[Chemistry4me]

You mean like a compound of sodium? Or the metal [] []

[Chemistry4me]

Assuming that it was a solution of sodium chloride, you would need to decrease the solubility so that it could precipitate out of solution. Sodium nitrate (or anything that contains Na⁺) and say, copper chloride (or anything that contains Cl^-) both share a common ion with sodium chloride. In the first case it is the sodium ion and in the second case it is the chloride ion. If either of these salts is added to the equlibrium system:
NaCl(s) + aq ↔ Na⁺(aq) + Cl^-(aq)
the solubility of sodium chloride decreases. In each case, the decrease is due to the common ion. The effect itself is known as (who would have thought []) the common ion effect. While equilibrium considerations allow this effect to be anticipated, a knowledge of solubility products (which I do not have) allows its extent to be calculated. So some compounds of sodium may be able to be precipitated out (the less soluble ones) whereas others might not...
And hence, if you wanted to increase solubility, you would do the opposite. []

[Chemistry4me]

Magnesium, silvery-white metallic element in group 2 of the periodic table. Magnesium is soft and ductile when heated. The metal is not attacked by oxygen, water, or alkalis at room temperature; it does react with acids. When heated to about 800° C, it reacts with oxygen and emits a bright, white light. It occurs in nature only in chemical combination with other elements, particularly as the minerals carnallite, dolomite, and magnesite; in many rock-forming silicates; and as salts, such as magnesium chloride, in ocean and saline-lake waters. Magnesium forms divalent compounds, chief among which are magnesium carbonate (MgCO3), which is formed by the reaction of a magnesium salt and sodium carbonate and is used as a refractory and insulating material; magnesium chloride (MgCl2•6H2O), which is formed by reacting magnesium carbonate or oxide with hydrochloric acid and is used as dressing and filler for cotton and woollen fabrics, in paper manufacture, and in cements and ceramics; magnesium citrate (Mg3(C6H 5O7)2•4H2O), which is formed by the reaction of magnesium carbonate with citric acid and is used in medicine and effervescent beverages; magnesium hydroxide (Mg(OH)2), formed by the reacting of magnesium salt and sodium hydroxide and used in medicine as the laxative “milk of magnesia”, in sugar refining and magnesium sulphate (MgSO4•7H2O), is well known as Epsom salt.

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Assuming that it was a solution of sodium chloride, you would need to decrease the solubility so that it could precipitate out of solution. Sodium nitrate (or anything that contains Na⁺) and say, copper chloride (or anything that contains Cl^-) both share a common ion with sodium chloride. In the first case it is the sodium ion and in the second case it is the chloride ion. If either of these salts is added to the equlibrium system:
NaCl(s) + aq ↔ Na⁺(aq) + Cl^-(aq)
the solubility of sodium chloride decreases.
In each case, the decrease is due to the common ion. The effect itself is known as (who would have thought []) the common ion effect.

No, it's not this the effect. You have it when, e.g., you add a concentrated CuCl₂ solution, (or HCl) to a NaCl solution (not to a NaNO₃ or other salt solution). Then you really can see the formation of a precipitate of NaCl (if you pass the solubility product).

What I intended is not simply to decrease the solubility of a sodium salt, but to form a new one which solubility is almost zero, so that you can make a quantitative analysis of the sodium originally present in solution (or to identify it qualitatively by observing the formation of a precipitate).

For example, if you have a BaCl₂ solution, you add e.g., a conc. solution of K₂SO₄ and you get (double exchange) BaSO₄ and KCl. The last one stay in solution, while the first one precipitates because is very insoluble (the amount of Ba⁺⁺ which remains in solution can be neglected). You filter out the solution, you wash and then dry the precipitate, you weight it and you have made a quantitative analysis of the Ba originally present in solution.
What would you use to precipitate sodium (insted of Barium) out of a solution?

I'm just trying to give you some homework to improve your skill in chemistry.

