Preparation of a colloidal solution of AgI. Qualitative reactions to chloride, bromide, and iodide ions Silver in nature

IODINE MOLECULAR.

a) on the chloride ion - the effect of a solution of silver nitrate → a white cheesy precipitate of silver chloride is formed:

Cl - + Ag + = AgCl↓

silver diammine chloride

b) for Bromide ion:

Br - + Ag + = AgBr↓

The reaction is pharmacopoeial.

2) with chlorine water

Cl 2 + 2 NaBr = 2 NaCl + Br 2

The reaction is pharmacopoeial.

c) for iodide ion:

KI + AgNO 3 = AgI↓ + KNO 3

I - + Ag + = AgI↓

The reaction is pharmacopoeial.

Cl 2 + 2 NaI = 2 NaCl + I 2

The reaction is pharmacopoeial.

Conclusions: a) on the chloride ion - the effect of a solution of silver nitrate → a white cheesy precipitate of silver chloride is formed:

NaCl + AgNO 3 = AgCl↓ + NaNO 3

Cl - + Ag + = AgCl↓

The precipitate is insoluble in nitric acid, but is easily soluble in ammonia to form a complex compound:

AgCl + 2 NH 3 = Cl

or AlCl + 2 NH 4 OH = Cl + 2 H 2 O

silver diammine chloride

When concentrated sulfuric acid is added to a solution of silver diammine chloride, a precipitate is released again:

Cl + 2 HNO 3 = AgCl↓ + 2 NH 4 NO 3

The reaction is pharmacopoeial.

AgCl – also soluble in sodium thiosulfate.

T.V.: add 2 drops of AgNO 3 solution to 2 drops of NaCl solution.

A concentrated ammonia solution is added to the solution containing the precipitate until the precipitate is completely dissolved. The resulting solution is acidified with concentrated nitric acid and the formation of a precipitate is observed.

b) for Bromide ion:

1) action of silver nitrate → yellowish-white precipitate of silver bromide:

NaBr + AgNO 3 = AgBr↓ + NaNO 3

Br - + Ag + = AgBr↓

The precipitate is insoluble in HNO3, poorly soluble in ammonia, unlike silver chloride, and readily soluble in sodium thiosulfate solution.

The reaction is pharmacopoeial.

T.V.: To 4 drops of NaBr solution add 4 drops of AgNO 3 solution. The solution with the precipitate is divided into two parts. A sodium thiosulfate solution is added to one part, and a concentrated ammonia solution is added to the other, and the dissolution of the AgBr precipitate in these reagents is compared.

2) with chlorine water

Chlorine water added to a bromide solution releases free bromine from it, which dissolves in carbon disulfide or chloroform, coloring the solvent layer in orange:

Cl 2 + 2 NaBr = 2 NaCl + Br 2

With a large excess of chlorine water, the color disappears due to the formation of BrCl, which has a lighter color.

The reaction is pharmacopoeial.

T.V. To 5 drops of NaBr solution add 1 ml of chloroform, 1-2 drops of diluted H2SO4 and then drop by drop, with vigorous shaking, 2-3 drops of chlorine water. Coloration of the chloroform layer is observed.

c) for iodide ion:

1) silver nitrate releases a light yellow cheesy precipitate of silver from iodides:

KI + AgNO 3 = AgI↓ + KNO 3

I - + Ag + = AgI↓

The precipitate is insoluble in nitric acid and ammonia solution and is poorly soluble in sodium thiosulfate solution.

The reaction is pharmacopoeial.

T.V.: A little AgNO3 solution is added to the KI solution. Check the dissolution of the precipitate in a sodium thiosulfate solution.

2) Chlorine water releases free iodine from iodide solutions, which colors carbon disulfide or chloroform reddish-violet and the starch solution blue.

Cl 2 + 2 NaI = 2 NaCl + I 2

T.V.: To 5 drops of NI (KI) solution add 1 ml of chloroform, 2-3 drops of diluted H2SO 4 and then drop by drop, with vigorous shaking, 2-3 drops of chlorine water. The chloroform layer is observed to turn reddish-violet. 1 drop of KI solution, 1 drop of chlorine water and 2 drops of starch solution are poured into another test tube. Observe the color change.

