Sodium Metal

Chapter I

Introduction

1.1. Background

Alkali metals is reactively metal elements which is located on the left side of elements periodic system. Alkali metals consist of six elements, that are Lithium(Li), Sodium(Na), Pottasium(K), Rubidium(Rb), Cesium(Cs), and Fransium(Fr). In this paper, it is described specificly about Sodium(Na) metal.

Sodium metal has atomic number is 11 and the atomic weight is 22,9898, and located in 3 periode. Sodium is discovered by Sir Humphry Davy in 1807, he isolated sodium for the first time by electrolysis of dried sodium hydroxide, which had been very slightly moistened. The electrolysis was powered by the combined output of three large batteries he had built.

Davy noted that the metal which formed at the wire electrode he placed in the sodium hydroxide was a liquid, but became solid on cooling and "appeared to have the lustre of silver".  It is exceedingly malleable and is much softer than any of the common metallic substances. This property does not diminish when it is cooled to 32 ^oF (0 ^oC). Davy also noted that, when added to water, sodium decomposed the water, releasing hydrogen.  He asked whether the new substance should be classed as a metal and noted that most other scientists thought it should, despite the fact that its density was much lower than the other metals then known:  "for amongst the metals themselves there are remarkable differences in this respect, platina [we now call it platinum] being nearly four times as heavy as tellurium.".  He named the new metal sodium, because he had used caustic soda or more simply soda as his source of the element.

Sodium is soft, silvery-white, highly reactive metal and is a member of the alkali metals. Its only stable isotope is ²³Na. It is an abundant element that exists in numerous minerals such as feldspars, sodalite and rock salt. Many salts of sodium are highly soluble in water and are thus present in significant quantities in the Earth's bodies of water, most abundantly in the oceans as sodium chloride.

Pure sodium is usually stored in a nonreactive substance, as it oxidizes rapidly when exposed to air, quickly forming a thick coating. The chemical element is also highly explosive when exposed to moisture and water, to the delight of many chemistry students. Since the element is so reactive, it is usually found naturally in compounds with other elements. Many of these compounds, such as salt, are extremely stable and perfectly safe to handle.

Many sodium compounds are useful, such as sodium hydroxide (lye) for soapmaking, and sodium chloride for use as a deicing agent and a nutrient. Sodium is an essential element for all animals and some plants. In animals, sodium ions are used against potassium ions to build up charges on cell membranes, allowing transmission of nerve impulses when the charge is dissipated; it is therefore classified as a dietary inorganic macro-mineral. The free metal, elemental sodium, does not occur in nature but must be prepared from sodium compounds.

Chapter II

Contents

2.1. Chemical and Physical Properties of Sodium Metals

2.1.1. Physical Properties

A positive flame test for sodium has a bright yellow color. Sodium at STP is a soft metal that can be readily cut with a knife and is a good conductor of electricity. Freshly exposed, sodium has a bright, silvery luster that rapidly tarnishes and forms a white oxide layer. These properties change at elevated pressures: at 1.5 Mbar, the color changes to black, then to red transparent at 1.9 Mbar, and finally clear transparent at 3 Mbar. All of these allotropes are insulators and electrides.

When sodium or its compounds are introduced into a flame, they turn it yellow, because the heat excites sodium atoms and moves their valence electrons from the 3s orbital to the 3p orbital; as those electrons fall back to 3s, they emit a photon with a wavelength corresponding to the D line at 589.3 nm. Spin-orbit interactions of the valence electron in the 3p orbital cause the D line to split into the D₁ (589.6 nm) and D₂ (589.0 nm) lines; hyperfine structures of both orbitals lead to many more lines. A practical use for lasers emitting light at the D line is to create artificial laser guide stars that assist in the adaptive optics for large land-based visible light telescopes.

As early as 1860, Kirchhoff and Bunsen noted the high sensitivity of a sodium flame test, and stated in Annalen der Physik und Chemie. It is stated:

“In a corner of our 60 m³ 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”.

2.1.2.     Chemical Properties of Sodium

Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts. Other alkali metals like potassium and lithium have more negative potentials. Hence, the extraction of sodium metal from its compounds (such as with sodium chloride) uses a significant amount of energy. In terms of handling properties, sodium is generally less reactive than potassium and more reactive than lithium. Like all the alkali metals, it reacts exothermically with water, to the point that sufficiently large pieces melt to a sphere and then explode; this reaction produces caustic sodium hydroxide and flammable hydrogen gas. When burned in dry air, it mainly forms sodium peroxide as well as some sodium oxide. In moist air, sodium hydroxide results.

