diethyl sulfate(C 2 H 5) 2 SO 4 . Melting point -26°C, boiling point 210°C, soluble in alcohols, insoluble in water. Obtained by the interaction of sulfuric acid with ethanol. It is an ethylating agent in organic synthesis. Penetrates through the skin.

dimethyl sulfate(CH 3) 2 SO 4 . Melting point -26.8°C, boiling point 188.5°C. Let's dissolve in alcohols, it is bad - in water. Reacts with ammonia in the absence of a solvent (explosively); sulfonates some aromatic compounds, such as phenol esters. Obtained by interaction of 60% oleum with methanol at 150°C. It is a methylating agent in organic synthesis. Carcinogen, affects the eyes, skin, respiratory organs.

Sodium thiosulfate Na 2 S 2 O 3

Na 2 SO 3 + S \u003d Na 2 S 2 O 3

Like thiosulfuric acid, it is a strong reducing agent. It is easily oxidized by chlorine to sulfuric acid:

Na 2 S 2 O 3 + 4Cl 2 + 5H 2 O \u003d 2H 2 SO 4 + 2NaCl + 6HCl

The use of sodium thiosulfate to absorb chlorine (in the first gas masks) was based on this reaction.

Sodium thiosulfate is oxidized somewhat differently by weak oxidizing agents. In this case, salts of tetrathionic acid are formed, for example:

2Na 2 S 2 O 3 + I 2 \u003d Na 2 S 4 O 6 + 2NaI

Sodium thiosulfate is a by-product in the production of NaHSO 3 , sulfur dyes, in the purification of industrial gases from sulfur. It is used to remove traces of chlorine after bleaching fabrics, to extract silver from ores; is a fixer in photography, a reagent in iodometry, an antidote for poisoning with arsenic, mercury compounds, an anti-inflammatory agent.

We take sodium thiosulfate and three acids (sulfuric, hydrochloric and orthophosphoric):

Na2S2O3 + H2SO4 = Na2SO4 + SO2 + S + H2O

Na2S2O3 + 2 HCl = 2 NaCl + SO2 + S + H2O

3 Na2S2O3 +2 H3PO4 = 2 Na3PO4 + 3 SO2 + 3 S + 3 H2O

Pour into three test tubes 8 ml of sodium thiosulfate solution. Pour 8 ml of sulfuric acid into the first test tube with a solution of sodium thiosulfate, mix quickly and note the time in seconds from the start of the reaction to the cloudiness of the solution. To better notice the end of the reaction, glue a strip of black paper on the opposite side of the test tube wall. We finish the time report at the moment when this strip is not visible through the cloudy solution.

Similarly, we carry out experiments with other acids. The results are entered in the table (Appendix 1, Table 1). The reaction rate is defined as a value inversely proportional to time: υ = 1/ t. Based on the table, we build a graph of the dependence of the reaction rate on the nature of the reactants (Appendix 2, graph 1).

Conclusion: thus, the nature of acids affects the rate of a chemical reaction. And, since the strength of acids is determined by the concentration of hydrogen ions, the reaction rate also depends on the concentration of the reactants.

B. Consider the reaction of the interaction of various metals with hydrochloric acid. The reaction rate will be determined by the volume of released hydrogen, which is collected by the method of water displacement (Appendix 3, Figure 1).

In four test tubes we place 0.05 g of metals: magnesium, zinc, iron and copper. In turn, pour equal volumes of hydrochloric acid (1:2) into each test tube (a). Hydrogen, which will be rapidly depleted, will enter the test tube (b). Note the time it takes for the tube to fill with hydrogen. Based on the results (Appendix 4, Table 2), we build a graph depending on the nature of the reactants (Appendix 4, Chart 2).

Conclusion: not all metals can interact with acids by removing hydrogen. Metals that displace hydrogen from acid solutions are located in the series N.N. Beketov to hydrogen, and metals that do not displace hydrogen - after hydrogen (in our case, this is copper). But the first group of metals also differ in the degree of activity: magnesium-zinc-iron, therefore, the intensity of hydrogen evolution is different.

Thus, the rate of a chemical reaction depends on the nature of the reactants.

2. Dependence of the rate of a chemical reaction on the concentration of interacting substances.

Target. Establish a graphical dependence of the effect of concentration on the reaction rate.

For the experiment, we use the same solutions of sodium thiosulfate and sulfuric acid that were used in the first experiment (A).