3) Iron (III) chloride, conc. H 2 SO 4 and some other oxidizing agents oxidize the I ion to free iodine; For example:

2 FeCl 3 + 2 KI = 2 FeCl 2 + 2 KCl + I 2

The reaction is pharmacopoeial.

T.V.: Solutions of KI, HCl, FeCl 3 are applied sequentially, 1 drop at a time, onto the filter paper in one place. Observe the appearance of a brown spot, turning blue from a drop of starch.

d) On molecular iodine → effect of starch → blue color.

Qualitative reactions:

a) on the chloride ion - the effect of a solution of silver nitrate → a white cheesy precipitate of silver chloride is formed; the precipitate is insoluble in nitric acid, but is easily soluble in ammonia to form the complex compound silver diammine chloride.

When concentrated sulfuric acid is added to a solution of silver diammine chloride, a precipitate is released again:

b) for Bromide ion:

1) action of silver nitrate → yellowish-white precipitate of silver bromide; the precipitate is insoluble in HNO3, poorly soluble in ammonia, unlike silver chloride, and soluble in sodium thiosulfate solution.

The reaction is pharmacopoeial.

2) with chlorine water

Chlorine water added to a bromide solution releases free bromine, which dissolves in carbon disulfide or chloroform, turning the solvent layer orange.

The reaction is pharmacopoeial.

c) for iodide ion:

1) silver nitrate releases a light yellow cheesy precipitate of silver from iodides.

The precipitate is insoluble in nitric acid and ammonia solution and is poorly soluble in sodium thiosulfate solution.

The reaction is pharmacopoeial.

2) Chlorine water releases free iodine from iodide solutions, which colors carbon disulfide or chloroform reddish-violet, and the starch solution blue.

3) Iron (III) chloride oxidizes the I ion to free iodine;

The reaction is pharmacopoeial.

d) On molecular iodine → effect of starch → blue color.

The second analytical group includes the cations Ag +, Pb 2+, 2+.

These elements are in different groups periodic system D.I. Mendeleev. They have either complete 18-electron outer layers or shells containing 18+2 electrons in the two outer layers, which causes the same ratio of their halide ions.

The group reagent for cations of analytical group II is a 2 mol/l solution of hydrochloric acid. Ag + , Pb 2+ , 2+ cations, when interacting with it, form precipitates that are sparingly soluble in water and dilute acids white:

Ag + + Cl - → AgCl

Pb 2+ + 2Cl - → PbCl 2

Excess of the reagent and the use of concentrated hydrochloric acid should be avoided, as soluble complex compounds may be formed:

AgCl + 2 HCl → H 2

PbCl 2 + HCl → H

The solubility of chlorides varies. At 20 0 C: lead chloride - 11.0 g/l, silver chloride - 1.8·10 -3 g/l, mercury chloride (I) - 2.0·10 -4 g/l. When the water temperature increases to 100 0 C, the solubility of PbCl 2 increases 3 times, while the solubility of AgCl and Hg 2 Cl 2 practically remains the same. This property is used to separate Pb 2+ cations from 2+ and Ag + cations.

Mercury (I) chloride, when interacting with a solution of ammonium hydroxide, forms mercury (I) amidochloride, which is unstable and decomposes into poorly soluble mercury (II) amidochloride and metallic mercury, which gives the precipitate a black color:

Hg 2 Cl 2 + 2 NH 4 OH → Cl + NH 4 Cl + 2H 2 O

Cl → Cl + Hg

This allows the 2+ cation to be separated from the Ag+ cation.

Silver chloride is soluble under the action of ammonium hydroxide solution to form the complex compound diammine silver chloride (I):

AgCl + 2 NH 4 OH → Cl + 2 H 2 O

From the above it follows that the precipitate of lead chloride is the most soluble, as a result of which it does not completely precipitate with this group of cations and partially remains in solution.

