Chemical reactions with mg. Characteristic chemical properties of Be, Mg and alkaline earth metals. Physical properties of simple substances

Group IIA contains only metals - Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium) and Ra (radium). The chemical properties of the first representative of this group, beryllium, differ most strongly from the chemical properties of the other elements of this group. His Chemical properties in many ways even more similar to aluminum than to the other metals of group IIA (the so-called "diagonal similarity"). Magnesium also differs markedly from Ca, Sr, Ba and Ra in chemical properties, but still has much more similar chemical properties with them than with beryllium. Due to the significant similarity of the chemical properties of calcium, strontium, barium and radium, they are combined into one family, called alkaline earth metals.

All elements of group IIA belong to s-elements, i.e. contain all of their valence electrons s-sublevel. Thus, the electronic configuration of the outer electron layer of all chemical elements this group has the form ns 2 , Where n– number of the period in which the element is located.

Due to the peculiarities of the electronic structure of group IIA metals, these elements, in addition to zero, are capable of having only one single oxidation state, equal to +2. simple substances, formed by elements IIA group, with participation in any chemical reactions can only oxidize, i.e. donate electrons:

Me 0 - 2e - → Me +2

Calcium, strontium, barium and radium are extremely reactive. The simple substances formed by them are very strong reducing agents. Magnesium is also a strong reducing agent. The reducing activity of metals obeys the general laws of the periodic law of D.I. Mendeleev and increases down the subgroup.

Interaction with simple substances

with oxygen

Without heating, beryllium and magnesium do not react with either atmospheric oxygen or pure oxygen due to the fact that they are covered with thin protective films consisting of BeO and MgO oxides, respectively. Their storage does not require any special methods of protection from air and moisture, unlike alkaline earth metals, which are stored under a layer of a liquid inert to them, most often kerosene.

Be, Mg, Ca, Sr, when burned in oxygen, form oxides of the composition MeO, and Ba - a mixture of barium oxide (BaO) and barium peroxide (BaO 2):

2Mg + O 2 \u003d 2MgO

2Ca + O 2 \u003d 2CaO

2Ba + O 2 \u003d 2BaO

Ba + O 2 \u003d BaO 2

It should be noted that during the combustion of alkaline earth metals and magnesium in air, the reaction of these metals with atmospheric nitrogen also proceeds side by side, as a result of which, in addition to compounds of metals with oxygen, nitrides with the general formula Me 3 N 2 are also formed.

with halogens

Beryllium reacts with halogens only at high temperatures, while the rest of the Group IIA metals already at room temperature:

Mg + I 2 \u003d MgI 2 - magnesium iodide

Ca + Br 2 \u003d CaBr 2 - calcium bromide

Ba + Cl 2 \u003d BaCl 2 - barium chloride

with non-metals of IV–VI groups

All metals of group IIA react when heated with all non-metals of groups IV-VI, but depending on the position of the metal in the group, as well as the activity of non-metals, a different degree of heating is required. Since beryllium is the most chemically inert among all metals of group IIA, its reactions with nonmetals require significantly more O high temperature.

It should be noted that the reaction of metals with carbon can form carbides of various nature. There are carbides related to methanides and conventionally considered derivatives of methane, in which all hydrogen atoms are replaced by a metal. They, like methane, contain carbon in the -4 oxidation state, and during their hydrolysis or interaction with non-oxidizing acids, methane is one of the products. There is also another type of carbides - acetylenides, which contain the C 2 2- ion, which is actually a fragment of the acetylene molecule. Carbides of the acetylenide type upon hydrolysis or interaction with non-oxidizing acids form acetylene as one of the reaction products. What type of carbide - methanide or acetylenide - will be obtained by the interaction of one or another metal with carbon depends on the size of the metal cation. As a rule, methanides are formed with metal ions having a small radius, and acetylides with larger ions. In the case of metals of the second group, methanide is obtained by the interaction of beryllium with carbon:

The remaining metals of group II A form acetylenides with carbon:

With silicon, group IIA metals form silicides - compounds of the Me 2 Si type, with nitrogen - nitrides (Me 3 N 2), phosphorus - phosphides (Me 3 P 2):

with hydrogen

All alkaline earth metals react when heated with hydrogen. In order for magnesium to react with hydrogen, heating alone, as in the case of alkaline earth metals, is not enough; in addition to high temperature, an increased pressure of hydrogen is also required. Beryllium does not react with hydrogen under any conditions.