Pour the indicated amounts of milliliters of sodium thiosulfate solution and water into numbered test tubes. Pour 8 ml of sulfuric acid solution into the first test tube, mix quickly and note the time from the beginning of the reaction to the cloudiness of the solution (see experiment 1 A). We carry out similar experiments with the rest of the test tubes. We enter the results in a table (Appendix 6, Table 3), on the basis of which we build a graph of the dependence of the rate of a chemical reaction on the concentration of reactants (Appendix 7, Chart 3). We obtained a similar result by keeping the concentration of sodium thiosulfate constant, but changing the concentration of sulfuric acid.

Conclusion: thus, the rate of a chemical reaction depends on the concentration of the reacting substances: the higher the concentration, the greater the reaction rate.

3. Dependence of the rate of a chemical reaction on temperature.

Purpose: To test whether the rate of a chemical reaction depends on temperature.

We carry out the experiment with solutions of sodium thiosulfate and sulfuric acid (see experiment 1), additionally we prepare a beaker, a thermometer.

Pour 8 ml of sodium thiosulfate solution into four test tubes, 8 ml of sulfuric acid solution into 4 others. We place all test tubes in a glass of water and measure the temperature of the water. After 5 minutes, we take out two test tubes with solutions of sodium thiosulfate and sulfuric acid, drain them, mix and note the time until the solution becomes cloudy. We heat a glass with water and test tubes by 10 ° C and repeat the experiment with the next two test tubes. We carry out the same experiments with the rest of the test tubes, each time increasing the water temperature by 10°C. The results obtained are recorded in a table (Appendix 8, Table 4) and we plot the dependence of the reaction rate on temperature (Appendix 9, Chart 4).

Conclusion: this experiment led to the conclusion that the rate of a chemical reaction increases with an increase in temperature for every 10°C by 2–4 times, i.e. proved the validity of Van't Hoff's law.

4. Effect of a catalyst on the rate of a chemical reaction.

Purpose: to check whether the rate of a chemical reaction depends on the catalyst, and whether the catalysts have specificity.

A. To test the specificity of the catalyst, we used the decomposition reaction of hydrogen peroxide: 2H2O2 = 2H2O + H2. They took a 3% solution, the decomposition of hydrogen peroxide is very weak, even a smoldering splinter dropped into a test tube does not flare up. We used silicon dioxide SiO2, manganese dioxide MnO2, potassium permanganate KMnO4, and sodium chloride NaCl as catalysts. Only when manganese (IV) oxide powder was added did a rapid evolution of oxygen occur, a smoldering splinter, lowered into a test tube, flared up brightly.

Thus, catalysts are substances that speed up a chemical reaction, and, most often, a particular reaction requires its own catalyst.

5. Kinetics of the catalytic decomposition of hydrogen peroxide.

Purpose: to find out the dependence of the reaction rate on the concentration of substances, temperature and catalyst.

The decomposition of a very weak solution of hydrogen peroxide begins under the influence of a catalyst. With the course of the reaction, the concentration of hydrogen peroxide decreases, as can be judged by the amount of oxygen released per unit time. We carry out the experiment in the device (Appendix 10, Figure 2): put 0.1 g of manganese dioxide powder into a test tube, attach it to a rubber tube, pour 40 ml of a 3% hydrogen peroxide solution into the flask, connect it with a test tube using a rubber tube. We fill the cylinder (burette) with water, lower it into the crystallizer, fix it vertically in the clamp of the tripod, and bring the gas outlet tube from the Wurtz flask under it. Without a catalyst, oxygen evolution is not observed. After adding manganese dioxide, every minute for 10 minutes we note and write down in the table the volume of released oxygen (Appendix 11, Table 5). Based on the data, we build a graph of the dependence of the volumes of released oxygen on time (Appendix 12, graph 5)

6. Influence of the contact surface of reactants on the rate of a chemical reaction.

Target. Find out whether the contact surface of the reactants affects the rate of a heterogeneous chemical reaction.