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...
 magnesium hydroxide (Mg(OH)2), formed by the reacting of magnesium salt and sodium hydroxide and used in medicine as the laxative “milk of magnesia”, in sugar refining...

Mg(OH)₂ is also used against stomach hyperacidity because it neutralizes it.

[Chemistry4me]

Are you asking me to find a compound of sodium that is insoluble? I could do with a bit of homework right now [] []

[Chemistry4me]

Well, according to Wikipedia: 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. So I guess any of those will do to precipitate Na out of solution... unless I have misunderstood your question []

[Chemistry4me]

Aluminium, the most abundant metallic element in the Earth's crust. The element is in group 13 of the periodic table. Aluminium is a strongly electropositive metal and extremely reactive. In contact with air, aluminium rapidly becomes covered with a tough, transparent layer of aluminium oxide that resists further corrosive action. For this reason, materials made of aluminium do not tarnish or rust. The metal reduces many other metallic compounds to their base metals. For example, when thermite (a mixture of powdered iron oxide and aluminium) is heated, the aluminium rapidly removes the oxygen from the iron; the heat of the reaction is sufficient to melt the iron. This phenomenon is used in the Thermit process for welding iron. The metal is becoming increasingly important architecturally, for both structural and ornamental purposes. Aluminium siding, storm windows, and foil make excellent insulators. The metal is also used as a material in low-temperature nuclear reactors because it absorbs relatively few neutrons. Aluminium becomes stronger and retains its toughness as it gets colder and is therefore used at cryogenic temperatures. Because of its light weight, ease of forming, and compatibility with foods and beverages, aluminium is widely used for containers, flexible packages, and easy-to-open bottles and cans.

[Chemistry4me]

Silicon, semimetallic element that is the second most common element on Earth, after oxygen. Silicon is prepared as a brown amorphous powder or as grey-black crystals. It is obtained by heating silica, or silicon dioxide (SiO2), with a reducing agent, such as carbon or magnesium, in an electric furnace. Silicon is not attacked by nitric, hydrochloric, or suphuric acids, but it dissolves in hydrofluoric acid. It dissolves in sodium hydroxide, forming sodium silicate and hydrogen gas. At ordinary temperatures silicon is impervious to air, but at high temperatures it reacts with oxygen, forming a layer of silica that does not react further. At high temperatures it also reacts with nitrogen and chlorine to form silicon nitride and silicon chloride. Silicon constitutes about 28 per cent of the Earth's crust. It does not occur in the free, elemental state, but is found in the form of silicon dioxide and in the form of complex silicates. Silicon-containing minerals constitute nearly 40 per cent of all common minerals, including more than 90 per cent of igneous-rock-forming minerals. The mineral quartz, varieties of quartz (such as cornelian, chrysoprase, onyx, flint, and jasper), and the minerals cristobalite and tridymite are the naturally occurring crystal forms of silica. Silicon dioxide is the principal constituent of sand. Silicon is a semiconductor, in which the resistivity to the flow of electricity at room temperature is in the range between that of metals and that of insulators.

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Well, according to Wikipedia: 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. So I guess any of those will do to precipitate Na out of solution... unless I have misunderstood your question []

It's too simple! To precipitate sodium from a solution into an insoluble salt of sodium, you have to find a chemical which instead is soluble! If, e.g., you want to precipitate sodium bismuthate NaBiO₃, you should show me how this pure compund can form adding some chemical to the Na⁺ solution. Does it exist, for example, a soluble bismuthate which forms BiO₃^- ions in solution and that this ions react with Na⁺ ions forming pure sodium bismuthate?

[Chemistry4me]

Oh... right, okay, let me see....

[Chemistry4me]

Hmm... this is harder than I thought.

[Chemistry4me]

What do you define as soluble or insoluble?

