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SODIUM

Sodium (pronounced /ˈsoʊdiəm/) is an element which has the symbol Na (Latin natrium, from Arabic natrun), atomic number 11, atomic mass 23 amu , and a common oxidation number +1. Sodium is a soft, silvery white, 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. Sodium quickly oxidizes in air and is violently reactive with water, so it must be stored in an inert medium, such as kerosene or mineral oil. Sodium is present in great quantities in the earth's oceans as sodium chloride (common salt). It is also a component of many minerals, and it is an essential element for animal life. As such, it is classified as a “dietary inorganic macro-mineral.”

Characteristics

At room temperature, sodium metal is so soft that it can be easily cut with a knife. In air, the bright silvery luster of freshly exposed sodium will rapidly tarnish. The density of alkali metals generally increases with increasing atomic number, but sodium is denser than potassium.

Chemical properties

Compared with other alkali metals, sodium is generally less reactive than potassium and more reactive than lithium,[2] in accordance with "periodic law": for example, their reaction in water, chlorine gas, etc.; the reactivity of their nitrates, chlorates, perchlorates, etc.

Sodium reacts exothermically with water: small pea-sized pieces will bounce across the surface of the water until they are consumed by it, whereas large pieces will explode. While sodium reacts with water at room temperature, the sodium piece melts with the heat of the reaction to form a sphere, if the reacting sodium piece is large enough. The reaction with water produces very caustic sodium hydroxide (lye) and highly flammable hydrogen gas. These are extreme hazards (see Precautions section below). When burned in air, sodium forms sodium peroxide Na2O2, or with limited oxygen, the oxide Na2O (unlike lithium, the nitride is not formed). If burned in oxygen under pressure, sodium superoxide NaO2 will be produced. In chemistry, most sodium compounds are considered soluble but nature provides examples of many insoluble sodium compounds such as the feldspars. There are other insoluble sodium salts such as sodium bismuthate NaBiO3, sodium octamolybdate Na2Mo8O25• 4H2O, sodium thioplatinate Na4Pt3S6, sodium uranate Na2UO4. Sodium meta-antimonate's 2NaSbO3•7H2O solubility is 0.3g/L as is the pyro form Na2H2Sb2O7• H2O of this salt. Sodium metaphosphate NaPO3 has a soluble and an insoluble form

SODIUM

Sodium (IPA: /ˈsəʊdiəm/) is a chemical element which has the symbol Na (Latin: natrium), atomic number 11, atomic mass 22.9898 g/mol, common oxidation number +1. Sodium is a soft, silvery white, 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. Sodium quickly oxidizes in air and is violently reactive with water, so it must be stored in an inert medium, such as kerosene. Sodium is present in great quantities in the earth's oceans as sodium chloride. It is also a component of many minerals, and it is an essential element for animal life. As such, it is classified as a “dietary inorganic macro-mineral.”

Notable characteristics

Compared with other alkali metals, sodium is generally less reactive than potassium and more so than lithium, in accordance with "periodic law": for example, their reaction in water, chlorine gas, etc.; the reactivity of their nitrates, chlorates, perchlorates, etc. An exception to the periodic law is regarding sodium's density. The density of the elements are expected to increase down the group. However, potassium is less dense than sodium.

Owing to its high reactivity, sodium is found in nature only as a compound and never as the free element. Sodium reacts exothermically with water: small pea-sized pieces will bounce across the surface of the water until they are consumed by it, whereas large pieces will explode. While sodium reacts with water at room temperature, the sodium piece melts with the heat of the reaction to form a sphere, if the reacting sodium piece is large enough. The reaction with water produces very caustic sodium hydroxide and highly flammable hydrogen gas. These are extreme hazards (see Precautions section below). When burned in air, sodium forms sodium peroxide Na2O2, or with limited oxygen, the oxide Na2O (unlike lithium, the nitride is not formed). If burned in oxygen under pressure, sodium superoxide NaO2 will be produced.

When sodium or its compounds are introduced into a flame it will contribute a bright yellow.

In chemistry, most sodium compounds are considered soluble but nature provides examples of many insoluble sodium compounds such as the feldspars. There are other insoluble sodium salts such as sodium bismuthate NaBiO3, sodium octamolybdate Na2Mo8O25• 4H2O, sodium thioplatinate Na4Pt3S6, sodium uranate Na2UO4. Sodium meta-antimonate's 2NaSbO3•7H2O solubility is 0.3g/L as is the pyro form Na2H2Sb2O7• H2O of this salt. Sodium metaphosphate NaPO3 has a soluble and an insoluble form.[1]

Sodium ions are necessary for regulation of blood and body fluids, transmission of nerve impulses, heart activity, and certain metabolic functions. Interestingly, sodium is needed by animals, which maintain high concentrations in their blood and extracellular fluids, but the ion is not needed by plants. A completely plant-based diet, therefore, will be very low in sodium. This requires some herbivores to obtain their sodium from salt licks and other mineral sources. The animal need for sodium is probably the reason for the highly-conserved ability to taste the sodium ion as "salty." Receptors for the pure salty taste respond best to sodium, and otherwise only to a few other small monovalent cations (Li+, NH4+, and to some extent also K+). Calcium chloride also tastes somewhat salty, but also quite bitter.

