is the fourth most abundant element on earth, comprising about 2.6% of the earth's crust; it is the most abundant of the alkali group of metals of which it is a member. The most common compound is sodium chloride, but it occurs in many other minerals, such as soda niter, cryolite, amphibole, zeolite, etc. It is a very reactive element and is never found free in nature. It is now obtained commercially by the electrolysis of absolutely dry fused sodium chloride. This method is much cheaper than that of electrolyzing sodium hydroxide, as was used several years ago. Sodium is a soft, bright, silvery metal which floats on water, decomposing it with the evolution of hydrogen and the formation of the hydroxide.
Although sodium is the sixth most abundant element on earth and comprises about 2.6% of the earth's crust, it is a very reactive element and is never found free in nature.
Sodium is used in the production of titanium, sodamide, sodium cyanide, sodium peroxide, and sodium hydride. Liquid sodium has been used as a coolant for nuclear reactors. Sodium vapor is used in streetlights and produces a brilliant yellow light.
A FASOR used at the Starfire Optical Range for LIDAR and laser guide star experiments is tuned to the Sodium D2a line and used to excite Sodium atoms in the upper atmosphere. FASOR stands for Frequency Addition Source of Optical Radiation, and for this system it is two single mode and single frequency solid state IR lasers, 1.064 and 1.319 microns, that are frequency summed in a LBO crystal within a doubly resonant cavity.
Sodium is relatively abundant in stars and the D spectral lines of this element are among the most prominent in star light. Sodium makes up about 2.6% by weight of the Earth's crust making it the fourth most abundant element overall and the most abundant alkali metal.
At the end of the 19th century, sodium was chemically prepared by heating Sodium Carbonate with Carbon to 110°C.
Na2CO3 (liquid) + 2C (solid, coke) 2Na (vapor) + 3CO (gas).
It is now produced commercially through the electrolysis of liquid Sodium Chloride, NaCl. This is done in a Downs Cell in which the NaCl is mixed with Calcium Chloride, CaCl2 to lower the melting point below 700°C. As Calcium is more electropositive than Sodium, no Calcium will be formed at the cathode. This method is less expensive than the previous Castner Process of electrolyzing Sodium Hydroxide, NaOH.
Very pure Sodium can be isolated by thermal decomposition Sodium Azide.
Metallic sodium costs about 15 to 20 cents per pound ($0.30/kg to $0.45/kg) in 1997 but reagent grade (ACS) sodium cost about $35 per pound ($75/kg) in 1990.
Phase Behavior Under Pressure
Under extreme pressure, Sodium departs from common melting behavior. Most materials require higher temperatures to melt under pressure than they do at normal atmospheric pressure. This is because they expand on melting due to looser molecular packing in the liquid, and thus pressure forces equilibrium in the direction of the denser solid phase.
At a pressure of 30 gigapascals (300,000 times sea level atmospheric pressure), the melting temperature of Sodium begins to drop. At around 100 gigapascals, Sodium will melt at near room temperature. A possible explanation for the aberrant behavior of Sodium is that this element has one free electron that is pushed closer to the other 10 electrons when placed under pressure, forcing interactions that are not normally present. While under pressure, solid Sodium assumes several odd crystal structures suggesting that the liquid might have unusual properties such as superconduction or superfluidity.
Sodium in its metallic form can be used to refine some reactive metals, such as Zirconium and Potassium, from their compounds. This alkali metal as the Na+ ion is vital to animal life. Other uses:
Sodium Chloride or Halite, better known as common salt, is the most common compound of Sodium, but Sodium occurs in many other minerals, such as Amphibole, Cryolite, Soda Niter and Zeolite. Sodium compounds are important to the chemical, glass, metal, paper, petroleum, soap, and textile industries. Hard soaps are generally Sodium salt of certain fatty acids (Potassium produces softer or liquid soaps).
The Sodium compounds that are the most important to industry are common salt (NaCl), Soda Ash (Na2CO3), Baking Soda (NaHCO3), Caustic Soda (NaOH), Chile Saltpeter (NaNO3), Di- and Tri-Sodium Phosphates, Sodium Thiosulfate (hypo, Na2S2O3 · 5H2O), and Borax (Na2B4O7 · 10H2O).
|Soda Ash, Na2CO3||Baking Soda, NaHCO3|
|Caustic Soda, NaOH||Chile Saltpeter, NaNO3|
|Borax, Na2B4O7 · 10H2O||Sodium Thiosulfate, Na2S2O3 · 5H2O|
|Halite, Sodium Chloride, NaCl||Pyromet, NaCl/(NH4)2HPO4|
There are thirteen isotopes of Sodium that have been recognized. The only stable isotope is 23Na. Sodium has two radioactive cosmogenic isotopes (22Na, half-life = 2.605 years; and 24Na, half-life 15 hours).
Acute neutron radiation exposure (e.g., from a nuclear criticality accident) converts some of the stable 23Na in human blood plasma to 24Na. By measuring the concentration of this isotope, the neutron radiation dosage to the victim can be computed.
|Na19||19.0139||< 40 ns|
|Extreme care is required in handling elemental/metallic Sodium. Sodium is potentially explosive in water (depending on quantity) and is a caustic poison, since it is rapidly converted to Sodium Hydroxide on contact with moisture. The powdered form may combust spontaneously in air or Oxygen.|
Sodium must be stored either in an inert (Oxygen and moisture free) atmosphere (such as Nitrogen or Argon), or under a liquid hydrocarbon such as mineral oil or kerosene.