[Chemistry4me]

Okay, sodium uranate can be prepared by reacting U3O8 with sodium carbonate
That's about the only one I can find at the moment

[Chemistry4me]

Phosphorus, reactive non-metallic element that is important to living organisms and has many industrial uses. Phosphorus is in group 15 of the periodic table. Phosphorus exists in three main allotropic forms: ordinary (or white) phosphorus, red phosphorus, and black phosphorus. Of these, only white and red phosphorus are of commercial importance. When freshly prepared, ordinary phosphorus is white, but it turns light yellow when exposed to sunlight. It is a crystalline, translucent, waxy solid, which glows faintly in moist air and is extremely poisonous. It ignites spontaneously in air at 34° C and must be stored under water. It is insoluble in water, slightly soluble in organic solvents, and very soluble in carbon disulphide. It does not occur in the free state but is found mostly as a phosphate, as in phosphate rock and apatite. It is also found in the combined state in all fertile soil and in many natural waters. Red phosphorus is a microcrystalline, non-poisonous powder. It sublimates at 416° C and has a relative density of 2.34. Black phosphorus is made by heating white phosphorus at 200° C  at very high pressure. The most important commercial compounds of phosphorus are phosphoric acid and the salts of phosphoric acid, called phosphates. The bulk of phosphorus-containing compounds are used as fertilizers. Phosphorous compounds are also used in clarifying sugar solutions, weighing silk, and fireproofing, and in such alloys as phosphor bronze and phosphor copper. White phosphorus is used in the making of rat poison, and red phosphorus is used in matches.

[Chemistry4me]

Sulphur, tasteless, odourless, light yellow non-metallic element. All forms of sulphur are insoluble in water, but the crystalline forms are soluble in carbon disulphide. When ordinary sulphur melts, it forms a straw-coloured liquid that turns darker with additional heating and then finally boils. When molten sulphur is slowly cooled, its physical properties change in accordance with the temperature, pressure, and method of crust formation. Sulphur thus exists in a variety of forms called allotropic modifications, which consist of the liquids Sλ, and Sµ, and several solid varieties. Sulphur combines with hydrogen and the metallic elements in the presence of heat to form sulphides. The most common sulphide is hydrogen sulphide, H2S, a colourless, poisonous gas with the odour of rotten eggs. Sulphur combines also with chlorine in several proportions to produce sulphur monochloride, S2Cl2, and sulphur dichloride, SCl2. When burned in air, sulphur combines with oxygen to form sulphur dioxide, SO2, a heavy, colourless gas with a characteristic, suffocating odour. Sulphur dioxide is released into the atmosphere in the combustion of fossil fuels, such as gas, petroleum, and coal, and constitutes one of the most troublesome air pollutants.

[Chemistry4me]

It's too simple! To precipitate sodium from a solution into an insoluble salt of sodium, you have to find a chemical which instead is soluble! If, e.g., you want to precipitate sodium bismuthate NaBiO₃, you should show me how this pure compund can form adding some chemical to the Na⁺ solution. Does it exist, for example, a soluble bismuthate which forms BiO₃^- ions in solution and that this ions react with Na⁺ ions forming pure sodium bismuthate?

Do you have an answer?

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It's too simple! To precipitate sodium from a solution into an insoluble salt of sodium, you have to find a chemical which instead is soluble! If, e.g., you want to precipitate sodium bismuthate NaBiO₃, you should show me how this pure compund can form adding some chemical to the Na⁺ solution. Does it exist, for example, a soluble bismuthate which forms BiO₃^- ions in solution and that this ions react with Na⁺ ions forming pure sodium bismuthate?

Do you have an answer?

A possible solution of the problem is to use potassium exahydroxyantimonate:

KSb(OH)₆

a solution of that salt added to an unknown solution, in the presence of sodium forms a white precipitate of sodium exahydroxyantimonate:

Sb(OH)₆^- + Na⁺ → NaSb(OH)₆↓

To prepare KSb(OH)₆ you make react KOH with antimonic acid HSbO₃:

KOH + HSbO₃ + 2H₂O → K⁺ + Sb(OH)₆^-

[Chemistry4me]

Ok, I did not know that...

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Ok, I did not know that...

Maybe you will study that in a few months or in the second year.