Nitrates of silver, lead and mercury (I) are highly soluble in water. The solubility of sulfates is low and decreases in the series Ag + - 2+ - Pb 2+. Carbonates and sulfides are poorly soluble in water. Silver hydroxide is unstable, mercury does not exist (only oxide), and lead hydroxide is amphoteric. Mercury (I) salts are unstable and are prone to disproportionation reactions with the release of free mercury and the formation of the corresponding mercury (II) compounds.

The effect of the group reagent HCl on cations of analytical group II (Ag +, Pb 2+, 2+).

Hydrochloric acid forms poorly soluble white precipitates with all group II cations. The reaction of HCl with Ag + cations is PHARMACOPOEIAL. (See the chemistry of the reaction above, in general characteristics groups).

Methodology: Take 3 test tubes. Place 3-4 drops of mercury (I) nitrate solution in the first, 3-4 drops of silver nitrate solution in the second, and 3-4 drops of lead nitrate solution in the third. Add 3-4 drops of 2mol/l hydrochloric acid solution to all three test tubes. The formation of white precipitates is observed in all three test tubes. Add (excess) 6-8 drops of ammonia solution to the first and second test tubes, 5 drops of water to the third test tube and heat.

Observe the events occurring.

Analytical reactions of the cation (Ag +).

Attention! Silver salts are poisonous! Work carefully!

1. Reaction with potassium iodide.

Potassium iodide with the Ag+ cation forms a yellow precipitate of silver iodide AgI, insoluble in a concentrated ammonia solution:

Ag + + I - → AgI

Methodology: Place 2-3 drops of silver nitrate solution into a test tube, add 2-3 drops of potassium or sodium iodide. Observe the formation of sediment yellow.

2. Reaction with potassium bromide.

Potassium bromide with Ag+ cations forms a pale yellow precipitate of silver bromide AgBr, which is partially soluble in concentrated ammonia solution:

Ag + + Br - → AgBr

Methodology: Place 2-3 drops of silver nitrate in a test tube, add 2-3 drops of potassium bromide. The formation of a pale yellow precipitate is observed.

3. Reaction with potassium chromate.

Potassium chromate with Ag + cations in a neutral or slightly acetic environment forms a brick-red precipitate of Ag 2 CrO 4:

Ag + + CrO 4 2- → Ag 2 CrO 4

The precipitate dissolves in a concentrated ammonia solution; in ammonia, strongly acidic media, a precipitate does not form.

Ions Pb 2+, Ba 2+ and others, which precipitate with CrO 4 2-, interfere with this reaction.

Method: Place 2-3 drops of silver nitrate solution into a test tube and add 1-2 drops of potassium chromate solution. Observe the formation of precipitate. Check the solubility of the precipitate in acetic acid and concentrated ammonia solution.

Analytical reactions of the 2+ cation.

Attention! All mercury salts are poisonous and require careful handling!

Easily form amalgams, do not allow contact with gold jewelry!

1. Reduction of 2+ to Hg with tin chloride (S).

When a solution of mercury (I) salt is exposed to a solution of tin (II) chloride, a white precipitate of Hg 2 Cl 2 is initially formed, which upon standing gradually darkens due to the reduction of 2+ ions to metallic mercury

2+ + 2Cl - → Hg 2 Cl 2

Hg 2 Cl 2 + Sn 2+ +2Cl - → 2Hg + Sn 4+ + 4Cl -

Mercury(II) ions interfere with the determination because they have a similar effect.

Methodology: Place 2-3 drops of a solution of mercury (I) nitrate into a test tube, add 2-3 drops of a solution of tin (II) chloride. A white precipitate is released, which gradually darkens.

2. Reduction of 2+ ions by metallic copper.

Method: A drop of mercuric (I) nitrate solution is applied to a copper plate cleaned with emery. After some time, a gray spot of amalgam appears, which, after removing the solution and wiping the surface with filter paper, becomes shiny:

2+ + Cu → Cu 2+ + 2 Hg

Mercury(II) salts have a similar effect.

