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Sodium

2007 Schools Wikipedia Selection. Related subjects: Chemical elements


                11                neon ← sodium → magnesium
                Li
                ↑
                Na
                ↓
                K

                                  Periodic Table - Extended Periodic Table

                                                                   General
                                       Name, Symbol, Number sodium, Na, 11
                                             Chemical series alkali metals
                                              Group, Period, Block 1, 3, s
                                                  Appearance silvery white
                                         Atomic mass 22.98976928 (2) g/mol
                                          Electron configuration [Ne] 3s^1
                                               Electrons per shell 2, 8, 1
                                                       Physical properties
                                                               Phase solid
                                      Density (near r.t.) 0.968 g·cm^−3
                                   Liquid density at m.p. 0.927 g·cm^−3
                                                   Melting point 370.87  K
                                                  (97.72 ° C, 207.9 ° F)
                                                      Boiling point 1156 K
                                                     (883 ° C, 1621 ° F)
                                             Critical point (extrapolated)
                                                            2573 K, 35 MPa
                                          Heat of fusion 2.60 kJ·mol^−1
                                   Heat of vaporization 97.42 kJ·mol^−1
                         Heat capacity (25 °C) 28.230 J·mol^−1·K^−1

   CAPTION: Vapor pressure

                                          P/Pa   1  10  100 1 k 10 k 100 k
                                         at T/K 554 617 697 802 946  1153

                                                         Atomic properties
                                     Crystal structure cubic body centered
                                                        Oxidation states 1
                                                    (strongly basic oxide)
                                    Electronegativity 0.93 (Pauling scale)
                                                       Ionization energies
                                           ( more) 1st: 495.8 kJ·mol^−1
                                                    2nd: 4562 kJ·mol^−1
                                                  3rd: 6910.3 kJ·mol^−1
                                                      Atomic radius 180 pm
                                              Atomic radius (calc.) 190 pm
                                                    Covalent radius 154 pm
                                               Van der Waals radius 227 pm
                                                             Miscellaneous
                                            Magnetic ordering paramagnetic
                               Electrical resistivity (20 °C) 47.7 nΩ·m
                        Thermal conductivity (300 K) 142 W·m^−1·K^−1
                         Thermal expansion (25 °C) 71 µm·m^−1·K^−1
                               Speed of sound (thin rod) (20 °C) 3200 m/s
                                                    Young's modulus 10 GPa
                                                     Shear modulus 3.3 GPa
                                                      Bulk modulus 6.3 GPa
                                                         Mohs hardness 0.5
                                                 Brinell hardness 0.69 MPa
                                             CAS registry number 7440-23-5
                                                         Selected isotopes

                  CAPTION: Main article: Isotopes of sodium

                                   iso   NA  half-life DM  DE ( MeV)  DP
                                  ^22Na syn  2.602 y   β^+ 0.546     ^22Ne
                                                       ε   -         ^22Ne
                                                       γ   1.2745    -
                                  ^23Na 100% Na is stable with 12 neutrons

                                                                References

   Sodium ( IPA: /ˈsəʊdiəm/) is a chemical element which has the symbol Na
   (Latin natrium), atomic number 11, atomic mass 22.9898 g/mol, oxidation
   number +1. Sodium is a soft, silvery, highly reactive element and is a
   member of the alkali metals within "group 1" (formally 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 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. It is also a
   component of many earthly minerals, and it is an essential element for
   animal life.

Notable Characteristics

   Compared with the other alkali metals, sodium is generally more
   reactive than lithium and less so than potassium, 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 swim around the surface of the
   water until they are consumed by it, whereas large pieces will explode.
   While sodium metal reacts with water, you can observe that the sodium
   piece melts with the heat of the reaction to form a perfect sphere
   shape if the reacting sodium is small enough. The reaction with water
   produces very caustic sodium hydroxide and highly flammable hydrogen
   gas. In any case these are considered an extreme hazard and will cause
   severe skin and eye injury. When burned in air, sodium forms sodium
   peroxide Na[2]O[2], or with limited oxygen, the oxide Na [2]O (unlike
   lithium, the nitride is not formed). If burned in oxygen under
   pressure, sodium superoxide NaO[2] 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 NaBiO[3], sodium octamolybdate Na[2]Mo[8]O[25]• 4H[2]O,
   sodium thioplatinate Na[4]Pt[3]S[6], sodium uranate Na[2]UO[4]. Sodium
   meta-antimonate's 2NaSbO[3]•7H[2]O solubility is 0.3g/L as is the pyro
   form Na[2]H[2]Sb[2]O[7]• H[2]O of this salt. Sodium metaphosphate
   NaPO[3] has a soluble and an insoluble form.

   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^+, NH[4]^+, 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 and food preservation, has been an important commodity in
   human activities (the English word salary refers to salarium, the
   prerequisite 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, 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.
   Enlarge
   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 spectral 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.
   Enlarge
   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 (English, 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 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 stems from
   Berzelius' publication of his system of atomic symbols in Thomas
   Thomson's Annals of Philosophy.

   Sodium imparts an intense yellow colour 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 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.
   Enlarge
   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 1100 °C.

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

   It is now produced commercially through the electrolysis of liquid
   sodium chloride. This is done in a Down's cell in which the NaCl is
   mixed with calcium chloride 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 method
   of electrolyzing sodium hydroxide.

   Metallic sodium cost 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 behaviour under pressure

   Under extreme pressure, sodium departs from common melting behaviour.
   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 behaviour 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. (Gregoryanz, et al., Phys.
   Rev. Lett. 94, 185502 (2005))

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 industry are common
   salt (NaCl), soda ash (Na[2]CO[3]), baking soda (NaHCO[3]), caustic
   soda (NaOH), Chile saltpeter (NaNO[3]), di- and tri-sodium phosphates,
   sodium thiosulfate (hypo, Na[2]S[2]O[3] · 5H[2]O), and borax
   (Na[2]B[4]O[7] · 10H[2]O).

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 metal, lye solution, and sometimes flame. This behaviour 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. When the metal itself
   catches fire (as opposed to just the hydrogen gas generated from it) it
   burns at high temperatures and also melts, which spreads the flame and
   exposes even more surface area to the air.

   Few common fire extinguishers work on sodium fires. Water, of course,
   exacerbates sodium fires, as do water-based foams. CO[2] and Halon are
   often ineffective on sodium fires, which reignite when the extinguisher
   dissipates. Among the very few materials effective on a sodium metal
   fire are Pyromet and Met-L-X. Pyromet is a NaCl/(NH[4])[2]HPO[4] 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 Na[2]CO[3]-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.

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