The most common sodium salt, sodium chloride (table salt), used for seasoning (for example the English word "salad" refers to salt) and warm-climate food preservation, such as pickling and making jerky (the high osmotic content of salt inhibits bacterial and fungal growth). As such, salt has been an important commodity in human activities (the English word salary refers to salarium, the perquisite ("perk") given to Roman soldiers for the purpose of buying salt). The human requirement for sodium in the diet is less than 500 mg per day,[citation needed] which is typically less than a tenth as much as many diets "seasoned to taste." Most people consume far more sodium than is physiologically needed. For certain people with salt-sensitive blood pressure, this extra intake may cause a negative effect on health.

Applications
A low pressure sodium lamp, glowing with the light of sodium D spectral lines.

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:

• In certain alloys to improve their structure.
• In soap, in combination with fatty acids. Sodium soaps are harder (higher melting) soaps than potassium soaps.
• To descale metal (make its surface smooth).
• To purify molten metals.
• In sodium vapor lamps, an efficient means of producing light from electricity (see the picture), often used for street lighting in cities. Low-pressure sodium lamps give a distinctive yellow-orange light which consists primarily of the twin sodium D lines. High-pressure sodium lamps give a more natural peach-colored light, composed of wavelengths spread much more widely across the spectrum.
• As a heat transfer fluid in some types of nuclear reactors and inside the hollow valves of high-performance internal combustion engines.
• NaCl, a compound of sodium ions and chloride ions, is an important heat transfer material.
• In organic synthesis, sodium is used as a reducing agent, for example in the Birch reduction.
• In chemistry, sodium is often used either alone or with potassium in an alloy, NaK as a desiccant for drying solvents. Used with benzophenone, it forms an intense blue coloration when the solvent is dry and oxygen-free.

History
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.

Sodium (the English word for which is soda) has long been recognized in compounds, but was not isolated until 1807 by Sir Humphry Davy through the electrolysis of caustic soda. In medieval Europe a compound of sodium with the Latin name of sodanum was used as a headache remedy.

Sodium's chemical abbreviation Na was first published by Jöns Jakob Berzelius in his system of atomic symbols (Thomas Thomson's Annals of Philosophy[2]) and is a contraction of the element's new Latin name natrium which refers to natron, a natural mineral salt whose primary ingredient is hydrated sodium carbonate and which historically had several important industrial and household uses later eclipsed by soda ash, baking soda and other sodium compounds.

Sodium imparts an intense yellow color to flames. 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.

Occurrence
A FASOR used at the Starfire Optical Range used to excite sodium atoms in the upper atmosphere.

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 1100 °C.

Na2CO3 (liquid) + 2 C (solid, coke) → 2 Na (vapor) + 3 CO (gas).

It is now produced commercially through the electrolysis of liquid sodium chloride.[3][4] This is done in a Downs Cell in which the NaCl is mixed with calcium chloride to lower the melting point below 700 °C. As calcium is less electropositive than sodium, no calcium will be formed at the anode. This method is less expensive than the previous Castner process of electrolyzing sodium hydroxide.

Very pure sodium can be isolated by thermal decomposition sodium azide.[5]

Metallic sodium costs about 15 to 20 US cents per pound (US\$0.30/kg to US\$0.45/kg) in 1997 but reagent grade (ACS) sodium cost about US\$35 per pound (US\$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.[6]

Compounds

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 industries 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).

Isotopes

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.

Precautions

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. (18.5 g explosion [1]) 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+]], a graphite based dry powder with an organophosphate flame retardant; and [[Na+]], 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.

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.

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 Spectral Lines

One notable atomic spectral line of sodium vapor is the so-called D-line. The D-line is one of the classified Fraunhofer lines observed in the visible spectrum of the sun's electromagnetic radiation. Sodium vapor in the upper layers of the sun creates a dark line in the emitted spectrum of electromagnetic radiation by absorbing visible light in a band of wavelengths around 589.5 nm. This wavelength corresponds to transitions in atomic sodium in which the valence-electron transitions from a 3p to 3s electronic state. Closer examination of the visible spectrum of atomic sodium reveals that the D-line actually consists of two lines called the D1 and D2 lines at 589.8 nm and 589.2 nm, respectively. This fine structure results from a spin-orbit interaction of the valence electron in the 3p electronic state. The spin-orbit interaction couples the spin angular momentum and orbital angular momentum of a 3p electron to form two states that are respectively notated as $3p(^2P^o_{1/2})$ and $3p(^2P^o_{3/2})$ in the LS coupling scheme. The 3s state of the electron gives rise to a single state which is notated as 3s(2S1 / 2) in the LS coupling scheme. The D1-line results from an electronic transition between 3s(2S1 / 2) lower state and $3p(^2P^o_{1/2})$ upper state. The D2-line results from an electronic transition between 3s(2S1 / 2) lower state and $3p(^2P^o_{3/2})$ upper state. Even closer examination of the visible spectrum of atomic sodium would reveal that the D-line actually consists of a lot more than two lines. These lines are associated with hyperfine structure of the 3p upper states and 3s lower states. Many different transitions involving visible light near 589.5 nm may occurr between the different upper and lower hyperfine levels.[7]

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