3. Reaction with potassium iodide.

Potassium iodide forms a precipitate of Hg 2 I 2 with mercury (I) cations:

2+ + 2I - → Hg 2 I 2

The precipitate is soluble in excess of the reagent to form potassium tetraiodohydrargyrate (II) and a black precipitate of metallic mercury:

Hg 2 I 2 + 2 I - → 2- + Hg

Methodology: Place 2-3 drops of mercury (I) nitrate solution into a test tube and add 2-3 drops of potassium iodide solution. A dirty green precipitate of Hg 2 I 2 is released. Excess reagent is added to the resulting precipitate. Observe the events occurring.

4. Reaction with potassium chromate.

Potassium chromate K 2 CrO 4 forms with cations 2+ a red precipitate of mercury (I) chromate, soluble in nitric acid:

2+ + CrO 4 2- → Hg 2 CrO 4

Methodology: 2-3 drops of a solution of mercury (I) nitrate Hg 2 (NO 3) 2 are placed in a test tube. Add 2-3 drops of potassium chromate. A red precipitate forms.

5. Reaction with ammonium hydroxide solution.

Mercury(I) nitrate reacts with ammonium hydroxide solution to form a dark precipitate of a mixture of metallic mercury and NO 3

2 Hg 2 2+ + NO 3 - + 4 NH 3 + H 2 O → NO 3 + 2 Hg + 3 NH 4 +

Methodology: Place 2-3 drops of mercury (I) nitrate solution in a test tube, add 6 drops of ammonium hydroxide solution. A black precipitate forms.

Analytical reactions of the Pb 2+ cation.

1. Reaction with sulfuric acid or soluble sulfates.

Sulfuric acid or soluble sulfates precipitate lead cations as a white precipitate of lead sulfate. The precipitate dissolves when heated in solutions of alkali metal hydroxides to form hydroxo complexes:

Pb 2+ + SO 4 2- → PbSO 4

PbSO 4 + 4 NaOH → Na 2 + Na 2 SO 4

Lead sulfate is also soluble in a 30% solution of ammonium acetate:

PbSO 4 + CH 3 COO - → + + SO 4 2-

Methodology: Place 5 drops of lead nitrate solution into a test tube, add an equal volume of sodium sulfate or potassium sulfate solution, and a white precipitate forms. Divide the sediment into 2 parts. Sodium or potassium hydroxide is added to one, and a 30% ammonium acetate solution is added to the other. The precipitate dissolves in both cases.

2. Reaction with potassium chromate.

Potassium chromate with lead cations forms a yellow crystalline precipitate PbCrO 4, soluble in alkali metal hydroxides, but insoluble in acetic acid:

Pb 2+ + CrO 4 2- → PbCrO 4

PbCrO 4 + 4OH - → 2- + CrO 4 2-

Method: Place 2-3 drops of lead salt solution into a test tube, add 3 drops of potassium chromate solution. A yellow crystalline precipitate forms. Check its solubility in acetic acid and sodium or potassium hydroxide.

3. Reaction with potassium or sodium iodide, “golden shower” reaction (PHARMACOPOEIAL).

Sodium or potassium iodide with lead cations forms a yellow crystalline precipitate of PbI 2, soluble in excess of the reagent to form the complex compound potassium tetraiodoplumbate (II):

Pb 2+ + 2 I - → PbI 2

PbI 2 + 2I - → 2-

Lead iodide is soluble in hot water and acetic acid. The solubility of the precipitate in hot water is used as an additional reaction for the detection of lead cations, since when the solution is cooled, a precipitate of lead iodide precipitates in the form of golden flakes. The reaction is specific.

Methodology: Place 3-5 drops of lead nitrate salt solution into a test tube, add 3 drops of potassium or sodium iodide solution. A yellow precipitate forms. Add a few drops of water and heat. The precipitate dissolves. Quickly cool the test tube under running cold tap water. A precipitate appears again in the form of shiny golden crystals.

Analytical reactions of group II cations

Silver is quite heavy (ρ = 10.5 g/cm3), shiny (light reflection coefficient is close to 100%), silvery-white metal, malleable and ductile (1 g of silver can stretch the thinnest wire almost 2 km long!), It is the best conductor of heat among metals (that’s why a silver spoon in a glass of hot tea heats up quickly) and electricity. Melting point 962°C.

Application

Silver has been known since ancient times. This is due to the fact that at one time silver, as well as gold, was found in its native form - it did not have to be smelted from ores.