The same amount (0.5 g) of chalk (CaCO3) in the form of a piece and powder was weighed on a balance, the weighed portions were placed in two test tubes, into which the same amount of hydrochloric acid (1:2) was poured. We observe the release of carbon dioxide, and in the first test tube (chalk in the form of a piece) the reaction is less vigorous than in the second (chalk in the form of powder) (Appendix 13, photos 1.2): CaCO3 + 2 HCl = CaCl2 + CO2 + H2O

An observable sign of the reaction is the formation of a white-yellow turbidity (insoluble sulfur). Thiosulfuric acid is unstable (see the reaction equation!), so it is obtained by reacting sodium thiosulfate with dilute sulfuric acid:

Na 2 S 2 O 3 + H 2 SO 4 \u003d H 2 S 2 O 3 + Na 2 SO 4

those. overall reaction:

Na 2 S 2 O 3 + H 2 SO 4 \u003d S + SO 2 + H 2 O + Na 2 SO 4

Carrying out the reaction: Pour 20 ml of 2M sulfuric acid into 2 identical glasses. In 1 of the glasses add 80 ml of water (reduce the concentration of acid). Simultaneously pour into both beakers (from 2 other beakers or cylinders) 20 ml of 2M sodium thiosulfate.

What to watch: In which of the beakers is the turbidity formed faster?

Catalysis

At the heart of the experiment hydrogen peroxide decomposition reaction

H 2 O 2 \u003d H 2 O + 1 / 2O 2

accelerating in the presence of manganese dioxide, as well as some salts of heavy metals, the enzyme catalase, etc. An observed sign of the reaction is the release of gas bubbles, in which a smoldering torch flares brightly.

Carrying out the reaction: Pour 10 ml of 30% H 2 O 2 into a tall cylinder (per 100 ml). Quickly pour in the MnO 2 powder (an option is to drop a few drops of blood). Insert a smoldering torch into the cylinder.

Catalysis

At the heart of the experiment catalytic oxidation of ammonia on chromium oxide.

4NH 3 + 5O 2 \u003d 4NO + 6H 2 O

The observed sign of the reaction is sparks (heating of chromium oxide particles due to the exothermic thermal effect of the reaction and their glow).

Carrying out the reaction: Thoroughly rinse a large flat-bottomed flask (500 ml) from the inside with concentrated ammonia solution (thus creating a high concentration of ammonia vapors in it). Throw chromium oxide (III) heated in an iron spoon into it.

A simple model experiment, on several topics at once.

In a dry beaker (simple disposable drinking cups can be used), place equal amounts (about the size of a pea each) of dry citric acid and baking soda (sodium bicarbonate).

The reaction does not proceed without water, and when a few drops of water are added, the mixture "boils".

NaHCO 3 + H 3 (C 5 H 5 O 7) = Na 3 (C 5 H 5 O 7) + CO 2 + H 2 O

You can carry out the same reaction by replacing soda with chalk. This proves that the reaction is reduced to the interaction of a carbonate ion with a proton:

CO 3 2- + 2H + = H 2 CO 3 = CO 2 + H 2 O

Then, in one glass, we prepare a saturated solution of soda (its solubility is 9.6 g per 100 g of water at room temperature). In two other glasses we put citric acid - in the first volume with a match head, in the second about 5 times more. Pour 10 ml of water into both glasses and dissolve the acid with stirring. In both glasses with citric acid, simultaneously add 5 ml of a saturated sodium bicarbonate solution. It can be seen that in a glass, where the concentration of citric acid is higher, gas evolution is more intense. Conclusion: the reaction rate is proportional to the concentration of the reactants.

Thiosulfuric acid. sodium thiosulfate. Obtaining, properties, application.

Sulfuric acid esters include dialkyl sulfates (RO2)SO2. These are high-boiling liquids; the lower ones are soluble in water; in the presence of alkalis, they form alcohol and salts of sulfuric acid. Lower dialkyl sulfates are alkylating agents.
Diethyl sulfate (C2H5)2SO4. Melting point -26°C, boiling point 210°C, soluble in alcohols, insoluble in water. Obtained by the interaction of sulfuric acid with ethanol. It is an ethylating agent in organic synthesis. Penetrates through the skin.
Dimethyl sulfate (CH3)2SO4. Melting point -26.8°C, boiling point 188.5°C. Let's dissolve in alcohols, it is bad - in water. Reacts with ammonia in the absence of a solvent (explosively); sulfonates some aromatic compounds, such as phenol esters. Obtained by interaction of 60% oleum with methanol at 150°C. It is a methylating agent in organic synthesis. Carcinogen, affects the eyes, skin, respiratory organs.
Sodium thiosulfate Na2S2O3