Interaction with complex substances

with water

All alkaline earth metals actively react with water to form alkalis (soluble metal hydroxides) and hydrogen. Magnesium reacts with water only during boiling, due to the fact that when heated, the protective oxide film of MgO dissolves in water. In the case of beryllium, the protective oxide film is very resistant: water does not react with it either when boiling or even at a red heat temperature:

with non-oxidizing acids

All metals of the main subgroup of group II react with non-oxidizing acids, since they are in the activity series to the left of hydrogen. In this case, a salt of the corresponding acid and hydrogen are formed. Reaction examples:

Be + H 2 SO 4 (razb.) \u003d BeSO 4 + H 2

Mg + 2HBr \u003d MgBr 2 + H 2

Ca + 2CH 3 COOH = (CH 3 COO) 2 Ca + H 2

with oxidizing acids

− dilute nitric acid

All Group IIA metals react with dilute nitric acid. In this case, the reduction products instead of hydrogen (as in the case of non-oxidizing acids) are nitrogen oxides, mainly nitrogen oxide (I) (N 2 O), and in the case of highly dilute nitric acid, ammonium nitrate (NH 4 NO 3):

4Ca + 10HNO 3 ( razb .) \u003d 4Ca (NO 3) 2 + N 2 O + 5H 2 O

4Mg + 10HNO3 (very disaggregated)\u003d 4Mg (NO 3) 2 + NH 4 NO 3 + 3H 2 O

− concentrated nitric acid

Concentrated nitric acid at ordinary (or low) temperature passivates beryllium, i.e. does not react with it. When boiling, the reaction is possible and proceeds mainly in accordance with the equation:

Magnesium and alkaline earth metals react with concentrated nitric acid to form a wide range of different nitrogen reduction products.

− concentrated sulfuric acid

Beryllium is passivated with concentrated sulfuric acid, i.e. does not react with it under normal conditions, however, the reaction proceeds during boiling and leads to the formation of beryllium sulfate, sulfur dioxide and water:

Be + 2H 2 SO 4 → BeSO 4 + SO 2 + 2H 2 O

Barium is also passivated by concentrated sulfuric acid due to the formation of insoluble barium sulfate, but reacts with it when heated, barium sulfate dissolves when heated in concentrated sulfuric acid due to its conversion to barium hydrogen sulfate.

The remaining metals of the main group IIA react with concentrated sulfuric acid under any conditions, including in the cold. Sulfur reduction can occur to SO 2, H 2 S and S, depending on the activity of the metal, the reaction temperature and the concentration of the acid:

Mg + H 2 SO 4 ( conc .) \u003d MgSO 4 + SO 2 + H 2 O

3Mg + 4H2SO4 ( conc .) \u003d 3MgSO 4 + S↓ + 4H 2 O

4Ca + 5H2SO4 ( conc .) \u003d 4CaSO 4 + H 2 S + 4H 2 O

with alkalis

Magnesium and alkaline earth metals do not interact with alkalis, and beryllium easily reacts both with alkali solutions and with anhydrous alkalis during fusion. Moreover, when the reaction is carried out in an aqueous solution, water also participates in the reaction, and the products are tetrahydroxoberyllates of alkali or alkaline earth metals and gaseous hydrogen:

Be + 2KOH + 2H 2 O \u003d H 2 + K 2 - potassium tetrahydroxoberyllate

When carrying out the reaction with solid alkali during fusion, beryllates of alkali or alkaline earth metals and hydrogen are formed.

Be + 2KOH \u003d H 2 + K 2 BeO 2 - potassium beryllate

with oxides

Alkaline earth metals, as well as magnesium, can reduce less active metals and some non-metals from their oxides when heated, for example:

The method of restoring metals from their oxides with magnesium is called magnesiumthermy.

Reaction with sodium hydrogen phosphate. a) Place drops of solution into a test tube, add 2-3 drops of solution to the resulting mixture. Thoroughly mix the contents of the test tube with a glass rod and then add to the solution until an alkaline reaction. A white crystalline precipitate of magnesium-ammonium phosphate falls out:

or in ionic form:

b) For microcrystalloscopic detection in the form, place a drop of the analyzed solution on a glass slide. To it, add from a capillary pipette, first a drop of the solution, then a drop of the concentrated solution. Finally, add a crystal of sodium hydrogen phosphate to the solution. It is recommended to gently heat the glass slide on the lid of the water bath. In this case, crystals are formed in the form of six-beam stars (Fig. 42).