In the old days, coins, vases, jewelry were made from it, and clothes were decorated with the finest silver threads. Now the use of silver is not limited to jewelry - it is used in the production of highly reflective mirrors (inexpensive mirrors are coated with aluminum), electrical contacts, batteries, used in dentistry, and used in gas mask filters, as a disinfectant for water disinfection. Some time ago, solutions of colloidal silver - protargol and collargol - were used to treat colds.

Silver iodide (AgI) is used for climate control (“cloud clearing”). The crystal lattice of silver iodide is very similar in structure to the lattice of ice, so the introduction is not large quantity iodide causes the formation of condensation centers in the clouds, thereby causing precipitation.

Silver is registered as food additives E-174.

Silver is used to make electrodes for powerful zinc-silver batteries. Thus, the batteries of the sunken American submarine Thrasher contained three tons of silver. The high thermal conductivity and chemical inertness of silver are used in electrical engineering: electrical contacts are made from silver and its alloys, and wires in critical devices are coated with silver. Dentures are made from silver-palladium alloy (75% Ag).

Huge amounts of silver used to be used to make coins. Nowadays, mainly commemorative and commemorative coins are made from silver. A lot of silver is spent to make jewelry and cutlery. On such products, as a rule, they put a test indicating the mass of pure silver in grams per 1000 g of alloy (modern test), or the number of spools in one pound of alloy (pre-revolutionary test). 1 pound contains 96 spools, therefore, for example, the old standard 84 corresponds to the modern one [(84/96) 1000] = 875. Soviet rubles and fifty dollars had a standard 900. Modern silver products can have a standard 960, 925, 916, 875, 800 and 750.

Silver compounds are often unstable to heat and light. The discovery of the photosensitivity of silver salts led to the advent of photography and a rapid increase in the demand for silver. Back in the mid-20s, about 10,000 tons of silver were mined annually around the world, and much more was spent (the deficit was covered by old reserves). Repression black and white photographs and color films has significantly reduced silver consumption.

“Silver does not oxidize in air,” wrote D.I. Mendeleev in his textbook “Fundamentals of Chemistry,” “and therefore is classified as a so-called noble metal.” But although silver does not react directly with oxygen, it can dissolve significant amounts of this gas. Even solid silver at a temperature of 450° C can absorb five times the volume of oxygen. Significantly more oxygen (up to 20 volumes per 1 volume of silver) dissolves in the liquid metal.

This property of silver leads to the beautiful (and dangerous) phenomenon of silver spattering, which has been known since ancient times. If molten silver has absorbed significant amounts of oxygen, then the solidification of the metal is accompanied by the release of large amounts of gas. The pressure of the released oxygen causes the crust on the surface of the solidifying silver to rupture, often with great strength. The result is a sudden explosive spatter of metal.

At 170° C, silver in air is covered with a thin film of Ag 2 O oxide, and under the influence of ozone higher oxides are formed (for example, Ag 2 O 3). But silver is especially “afraid” of iodine (iodine tincture) and hydrogen sulfide. Over time, silver items often become tarnished and may even turn black. The reason is the action of hydrogen sulfide. Its source can be not only rotten eggs, but also rubber, some polymers and even food. In the presence of moisture, silver easily reacts with hydrogen sulfide to form a thin film of Ag 2 S sulfide on the surface; due to surface irregularities and the play of light, such a film sometimes appears iridescent. Gradually the film thickens, darkens, turns brown, and then black.

One of the important areas of use of silver was medicine. The ancient Egyptians, for example, applied a silver plate to wounds to ensure rapid healing. The Persian king Cyrus transported water only in silver vessels during his military campaigns. The famous medieval physician Paracelsus treated some diseases with AgNO 3 - silver nitrate (lapis). This remedy is still used in medicine today.

Relatively recently, studies of body cells for silver content led to the conclusion that it is elevated in brain cells.

The bactericidal effect of small concentrations of silver on drinking water. At a content of 0.05 mg/l, water can be drunk without harm to health. Its taste does not change. (For cosmonauts’ drinking, the Ag+ concentration is allowed up to 0.1 – 0.2 mg/l.).