Salt of thiosulfuric acid, in which two sulfur atoms have different oxidation states: +6 and -2. Crystalline substance, highly soluble in water. It is produced in the form of Na2S2O3 5H2O crystalline hydrate, commonly called hyposulfite. Obtained by the interaction of sodium sulfite with sulfur during boiling:
Na2SO3+S=Na2S2O3
Like thiosulfuric acid, it is a strong reducing agent. It is easily oxidized by chlorine to sulfuric acid:
Na2S2O3+4Сl2+5Н2О=2H2SO4+2NaCl+6НCl
The use of sodium thiosulfate to absorb chlorine (in the first gas masks) was based on this reaction.
Sodium thiosulfate is oxidized somewhat differently by weak oxidizing agents. In this case, salts of tetrathionic acid are formed, for example:
2Na2S2O3+I2=Na2S4O6+2NaI
Sodium thiosulfate is a by-product in the production of NaHSO3, sulfur dyes, in the purification of industrial gases from sulfur. It is used to remove traces of chlorine after bleaching fabrics, to extract silver from ores; is a fixer in photography, a reagent in iodometry, an antidote for poisoning with arsenic, mercury compounds, an anti-inflammatory agent.

Thiosulfuric acid- an inorganic compound, a dibasic strong acid with the formula H 2 SO 3 S. A colorless viscous liquid that reacts with water. Forms salts - inorganic thiosulfates. Thiosulfuric acid contains two sulfur atoms, one of which has an oxidation state of +4, and the second is electrically neutral.

Receipt

The reaction of hydrogen sulfide and sulfur trioxide in ethyl ether at low temperatures:

The action of gaseous hydrogen chloride on sodium thiosulfate:

Physical Properties

Thiosulfuric acid forms a colorless viscous liquid that does not freeze even at very low temperatures. Thermally unstable - decomposes already at room temperature.

Quickly, but not instantly, decomposes in aqueous solutions. In the presence of sulfuric acid, it decomposes instantly.

Chemical properties

Thermally very unstable:

Decomposes in the presence of sulfuric acid:

Reacts with alkalis:

Reacts with halogens:

Forms esters - organic thiosulfates.

Sodium thiosulfate (antichlor, hyposulfite, sodium sulfidotrioxosulfate) - Na 2 S 2 O 3 or Na 2 SO 3 S, a salt of sodium and thiosulfuric acid, forms a crystalline Na 2 S 2 O 3 5H 2 O.

Receipt

Oxidation of Na polysulfides;

Boiling excess sulfur with Na 2 SO 3:

The interaction of H 2 S and SO 2 with NaOH (a by-product in the production of NaHSO 3, sulfur dyes, in the purification of industrial gases from S):

Boiling excess sulfur with sodium hydroxide:

then, according to the above reaction, sodium sulfide adds sulfur, forming sodium thiosulfate.

At the same time, during this reaction, sodium polysulfides are formed (they give the solution a yellow color). For their destruction, SO 2 is passed into the solution.

Pure anhydrous sodium thiosulfate can be obtained by reacting sulfur with sodium nitrite in formamide. This reaction proceeds quantitatively (at 80 °C in 30 minutes) according to the equation:

Dissolution of sodium sulfide in water in the presence of atmospheric oxygen:

Physical and chemical properties

Colorless monoclinic crystals. Molar mass 248.17 g/mol (pentahydrate).

Soluble in water (41.2% at 20°C, 69.86% at 80°C).

At 48.5 °C, the crystalline hydrate dissolves in its water of crystallization, forming a supersaturated solution; dehydrated at about 100°C.

When heated to 220 ° C, it decomposes according to the scheme:

Sodium thiosulfate is a strong reducing agent:

With strong oxidizing agents, such as free chlorine, it oxidizes to sulfates or sulfuric acid:

Weaker or slower acting oxidizing agents, for example, iodine, are converted into salts of tetrathionic acid:

The above reaction is very important, as it serves as the basis of iodometry. It should be noted that in an alkaline medium, sodium thiosulfate can be oxidized with iodine to sulfate.

It is impossible to isolate thiosulfuric acid (hydrogen thiosulfate) by the reaction of sodium thiosulfate with a strong acid, since it is unstable and immediately decomposes:

Molten hydrated Na 2 S 2 O 3 ·5H 2 O is very prone to supercooling.