Others chemical properties of Sodium are:

Atomic number	11
Atomic mass	22.98977 g.mol -1
Electronegativity according to Pauling	0.9
Density	0.97 g.cm -3 at 20 °C
Melting point	97.5 °C
Boiling point	883 °C
Vanderwaals radius	0.196 nm
Ionic radius	0.095 (+1) nm
Isotopes	3
Electronic shell	[Ne] 3s1
Energy of first ionisation	495.7 kJ.mol -1
Standard potential	- 2.71 V

2.2. Isotope of Sodium

20 isotopes of sodium are known, but only ²³Na is stable. Two radioactive, cosmogenic isotopes are the byproduct of cosmic ray spallation. ²²Na with a half-life of  2.6 years and ²⁴Na with a half-life of 15 hours. All other isotopes have a half-life of less than one minute. Two nuclear isomers have been discovered, the longer-lived one being ²⁴Na with a half-life of around 20.2 microseconds. Acute neutron radiation, such as from a nuclear criticality accident, converts some of the stable ²³Na in human blood to ²⁴Na, by measuring the concentration of ²⁴Na in relation to ²³Na, the neutron radiation dosage of the victim can be calculated.

2.3. Exist of Sodium in Nature

In nature, sodium exist 2,8% in Earth crust as salt (NaCl), chili gun-powder NaNO₃, Carnalit KMgCl₃.6H₂O,  Trona Na₅(CO₃)₂.(HCO₃).2H₂O, and sea water. ²³Na is created in the carbon-burning process by fusing two carbon atoms together; this requires temperatures above 600 megakelvins and a star with at least three solar masses. The Earth's crust has 2.8% sodium by weight, making it the sixth most abundant element there. Because of its high reactivity, it is never found as a pure element.

It is found in many different minerals, some very soluble, such as halite and natron, others much less soluble such as amphibole, and zeolite. The insolubility of certain sodium minerals such as cryolite and feldspar arises from their polymeric anions, which in the case of feldspar is a polysilicate. In the interstellar medium, sodium is identified by the D line; though it has a high vaporization temperature, its abundance allowed it to be detected by Mariner 10 in Mercury's atmosphere.

2.4. Biological Role of Sodium

Sodium is an essential nutrient that regulates blood volume, blood pressure, osmotic equilibrium and pH; the minimum physiological requirement for sodium is 500 milligrams per day. Sodium chloride is the principal source of sodium in the diet, and is used as seasoning and preservative, such as for pickling and jerky; most of it comes from processed foods. The DRI for sodium is 2.3 grams per day, but on average people in the United States consume 3.4 grams per day, the minimum amount that promotes hypertension. This in turn causes 7.6 million premature deaths worldwide.

The renin-angiotensin system and the atrial natriuretic peptide regulate the amount of fluid in the body. Reduction of blood pressure and sodium concentration in the kidney result in the production of renin, which in turn produces aldosterone, retaining sodium in the urine. Because of this, the osmotic pressure changes and osmoregulation systems generate the antidiuretic hormone, causing the body to retain water and restore its total amount of fluid. Receptors in the heart and blood vessels sense the resulting distension and pressure, leading to production of the atrial natriuretic peptide, causing the body to lose sodium in the urine; the osmoregulation systems detect this and remove water, restoring the total fluid levels.

Sodium is also important in neuron function and osmoregulation between cells and the extracellular fluid; their distribution is mediated in all animals by Na⁺/K⁺-ATP. Hence, sodium is the most prominent cation in extracellular fluid: the 15 liters of it in a 70 kg human have around 50 grams of sodium, 90% of the body's total sodium content.

In C4 plants, sodium is a micronutrient that aids in metabolism, specifically in regeneration of phosphoenolpyruvate and synthesis of chlorophyll. In others, it substitutes for potassium in several roles, such as maintaining turgor pressure and aiding in the opening and closing of stomata. Excess sodium in the soil limits the uptake of water due to decreased water potential, which may result in wilting; similar concentrations in the cytoplasm can lead to enzyme inhibition, which in turn causes necrosis and chlorosis. To avoid these problems, plants developed mechanisms that limit sodium uptake by roots, store them in cell vacuoles, and control them over long distances, excess sodium may also be stored in old plant tissue, limiting the damage to new growth.

2.5. Production of Sodium

Enjoying rather specialized applications, only about 100,000 tonnes of metallic sodium are produced annually. Metallic sodium was first produced commercially in 1855 by carbothermal reduction of sodium carbonate at 1100 °C, in what is known as the Deville process.

Na₂CO₃ + 2 C → 2 Na + 3 CO

A related process based on the reduction of sodium hydroxide was developed in 1886.

Sodium is now produced commercially through the electrolysis of molten sodium chloride, based on a process patented in 1924. 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.