Crystals of a different type stand out from dilute solutions (Fig. 43).

Rice. 42. Crystals isolated from concentrated solutions.

Rice. 43. Crystals isolated from dilute solutions.

The resulting precipitate is soluble in acids. Reactions are directed towards the formation of weak electrolytes: hydrophosphate and dihydrophosphate ions. Under the action of strong acids, phosphoric acid is also formed:

The formation of certain reaction products depends on the acidity of the solution, that is, on the strength and concentration of the acid taken to dissolve the precipitate. When exposed to, only and are formed and are not formed, since acetic acid is a weaker acid than. Therefore, the reaction of dissolution in acetic acid should be represented as follows:

However, it should be borne in mind that when dissolved in strong acids, phosphoric acid is predominantly formed.

Reaction conditions. 1. Precipitation is recommended at .

2. and other cations (except cations of analytical group I) must be removed beforehand, because most of the cations of other analytical groups form insoluble phosphates under these conditions.

When carrying out a microcrystalloscopic reaction in the presence of, often accompanying, citric acid is added to the test solution.

This makes it possible to carry out the reaction in the presence of .

3. When depositing, a slight excess should be added to avoid the precipitation of an amorphous precipitate in an alkaline environment. However, a large excess prevents precipitation due to the formation of complex ions:

4. Heating the solution to favors the formation of a crystalline precipitate.

5. Solutions are prone to supersaturation, therefore, to accelerate the precipitation, it is recommended to rub a glass rod against the walls of the test tube.

6. With a low content or when working with dilute solutions, the final conclusion about the presence or absence can be made only after the reaction.

Reaction with -oxyquinoline (oxine). Place a drop of the solution containing , in a test tube or on a porcelain plate, add a drop of solutions and -oxyquinoline. This forms a greenish-yellow crystalline precipitate of magnesium oxyquinolate:

The ions do not precipitate with α-hydroxyquinoline.

This reaction is used to separate from the rest of the cations of group I, including from, as well as for the quantitative determination of magnesium.

Reaction conditions. 1. Precipitation is recommended for

Oxyquinolates of other ions precipitate at different values ​​of :

2. The reagent precipitates cations of many other elements, so cations, in addition to analytical groups I and II, should be absent.

3. If the reaction has to be carried out in the presence of other cations precipitated by oxyquinoline, then methods of masking interfering ions are used (see Chapter III, § 14).

4. Precipitation is best done by heating.

Reaction with -nitrobenzolaresorcinol ("magneson"). Place 2-3 drops of the studied neutral or slightly acidic solution on a drip plate, add 1-2 drops of magneson solution, which has a red-violet color in an alkaline medium. If the solution turns yellow (indicating the acidic nature of the medium), add 1-3 drops of the solution and KOH. In the presence of magnesium ions, the solution turns blue or a precipitate of the same color precipitates.

The reaction mechanism is based on precipitation, accompanied by the phenomenon of adsorption of the dye on the surface of magnesium hydroxide. The adsorption of some dyes of the so-called anthraquinone series is accompanied by a change in the original color of the non-adsorbed dye. Since the adsorption of the dye on the surface proceeds instantaneously, this phenomenon serves as an excellent means for detecting magnesium ions. do not interfere with this reaction. Ammonium salts prevent precipitation, so they must first be removed.

Drip reaction N. A. Tananaeva. Place a drop of phenolphthalein solution, a drop of neutral solution of the test substance and a drop of ammonia solution on filter paper. In this case, a red spot appears, due to the alkalinity of the ammonia solution and the resulting magnesium hydroxide. The appearance of staining does not yet give grounds to draw any conclusions regarding the presence of . When the wet spot is dried over a burner flame, the excess evaporates, the magnesium hydroxide dehydrates, and the red spot becomes discolored. If the dried spot is then moistened with distilled water, the red color appears again due to the formation.

Table 8. Effect of reagents on cations of the first analytical group

Continuation of the table. 8.