Application

for removing traces of chlorine after bleaching fabrics

for the extraction of silver from ores;

fixer in photography;

Reagent in iodometry

antidote for poisoning: As, Br, Hg and other heavy metals, cyanides (translates them into thiocyanates), etc.

for intestinal disinfection;

for the treatment of scabies (together with hydrochloric acid);

Anti-inflammatory and anti-burn agent;

can be used as a medium for determining molecular weights by freezing point depression (cryoscopic constant 4.26°)

Registered in the food industry as a food additive E539.

admixtures for concrete.

for cleansing tissues from iodine

· Gauze bandages impregnated with a solution of sodium thiosulfate were used to protect the respiratory organs from the poisonous substance chlorine in the First World War.

Lecturer: Korableva A.A.

REPORT

ABOUT LABORATORY WORK

COURSE: GENERAL CHEMISTRY

" REACTION RATE IN SOLUTIONS "

OF 62 5528 1.04 LR

I've done the work

group student

Saint Petersburg

Goal of the work:

Determine the rate constant, temperature coefficient, activation energy of the reaction of interaction of sodium thiosulfate with sulfuric acid.

This laboratory work studies the reaction between sodium thiosulfate (hyposulfite) Na2S2O3 and sulfuric acid H2SO4.

This reaction proceeds in two stages:

1) (fast)

The first stage of ion exchange proceeds almost instantaneously. Thiosulfuric acid is an unstable compound that decomposes with the release of a white sulfur precipitate.

2) (slowly)

The reaction rate can be judged by the appearance of opalescence and further turbidity of the solution from precipitated sulfur.

The overall reaction is determined by the second stage of the process and depends on the concentration of H2SO4, and hence Na2S2O3 (the reaction is pseudomolecular).

The kinetic equation has the form:

Instruments and reagents:

Thermostats, thermometers, measuring cylinders, test tubes, test tube holders, stopwatch, Na2S2O3 and H2SO4 solutions.

Experience #1:

Effect of thiosulfate on the rate of a chemical reaction.

Dependence of the reaction rate on the concentration of sodium thiosulfate.

Processing the results of the experience:

We calculate the relative reaction rate using the formula:

2. Based on the kinetic equation, we determine the value of the reaction rate constant:

3. Determine the average value of the constant for a given room temperature, in this case T = 14 degrees Celsius.

4
. Express the dependence of the reaction rate on the concentration of thiosulfate - graphically. (see fig. No. 1).

5. Graphically, we determine the reaction rate constant as the tangent of the slope of the straight line OA to the abscissa axis. We compare a graphically defined constant with its analytical value.

KGR = tg = 0.162 KSR = 0.17 KGR  KSR

Experience #2:

Effect of temperature on the rate of a chemical reaction.

experience temperature, T, degrees Celsius.	reactions t, s	Relates speed react. V, 1/s	Const. speed react. K, l/mol*s

Processing the results of the experience:

1. Calculate the relative reaction rate at each temperature:

See the results in the table above.

2. Based on the kinetic equation, we determine the value of the constant for each temperature:

R
See the results in the table above.

3. We express graphically the effect of temperature on the rate of a chemical reaction. (see Fig. No. 2).

4. Based on the Van Hoff equation, we determine the value of the temperature coefficient for each temperature interval and calculate its average value:

K2/K1 = 1 = 2.42

K3/K2 = 2 = 1.97 medium = 2.3

K4/K3 = 3 = 2.49

5
. Based on the Arrhenius equation, we calculate the analytical value of the activation energy for each temperature interval:

E
a1 = 61785 J/mol Ea2 = 50729 J/mol Ea3 = 72882 J/mol

And we calculate its average value:

EaAVG = 61798 J/mol

6. We build a graphical dependence of lgK on 1/T according to the calculated rate constants at different temperatures and determine the activation energy graphically (see Fig. No. 3).

tg = - Еа / 2.3 R, therefore

ЕаГР = -2.3 R tg = -2.3 * 8.3 * tg = 19.09* 3230 = 61660 J/mol

7. Compare the activation energy values obtained graphically and analytically:

EaGR = 61660 J/mol EaAVED = 61798 J/mol EaGR  EaGR

Conclusion:

At a temperature equal to const, the rate of a chemical reaction is proportional to the concentration of substances involved in this reaction. (see fig. No. 1)

As the temperature increases, the rate of a chemical reaction increases

Provided that the concentration remains unchanged. This can be explained by the fact that with increasing temperature, the atoms of substances pass into a more excited state, that is, they receive additional energy - the activation energy necessary to break the chemical bond and form a new substance.