Reagent-grade sodium in tonne quantities sold for about US$3.30/kg in 2009; lower purity metal sells for considerably less. The market for sodium is volatile due to the difficulty in its storage and shipping; it must be stored under a dry inert gas atmosphere or anhydrous mineral oil to prevent the formation of a surface layer of sodium oxide or sodium superoxide. These oxides can react violently in the presence of organic materials. Sodium will also burn violently when heated in air. Smaller quantities of sodium cost far more, in the range of US$165/kg; the high cost is partially due to the expense of shipping hazardous material.

2.6. Precautions of Sodium

Care is required in handling elemental sodium, as it is potentially explosive and generates flammable hydrogen and caustic sodium hydroxide upon contact with water; powdered sodium may combust spontaneously in air or oxygen. Excess sodium can be safely removed by hydrolysis in a ventilated cabinet. This is typically done by sequential treatment with isopropanol, ethanol and water. Isopropanol reacts very slowly, generating the corresponding alkoxide and hydrogen.

Fire extinguishers based on water accelerate sodium fires; those based on carbon dioxide and bromochlorodifluoromethane lose their effectiveness when they dissipate. An effective extinguishing agent is Met-L-X, which comprises approximately 5%, in sodium chloride together with flow agents; it is most commonly hand-applied with a scoop. Other materials include Lith+, which has graphite powder and an organophosphate flame retardant, and dry sand.

2.7. Application of Sodium

Sodium in its metallic form is very important in making esters and in the manufacture of organic compounds. Sodium is also a component of sodium chloride (NaCl) a very important compount found everywhere in the living environment. Other uses are: to improve the structure of certain alloys; in soap, in combination with fatty acids, in sodium vapor lamps, to descal metals, to purify molten metals. Solid sodium carbonate is needed to make glass.

2.8. Health Effects of Sodium

Sodium is a compound of many foodstuffs, for instance of common salt. Sodium occurs naturally in most foods. The most common form of sodium is sodium chloride, which is table salt. Milk, beets, and celery also naturally contain sodium, as does drinking water, although the amount varies depending on the source.

Sodium is also added to various food products. Some of these added forms are monosodium glutamate, sodium nitrite, sodium saccharin, baking soda (sodium bicarbonate), and sodium benzoate. These are ingredients in condiments and seasonings such as Worcestershire sauce, soy sauce, onion salt, garlic salt, and bouillon cubes.

Processed meats, such as bacon, sausage, and ham, and canned soups and vegetables are all examples of foods that contain added sodium. Fast foods are generally very high in sodium.

Sodium is necessary for humans to maintain the balance of the physical fluids system. Sodium is also required for nerve and muscle functioning. Too much sodium can damage our kidneys and increases the chances of high blood pressure. The amount of sodium a person consumes each day varies from individual to individual and from culture to culture, some people get as little as 2 g/day, some as much as 20 grams. Sodium is essential, but controversely surrounds the amount required.

Contact of sodium with water, including perspiration causes the formation of sodium hydroxide fumes, which are highly irritating to skin, eyes, nose and throat. This may cause sneezing and coughing. Very severe exposures may result in difficult breathing, coughing and chemical bronchitis. Contact to the skin may cause itching, tingling, thermal and caustic burns and permanent damage. Contact with eyes may result in permanent damage and loss of sight.

2.8.1.     Recomendation Uses of Sodium

Dietary sodium is measured in milligrams (mg). Table salt is 40% sodium; 1 teaspoon of table salt contains 2,300 mg of sodium.

Healthy adults should limit sodium intake to 2,300 mg per day while individuals with high blood pressure should consume no more than 1,500 mg per day. Those with congestive heart failure, liver cirrhosis, and kidney disease may need much lower amounts.

Specific recommendations regarding sodium intake do not exist for infants, children, and adolescents. Eating habits and attitudes about food formed during childhood are likely to influence eating habits for life. For this reason, moderate intake of sodium is suggested.

Chapter III

Conclusion

3.1. Conclusion

From the explanation above, we known that Sodium is one of member of alkali metals. It has atomic number 11, located on 3 periode, the atomic weight is 22,9898, and the form is solid. It is a reactive metals. Sodium is discovered by Sir Humphry Davy in 1807. In nature, sodium exist 2,8% in Earth crust as salt (NaCl), chili gun-powder NaNO₃, Carnalit KMgCl₃.6H₂O,  Trona Na₅(CO₃)₂.(HCO₃).2H₂O, and sea water. And sodium is used in foodstuff as preservative. Some of these added forms are monosodium glutamate, sodium nitrite, sodium saccharin, baking soda (sodium bicarbonate), and sodium benzoate. Contact of sodium with water, including perspiration causes the formation of sodium hydroxide fumes, which are highly irritating to skin, eyes, nose and throat. This may cause sneezing and coughing. Very severe exposures may result in difficult breathing, coughing and chemical bronchitis. Contact to the skin may cause itching, tingling, thermal and caustic burns and permanent damage. Contact with eyes may result in permanent damage and loss of sight.

References

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