To disinfect water in swimming pools, it was proposed to saturate it with silver bromide. A saturated AgBr solution contains 0.08 mg/l, which is harmless to human health, but harmful to microorganisms and algae.

However, as often happens, what is beneficial in small doses is harmful in large ones. Ag is no exception.

Silver, when excessively introduced into the body, causes a decrease in immunity, changes in the tissues of the brain and spinal cord, and leads to diseases of the liver, kidneys, and thyroid gland. Cases described serious violation mental health in people poisoned with silver preparations. Fortunately, after 1-2 weeks only 0.02 - 0.1% of the injected silver remains in our body, the rest is excreted from the body.

After many years of working with silver and its salts, when they enter the body for a long time, But small doses, an unusual disease may develop - argyria. Silver entering the body can be slowly deposited as metal in the connective tissue and capillary walls of various organs, including the kidneys, bone marrow, and spleen. Accumulating in the skin and mucous membranes, silver gives them a gray-green or bluish color, especially strong on open areas of the body exposed to light. Occasionally, the coloring can be so intense that the skin resembles the skin of blacks.

Argyria develops very slowly, its first signs appear after 2–4 years of continuous work with silver, and severe darkening of the skin is observed only after decades. The lips, temples and conjunctiva of the eyes darken first, then the eyelids. The mucous membranes of the mouth and gums, as well as the sockets of the nails, can be severely stained. Sometimes argyria appears as small blue-black spots. Once it appears, argyria does not disappear, and the skin cannot be returned to its original color. Apart from purely cosmetic inconveniences, a patient with argyria may not experience any painful sensations or health disorders (if the cornea and lens of the eye are not affected); in this regard, argyria can be called a disease only conditionally. This disease also has its own “spoon of honey” - with argyria there are no infectious diseases: a person is so “impregnated” with silver that it kills all pathogenic bacteria that enter the body.

Silver in nature

This beautiful metal has been known to people since ancient times. Silver products found in Western Asia are more than 6 thousand years old. The world's first coins were made from an alloy of gold and silver (electrum). And for several millennia, silver was one of the main coin metals.

The Ore Mountains, the Harz, and the mountains of Bohemia and Saxony located in Central Europe were especially rich in silver. Millions of coins were minted from silver mined near the city of Joachimsthal (now Jáchymov in the Czech Republic). At first they were called “Joachimsthalers”; then the name was shortened to “taler” (in Russia, the first part of the word is “efimka”). These coins were in use throughout Europe, becoming the most common silver coin in history. The name of the dollar comes from the thaler.

After the discovery of America, many silver nuggets were found in the territory of modern Peru, Chile, Mexico, and Bolivia. Thus, a nugget in the form of a plate weighing 1420 kg was discovered in Chile. Many elements have “geographical” names, but Argentina is the only country named after an already known element. The last of the largest silver nuggets were found already in the 20th century in Canada (Ontario). One of them, called the “silver sidewalk,” was 30 m long and went 18 m deep into the earth. When pure silver was smelted from it, it turned out to be 20 tons!

Native silver is rarely found; The bulk of silver in nature is concentrated in minerals, the main one being argentite Ag 2 S. Even more silver is scattered among various rocks.

When describing any element, it is customary to indicate its discoverer and the circumstances of its discovery. Humanity does not have such data about element No. 47. People began to use silver even when there were no scientists.

The Latin name for silver Argentum comes from the Greek “argos” - white, shiny. The Russian word “silver,” according to scientists, comes from the word “sickle” (sickle of the moon). The shine of silver also reminded alchemists of the moon, who used the sign of the moon as a symbol of the element.

Silver and glass. These two substances are found not only in the production of mirrors. Silver is needed to make signal glasses and light filters. A small addition (0.15 - 0.20%) of silver nitrate (or silver nitrate) gives the glass an intense golden-yellow color. Orange glass is obtained by introducing gold and silver into the glass melt at the same time.

Silver resists the action of alkalis better than many other metals. That is why the walls of pipelines, autoclaves, reactors and other apparatus in the chemical industry are coated with silver as a protective metal.