Tananaev's color reaction makes it possible to open in the presence of . Cations of other analytical groups must be removed. The performance of the reaction on filter paper is shown in fig. 12 (see Ch. III, § 5).

Reaction with hypoiodine. The freshly precipitated white precipitate turns red-brown under the action of hypoiodite due to the adsorption of elemental iodine on the surface of the magnesium hydroxide precipitate. The red-brown color is discolored when the precipitate is treated with iodide or potassium hydroxide, alcohol and other solvents that dissolve iodine, and also when sulfite or thiosulfate acts to reduce elemental iodine.

2. Ammonium salts and ions of III, IV and V analytical groups should be absent.

3. Reducing agents interfere with the reaction.

4. Phosphates and oxalates also interfere with the reaction due to the formation of compact precipitates of phosphate and magnesium oxalate, which are not able to adsorb elemental iodine, in contrast to the well-developed surface of an amorphous precipitate.

To the family alkaline earth elements include calcium, strontium, barium and radium. D. I. Mendeleev included magnesium in this family. Alkaline earth elements are named for the reason that their hydroxides, like alkali metal hydroxides, are soluble in water, that is, they are alkalis. “... They are called earthy because in nature they are found in the state of compounds that form an insoluble mass of the earth, and themselves, in the form of RO oxides, have an earthy appearance,” Mendeleev explained in Fundamentals of Chemistry.

General characteristics of the elements of group IIa

Metals of the main subgroup of group II have an electronic configuration of the external energy level ns², and are s-elements.

Easily donate two valence electrons, and in all compounds they have an oxidation state of +2

Strong reducing agents

The activity of metals and their reducing ability increases in the series: Be–Mg–Ca–Sr–Ba

Alkaline earth metals include only calcium, strontium, barium and radium, less often magnesium

Beryllium is closer to aluminum in most properties.

Physical properties of simple substances

Alkaline earth metals (compared to alkali metals) have higher t°pl. and t ° boiling., ionization potentials, densities and hardness.

Chemical properties of alkaline earth metals + Be

1. Reaction with water.

Under normal conditions, the surface of Be and Mg is covered with an inert oxide film, so they are resistant to water. In contrast, Ca, Sr and Ba dissolve in water with the formation of alkalis:

Mg + 2H 2 O - t ° → Mg (OH) 2 + H 2

Ca + 2H 2 O → Ca (OH) 2 + H 2

2. Reaction with oxygen.

All metals form oxides RO, barium peroxide - BaO 2:

2Mg + O 2 → 2MgO

Ba + O 2 → BaO 2

3. Form binary compounds with other non-metals:

Be + Cl 2 → BeCl 2 (halides)

Ba + S → BaS (sulfides)

3Mg + N 2 → Mg 3 N 2 (nitrides)

Ca + H 2 → CaH 2 (hydrides)

Ca + 2C → CaC 2 (carbides)

3Ba + 2P → Ba 3 P 2 (phosphides)

Beryllium and magnesium react relatively slowly with non-metals.

4. All alkaline earth metals dissolve in acids:

Ca + 2HCl → CaCl 2 + H 2

Mg + H 2 SO 4 (dec.) → MgSO 4 + H 2

5. Beryllium dissolves in aqueous solutions alkalis:

Be + 2NaOH + 2H 2 O → Na 2 + H 2

6. Volatile compounds of alkaline earth metals give the flame a characteristic color:

calcium compounds - brick red, strontium - carmine red, and barium - yellowish green.

Beryllium, like lithium, is an s-element. The fourth electron that appears in the Be atom is placed in the 2s orbital. The ionization energy of beryllium is higher than that of lithium due to the larger nuclear charge. In strong bases, it forms the BeO 2-2 beryllate ion. Therefore, beryllium is a metal, but its compounds are amphoteric. Beryllium, although a metal, is much less electropositive than lithium.

The high ionization energy of the beryllium atom differs markedly from other elements of the PA subgroup (magnesium and alkaline earth metals). Its chemistry is largely similar to that of aluminum (diagonal similarity). Thus, this is an element with the presence of amphoteric qualities in its compounds, among which the basic ones still prevail.

The electronic configuration of Mg: 1s 2 2s 2 2p 6 3s 2 has one significant feature compared to sodium: the twelfth electron is placed in the 2s orbital, where there is already 1e - .