The reaction of Sodium and water is a familiar one in chemistry labs, and is reasonably safe if amounts of Sodium smaller than a pencil eraser are used and the reaction is done behind a plastic shield by people wearing eye protection. However, the Sodium-water reaction does not scale up well, and is treacherous when larger amounts of Sodium are used. Larger pieces of Sodium melt under the heat of the reaction, and the molten ball of metal is buoyed up by Hydrogen and may appear to be stably reacting with water, until splashing covers more of the reaction mass, causing thermal runaway and an explosion which scatters molten Sodium, Lye solution, and sometimes flame. This behavior is unpredictable, and among the alkali metals it is usually Sodium which invites this surprise phenomenon, because Lithium is not reactive enough to do it, and Potassium is so reactive that chemistry students are not tempted to try the reaction with larger Potassium pieces.
Sodium is much more reactive than Magnesium; a reactivity which can be further enhanced due to Sodium's much lower melting point. When Sodium catches fire in air (as opposed to just the Hydrogen gas generated from water by means of its reaction with Sodium) it more easily produces temperatures high enough to melt the Sodium, exposing more of its surface to the air and spreading the fire.
Few common fire extinguishers work on Sodium fires. Water, of course, exacerbates Sodium fires, as do water-based foams. CO2 and Halon are often ineffective on Sodium fires, which reignite when the extinguisher dissipates. Among the very few materials effective on a Sodium fire are Pyromet and Met-L-X. Pyromet is a NaCl/(NH4)2HPO4 mix, with flow/anti-clump agents. It smothers the fire, drains away heat, and melts to form an impermeable crust. This is the standard dry-powder canister fire extinguisher for all classes of fires. Met-L-X is mostly Sodium Chloride, NaCl, with approximately 5% Saran plastic as a crust-former, and flow/anti-clumping agents. It is most commonly hand-applied, with a scoop. Other extreme fire extinguishing materials include Lith-X, a graphite based dry powder with an organophosphate flame retardant; and Na-X, a Na2CO3 - based material.
Because of the reaction scale problems discussed above, disposing of large quantities of Sodium (more than 10 to 100 grams) must be done through a licensed hazardous materials disposer. Smaller quantities may be broken up and neutralized carefully with Ethanol (which has a much slower reaction than water), or even Methanol (where the reaction is more rapid than Ethanol's but still less than in water), but care should nevertheless be taken, as the caustic products from the Ethanol or Methanol reaction are just as hazardous to eyes and skin as those from water. After the alcohol reaction appears complete, and all pieces of reaction debris have been broken up or dissolved, a mixture of Alcohol and water, then pure water, may then be carefully used for a final cleaning. This should be allowed to stand a few minutes until the reaction products are diluted more thoroughly and flushed down the drain. The purpose of the final water soak and wash of any reaction mass which may contain Sodium is to ensure that Alcohol does not carry unreacted Sodium into the sink trap, where a water reaction may generate Hydrogen in the trap space which can then be potentially ignited, causing a confined sink trap explosion.
Physiology and Sodium Ions
Sodium ions play a diverse and important role in many physiological processes. Excitable animal cells, for example, rely on the entry of Na+ to cause a depolarization An example of this is signal transduction in the human central nervous system, which depends on Sodium ion motion across the nerve cell membrane, in all nerves.
|Some potent neurotoxins, such as batrachotoxin, increase the Sodium ion permeability of the cell membranes in nerves and muscles, causing a massive and irreversible depolarization of the membranes, with potentially fatal consequences.|
However, drugs with smaller effects on sodium ion motion in nerves may have diverse pharmacological effects which range from anti-depressant to anti-seizure actions.
Sodium is the primary cation (positive ion) in extracellular fluids in animals and humans. These fluids, such as blood plasma and extracellular fluids in other tissues, bathe cells and carry out transport functions for nutrients and wastes. Sodium is also the principal cation in seawater, although the concentration there is about 3.8 times what it is normally in extracellular body fluids. This suggests that animal life moved from the sea to dry land at a time when the seas were far less salty than they are now.
Although the system for maintaining optimal salt and water balance in the body is a complex one, one of the primary ways in which the human body keeps track of loss of body water is that osmoreceptors in the hypothalamus sense a balance of sodium and water concentration in extracellular fluids. Relative loss of body water will cause sodium concentration to rise higher than normal, a condition known as hypernatremia. This ordinarily results in thirst. Conversely, an excess of body water caused by drinking will result in too little sodium in the blood (hyponatremia), a condition which is again sensed by the hypothalamus, causing a decrease in vasopressin hormone secretion from the posterior pituitary, and a consequent loss of water in the urine, which acts to restore blood Sodium concentrations to normal.
Severely dehydrated persons, such as people rescued from ocean or desert survival situations, usually have very high blood Sodium concentrations. These must be very carefully and slowly returned to normal, since too-rapid correction of hypernatremia may result in brain damage from cellular swelling, as water moves suddenly into cells with high osmolar content.
Because the hypothalamus/osmoreceptor system ordinarily works well to cause drinking or urination to restore the body's Sodium concentrations to normal, this system can be used in medical treatment to regulate the body's total fluid content, by first controlling the body's Sodium content. Thus, when a powerful diuretic drug is given which causes the kidneys to excrete Sodium, the effect is accompanied by an excretion of body water (water loss accompanies Sodium loss). This happens because the kidney is unable to efficiently retain water while excreting large amounts of Sodium. In addition, after Sodium excretion, the osmoreceptor system may sense lowered Sodium concentration in the blood, and then directs compensatory urinary loss of water, in order to correct the hyponatremia, or (low-blood-Sodium) state.
Atomic Radius (Å): 2.23Å
Electrochemical Equivalents: 0.85775 g/amp-hr
Atomic Mass Average: 22.98977