And in terms of sonority, silver stands out noticeably among other metals. It is not for nothing that silver bells appear in many fairy tales. Bellmakers have long added silver to bronze “for a crimson ringing.” Nowadays, the strings of some musical instruments are made from an alloy that contains 90% silver.

If the silver has turned black...

During long-term storage, silver items become dull and become covered with a thin layer of silver sulfide Ag 2 S. To restore the item to its former shine, it is necessary to remove the sulfide film. This can be done in several ways.

1) Mix water, ammonia and tooth powder in the form of gruel. Apply this product to soft cloth and clean the products until the darkening is removed.

2) Boil the silver item (about 20 minutes) in water with the addition of baking soda and pieces aluminum foil or wire (or in an aluminum container).

3) Regular tooth powder or toothpaste are still not inferior to any of the newest means. By rubbing the product with a former toothbrush, you will restore its original shine.

No matter what product you choose to clean your products, be sure to rinse them thoroughly after the procedure and wipe them dry with a cloth.

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(chemical condensation method)

AgNO 3 + KI = AgI + KNO 3

The dispersion medium is water, it contains AgI particles. If we take a strictly equivalent amount of reagents (without excess or deficiency), then an AgI precipitate will form (since AgI is insoluble in water).

How is precipitate formed?

After mixing the two solutions, AgI molecules are formed throughout the volume. Next, the nearest molecules upon collision stick together and larger ones are formed.

These particles are coarse and they precipitate. To form a colloidal solution, the substance must be in excess.

A) if there is an excess of AgNO 3, then after the reaction AgI molecules are formed throughout the entire volume, and K + ions remain in the solution; NO 3 - ; Ag+. When colloidal particles are formed in a solution, the adsorption of Ag ions immediately begins on them - specific adsorption: those ions that are already in the lattice are adsorbed on the surface of the solid. Equilibrium quickly occurs in the solution, since positively charged Ag ions attached to the colloidal AgI particle do not allow AgI molecules to disintegrate, and also do not allow other Ag ions to attach. Then this complex particle with a positive charge begins to attach NO 3 - ions. But the adsorption of NO 3 - does not completely neutralize the charge, since these ions are not enough to completely neutralize the positive charge of the particle (per 100 Ag + 92 NO 3 -). NO 3 - ions are attracted by the particle and are retained near it in the diffusion layer. Not every force is capable of attracting and holding an ion as it moves.

Thus, in a solution, all colloidal particles are charged with the same charge, and when they collide, they repel. In this case, excess AgNO 3 acts as a stabilizer.

Structure of a colloidal particle

(n Ag + (n-x) NO 3 - ) x+ x NO 3 -

– what substance does the colloidal particle consist of?

m is the number of particles that make up the colloidal particle.

n is the number of ions adsorbed first on the surface of charged particles (usually n=100). This potential-determining ions or adsorption layer of counterions.

x – number of ions of the opposite sign located in the diffusion layer of the particle (usually x=8) – diffuse layer of counterions.

(n-x) – number of particles on the surface.

– the nucleus, the particle itself.

( ) – core with ions adsorbed on the surface; granule.

All together - a colloidal micelle.

B) if there is an excess of KI, then the formula of the colloidal particle is

(n I - (n-x) K + ) x- x K +

I - are adsorbed on the surface of the colloidal particle, and K + ions are adsorbed in the diffusion layer. A colloidal solution with negatively charged particles is formed. To obtain a colloidal solution, the initial concentrations of substances must be moderate (0.001 N). If the solutions are concentrated, a precipitate will form.

AgNO 3 + KI = AgI ↓+ KNO 3

At high concentrations, many AgI molecules are formed, that is, many colloidal particles, the distance between them is small, Ag + ions do not have time to be adsorbed and the particles stick together. If the concentration is too low, then there are few AgI particles and there will not be enough of them to grow to the size of a colloidal particle. The required concentration is determined experimentally. To verify the colloidality of a solution, various methods are used.

1. filtration– coarse solutions do not pass through the filter completely

2. opalescence– change in the color of the solution when light falls from different viewing angles. This phenomenon is typical only for colloidal solutions.