Magnesium and calcium ions are indispensable elements of the vital activity of any cell. Their ratio in the body must be strictly defined. Magnesium ions are involved in the activity of enzymes (for example, carboxylase), calcium - in the construction of the skeleton and metabolism. Increasing the calcium content improves the absorption of food. Calcium excites and regulates the work of the heart. Its excess sharply increases the activity of the heart. Magnesium plays part of the role of a calcium antagonist. The introduction of Mg 2+ ions under the skin causes anesthesia without a period of excitement, paralysis of muscles, nerves and heart. Getting into the wound in the form of metal, it causes long-term non-healing purulent processes. Magnesium oxide in the lungs causes the so-called foundry fever. Frequent contact of the skin surface with its compounds leads to dermatitis. The most widely used calcium salts in medicine are CaSO 4 sulfate and CaCL 2 chloride. The first is used for plaster casts, and the second is used for intravenous infusions and as an internal remedy. It helps fight swelling, inflammation, allergies, relieves spasms of the cardiovascular system, and improves blood clotting.

All barium compounds except BaSO 4 are poisonous. Cause mengoencephalitis with damage to the cerebellum, damage to smooth heart muscles, paralysis, and in large doses - degenerative changes in the liver. In small doses, barium compounds stimulate the activity of the bone marrow.

When strontium compounds are introduced into the stomach, its disorder, paralysis, and vomiting occur; lesions are similar in signs to lesions from barium salts, but strontium salts are less toxic. Of particular concern is the appearance in the body of the radioactive isotope of strontium 90 Sr. It is extremely slowly excreted from the body, and its long half-life and, therefore, the duration of action can cause radiation sickness.

Radium is dangerous for the body with its radiation and a huge half-life (T 1/2 = 1617 years). Initially, after the discovery and production of radium salts in a more or less pure form, it began to be used quite widely for fluoroscopy, the treatment of tumors, and some serious diseases. Now, with the advent of other more accessible and cheaper materials, the use of radium in medicine has practically ceased. In some cases, it is used to produce radon and as an additive to mineral fertilizers.

The filling of the 4s orbital is completed in the calcium atom. Together with potassium, it forms a pair of s-elements of the fourth period. Calcium hydroxide is a fairly strong base. In calcium - the least active of all alkaline earth metals - the nature of the bond in the compounds is ionic.

According to its characteristics, strontium occupies an intermediate position between calcium and barium.

The properties of barium are closest to those of alkali metals.

Beryllium and magnesium are widely used in alloys. Beryllium bronzes are elastic copper alloys with 0.5-3% beryllium; aviation alloys (density 1.8) contain 85-90% magnesium ("electron"). Beryllium differs from other metals of group IIA - it does not react with hydrogen and water, but it dissolves in alkalis, since it forms amphoteric hydroxide:

Be + H 2 O + 2NaOH \u003d Na 2 + H 2.

Magnesium actively reacts with nitrogen:

3 Mg + N 2 \u003d Mg 3 N 2.

The table shows the solubility of hydroxides of elements of group II.

Traditional technical problem - hardness of water associated with the presence of Mg 2+ and Ca 2+ ions in it. Magnesium and calcium carbonates and calcium sulfate are deposited on the walls of heating boilers and pipes with hot water from bicarbonates and sulfates. They especially interfere with the work of laboratory distillers.

S-elements in a living organism perform an important biological function. The table shows their content.

The extracellular fluid contains 5 times more sodium ions than inside the cells. An isotonic solution (“physiological fluid”) contains 0.9% sodium chloride, it is used for injections, washing wounds and eyes, etc. Hypertonic solutions (3-10% sodium chloride) are used as lotions in the treatment of purulent wounds (“stretching » pus). 98% of potassium ions in the body are inside the cells and only 2% in the extracellular fluid. A person needs 2.5-5 g of potassium per day. 100 g of dried apricots contains up to 2 g of potassium. In 100 g of fried potatoes - up to 0.5 g of potassium. In intracellular enzymatic reactions ATP and ADP participate in the form of magnesium complexes.

Every day a person needs 300-400 mg of magnesium. It enters the body with bread (90 mg of magnesium per 100 g of bread), cereals (up to 115 mg of magnesium in 100 g of oatmeal), nuts (up to 230 mg of magnesium per 100 g of nuts). In addition to building bones and teeth based on hydroxylapatite Ca 10 (PO 4) 6 (OH) 2, calcium cations are actively involved in blood coagulation, transmission of nerve impulses, and muscle contraction. Adults need to consume about 1 g of calcium per day. 100 g of hard cheeses contain 750 mg of calcium; in 100 g of milk - 120 mg of calcium; in 100 g of cabbage - up to 50 mg.

The 4th analytical group includes cations Mg 2+ , Mn 2+ , Fe 2+ , Fe 3+ .

Hydroxides of group IV cations are insoluble in excess alkali and ammonia solution. They are quantitatively precipitated with an excess of NaOH solution, in the presence of hydrogen peroxide, which is a group reagent for ions of this group. All cations form sparingly soluble phosphates, oxalates, sulfides (except Mg 2+). Mn 2+ , Fe 2+ , Fe 3+ exhibit redox properties.

Reactions of magnesium ions

Reaction with alkalis.

Caustic alkalis form a white gelatinous precipitate of magnesium hydroxide:

MgCl 2 + 2NaOH \u003d Mg (OH) 2  + 2NaCl

Magnesium hydroxide is soluble in acids and ammonium salts, but insoluble in excess alkali.

Reaction with an aqueous solutionNH 3 .

Ammonia with magnesium ions forms a precipitate of magnesium hydroxide:

Mg 2+ + 2NH 3 ˙ H 2 O \u003d Mg (OH) 2  + 2NH 4 +,

which does not settle completely. In the presence of ammonium salts, the dissociation of NH 3 ˙ H 2 O decreases so much that the concentration of OH ions becomes less than necessary in order for the solubility product of Mg (OH) 2 to be exceeded. In other words, NH 4 Cl and NH 3 form a buffer solution with pH = 8.3, at which magnesium hydroxide does not precipitate.

3. Reaction with sodium hydrogen phosphate.

MgCl 2 + Na 2 HPO 4 \u003d MgHPO 4  + 2NaCl

Magnesium hydrogen phosphate is a white amorphous precipitate, soluble in mineral acids, and when heated, in acetic acid.

Executing a reaction: when carrying out the reaction in the presence of NH 3 ˙ H 2 O and NH 4 Cl precipitate a white crystalline precipitate of magnesium and ammonium phosphate. 3–4 drops of magnesium salt (tasks) are placed in a test tube, ammonia solution is added to a weak turbidity, an NH 4 Cl solution until it dissolves, and 2–3 drops of a Na 2 HPO 4 solution. The test tube is cooled under cold water by rubbing a glass rod against the inner walls of the test tube. In the presence of magnesium ions, a white crystalline precipitate forms over time:

MgCl 2 + Na 2 HPO 4 + NH 3 ˙ H 2 O \u003d MgNH 4 PO 4  + 2NaCl + H 2 O

The reaction can also be carried out as microcrystalloscopic. A drop of magnesium salt (tasks), a drop of NH 4 Cl is applied to a glass slide, kept over a bottle with a concentrated solution of NH 3 (drop down), a crystal of dry Na 2 HPO 4 12H 2 O is added and after a minute crystals of MgNH 4 PO 4 are observed in the form of dendrites (leaves) under a microscope.

Reaction with ammonium carbonate.

2MgCl 2 + 2(NH 4) 2 CO 3 + H 2 O \u003d Mg 2 (OH) 2 CO 3  + 4NH 4 Cl + CO 2 

The precipitate is slightly soluble in water and precipitates only at pH> 9. It is soluble in ammonium salts, which can be explained based on the following equilibrium: Mg 2 (OH) 2 CO 3  Mg 2 (OH) 2 CO 3  2Mg 2+ + 2OH – + CO 3 2–

With the introduction of NH 4 Cl, it dissociates NH 4 Cl  NH 4 + + Cl - . Ions NH 4 + bind with hydroxide ions into a low-dissociating compound NH 3 ˙ H 2 O, as a result of which the concentration of OH ions decreases and is not reached and the precipitate dissolves.

5. Reaction with 8-hydroxyquinoline.

8-oxyquinoline in an ammonia medium at pH 9.5–12.7 forms a greenish-yellow crystalline precipitate of the intra-complex salt of magnesium hydroxyquinolate Mg (C 9 H 6 NO) 2 2H 2 O with magnesium ions:

Mg 2+ + 2C 9 H 6 NOH + 2NH 4 OH \u003d Mg (C 9 H 6 NO) 2 + 2NH 4 +

The precipitate is soluble in acetic and mineral acids. Alkali and alkaline earth metal cations do not interfere with the reaction.

Executing a reaction: 2 drops of phenolphthalein solution and 2 M ammonia solution are added dropwise to 3–4 drops of the test solution until a pink color appears. The contents of the tube are heated to a boil and 4-5 drops of a 5% alcohol solution of 8-hydroxyquinoline are added. In the presence of magnesium, a greenish-yellow precipitate forms. The reactions are not interfered with by alkali and alkaline earth metal ions.

From this article, you will learn what magnesium is and see a real chemical miracle - the burning of magnesium in water!

In the 17th century, in the English town of Epsom, a bitter substance was isolated from a mineral spring, which had a laxative effect. This substance turned out to be magnesium sulfate hydrate or MgSO₄∙7H₂O. Because of the specific taste, pharmacists dubbed this compound "bitter salt." In 1808, the English chemist Humphrey Davy obtained an amalgam of the twelfth element using magnesia and mercury. Eleven years later, the French chemist Antoine Bussy obtained the substance in question with the help of magnesium chloride and potassium, reducing magnesium.

Magnesium is one of the most abundant elements in earth's crust. Most magnesium compounds are found in sea ​​water. This element plays important role in human life, animals and.

As a metal, magnesium is not used in its pure form - only in alloys (for example, with titanium). Magnesium allows you to create ultra-light alloys.

Physical properties of magnesium

It is a light and ductile metal of a silvery-light color with a characteristic metallic sheen.

Magnesium is oxidized by air, a sufficiently strong MgO film is formed on its surface, which protects the metal from corrosion.

The melting point of the silvery metal is 650°C and the boiling point is 1091°C.

Chemical properties of magnesium

This metal is covered with a protective oxide film. If it is destroyed, magnesium will quickly oxidize in air. Under the influence of temperature, the metal actively interacts with halogens and many non-metals. Magnesium reacts with hot water to form magnesium hydroxide as a precipitate:

Mg + 2H₂O = Mg(OH)₂ + H₂

If magnesium powder is ignited on a gas burner in a special chemical spoon and then lowered into water, the powder will begin to burn more intensively.

Here's how it goes:

Due to the intensely released hydrogen, it will be accompanied by . In this case, magnesium oxide is formed, and then its hydroxide.

Magnesium belongs to active metals, and therefore violently interacts with acids. However, this does not occur as violently as in the case of the alkali metal potassium, that is, the reaction proceeds without ignition. But with a characteristic hiss, hydrogen bubbles are actively released. And although the hydrogen bubbles lift the metal, it is not light enough to stay afloat.

The equation for the reaction of magnesium and hydrochloric acid:

Mg + 2HCl = MgCl₂ + H₂

At temperatures above 600 °C, magnesium ignites in air, emitting extremely bright light in almost the entire spectrum, similar to the Sun.

Attention! Do not try to repeat these experiments yourself!

Such a blinding flash can injure the eyes: you can get retinal burns, and in the worst case, you can lose your eyesight. Therefore, such an experience is not only among the most beautiful, but also among the most dangerous. It is not recommended to carry out this experiment without special protective dark glasses. you will find a magnesium burning experiment that can be done safely at home.

During the reaction, a white powder of magnesium oxide (also called magnesia) is formed, as well as magnesium nitride. Combustion equations:

2Mg + O₂ = 2MgO;

3Mg + N₂ = Mg₃N₂.

Magnesium continues to burn both in water and in the atmosphere carbon dioxide, so putting out such a fire is quite difficult. Extinguishing with water only exacerbates the situation, as hydrogen begins to be released, which also ignites.

An Unusual Use of Magnesium as a Light Source (1931)

The 12th element is very similar to the alkali metal. For example, it also reacts with nitrogen to form nitride:

3Mg + N₂ = Mg₃N₂.

Also, like lithium, magnesium nitride can be easily decomposed with water:

Mg₃N₂ + 6Н₂О = 3Mg(OH)₂ + 2NH₃.