OVRs are specially highlighted in color in the article. Pay attention to them Special attention. These equations may appear on the Unified State Exam.
Dilute sulfuric acid behaves like other acids, hiding its oxidative capabilities:
And one more thing to remember about dilute sulfuric acid: she does not react with lead. A piece of lead thrown into dilute H2SO4 becomes covered with a layer of insoluble (see solubility table) lead sulfate and the reaction immediately stops.
Oxidizing properties of sulfuric acid
– heavy oily liquid, non-volatile, tasteless and odorless
Due to sulfur in the oxidation state +6 (highest) sulfuric acid acquires strong oxidizing properties.
Rule for task 24 (old A24) when preparing sulfuric acid solutions You should never pour water into it. Concentrated sulfuric acid should be poured into water in a thin stream, stirring constantly.
Reaction of concentrated sulfuric acid with metals
These reactions are strictly standardized and follow the scheme:
H2SO4(conc.) + metal → metal sulfate + H2O + reduced sulfur product.
There are two nuances:
1) Aluminum, iron And chromium with H2SO4 (conc) in normal conditions do not react due to passivation. Needs to be heated.
2) C platinum And gold H2SO4 (conc) does not react at all.
Sulfur V concentrated sulfuric acid- oxidizer
- This means that it will recover itself;
- the degree of oxidation to which sulfur is reduced depends on the metal.
Let's consider sulfur oxidation state diagram:
- Before -2 sulfur can only be reduced by very active metals - in a series of voltages up to and including aluminum.
The reactions will go like this:
8Li+5H 2 SO 4( conc. .) → 4Li 2 SO 4 + 4H 2 O+H 2 S
4Mg + 5H 2 SO 4( conc. .) → 4MgSO 4 + 4H 2 O+H 2 S
8Al + 15H 2 SO 4( conc. .) (t)→ 4Al 2 (SO 4 ) 3 +12H 2 O+3H 2 S
- upon interaction of H2SO4 (conc) with metals in a series of voltages after aluminum, but before iron, that is, with metals with average activity, sulfur is reduced to 0 :
3Mn + 4H 2 SO 4( conc. .) → 3MnSO 4 + 4H 2 O+S↓
2Cr + 4H 2 SO 4( conc. .) (t)→Cr 2 (SO 4 ) 3 + 4H 2 O+S↓
3Zn + 4H 2 SO 4( conc. .) → 3ZnSO 4 + 4H 2 O+S↓
- all other metals starting with hardware in a number of voltages (including those after hydrogen, except for gold and platinum, of course), they can only reduce sulfur to +4. Since these are low-active metals:
2 Fe + 6 H 2 SO 4(conc.) ( t)→ Fe 2 ( SO 4 ) 3 + 6 H 2 O + 3 SO 2
(note that iron oxidizes to +3, the highest possible oxidation state, since it is a strong oxidizing agent)
Cu+2H 2 SO 4( conc. .) → CuSO 4 + 2H 2 O+SO 2
2Ag + 2H 2 SO 4( conc. .) → Ag 2 SO 4 + 2H 2 O+SO 2
Of course, everything is relative. The depth of recovery will depend on many factors: acid concentration (90%, 80%, 60%), temperature, etc. Therefore, it is impossible to predict products completely accurately. The above table also has its own approximate percentage, but you can use it. It is also necessary to remember that in the Unified State Examination, when the product of reduced sulfur is not indicated and the metal is not particularly active, then, most likely, the compilers mean SO 2. You need to look at the situation and look for clues in the conditions.
SO 2 - this is generally a common product of ORR with the participation of conc. sulfuric acid.
H2SO4 (conc) oxidizes some nonmetals(which exhibit reducing properties), as a rule, to a maximum - the highest degree of oxidation (an oxide of this non-metal is formed). In this case, sulfur is also reduced to SO 2:
C+2H 2 SO 4( conc. .) → CO 2 + 2H 2 O+2SO 2
2P+5H 2 SO 4( conc. .) → P 2 O 5 +5H 2 O+5SO 2
Freshly formed phosphorus oxide (V) reacts with water to produce orthophosphoric acid. Therefore, the reaction is recorded immediately:
2P+5H 2 SO 4( conc. ) → 2H 3 P.O. 4 + 2H 2 O+5SO 2
The same thing with boron, it turns into orthoboric acid:
2B+3H 2 SO 4( conc. ) → 2H 3 B.O. 3 +3SO 2
The interaction of sulfur with an oxidation state of +6 (in sulfuric acid) with “other” sulfur (located in a different compound) is very interesting. Within the framework of the Unified State Examination, the interaction of H2SO4 (conc) is considered with sulfur (a simple substance) and hydrogen sulfide.
Let's start with interaction sulfur (a simple substance) with concentrated sulfuric acid. In a simple substance the oxidation state is 0, in an acid it is +6. In this ORR, sulfur +6 will oxidize sulfur 0. Let's look at the diagram of the oxidation states of sulfur:
Sulfur 0 will oxidize, and sulfur +6 will be reduced, that is, lower the oxidation state. Sulfur dioxide will be released:
2 H 2 SO 4(conc.) + S → 3 SO 2 + 2 H 2 O
But in the case of hydrogen sulfide:
Both sulfur (a simple substance) and sulfur dioxide are formed:
H 2 SO 4( conc. .) +H 2 S → S↓ + SO 2 + 2H 2 O
This principle can often help in determining the product of ORR, where the oxidizing and reducing agent are the same element, in different degrees oxidation. The oxidizing agent and the reducing agent “meet each other halfway” according to the oxidation state diagram.
H2SO4 (conc), one way or another, interacts with halides. Only here you need to understand that fluorine and chlorine are “themselves with a mustache” and ORR does not occur with fluorides and chlorides, undergoes a conventional ion exchange process during which hydrogen halide gas is formed:
CaCl 2 + H 2 SO 4 (conc.) → CaSO 4 + 2HCl
CaF 2 + H 2 SO 4 (conc.) → CaSO 4 + 2HF
But the halogens in the composition of bromides and iodides (as well as in the composition of the corresponding hydrogen halides) are oxidized to free halogens. Only sulfur is reduced in different ways: iodide is a stronger reducing agent than bromide. Therefore, iodide reduces sulfur to hydrogen sulfide, and bromide to sulfur dioxide:
2H 2 SO 4( conc. .) + 2NaBr → Na 2 SO 4 + 2H 2 O+SO 2 +Br 2
H 2 SO 4( conc. .) + 2HBr → 2H 2 O+SO 2 +Br 2
5H 2 SO 4( conc. .) + 8NaI → 4Na 2 SO 4 + 4H 2 O+H 2 S+4I 2 ↓
H 2 SO 4( conc. .) + 8HI → 4H 2 O+H 2 S+4I 2 ↓
Hydrogen chloride and hydrogen fluoride (as well as their salts) are resistant to the oxidizing action of H2SO4 (conc).
And finally, the last thing: this is unique for concentrated sulfuric acid, no one else can do this. She has water-removing property.
This allows concentrated sulfuric acid to be used in a variety of ways:
First, drying of substances. Concentrated sulfuric acid removes water from the substance and it “becomes dry.”
Secondly, a catalyst in reactions in which water is eliminated (for example, dehydration and esterification):
H 3 C–COOH + HO–CH 3 (H 2 SO 4 (conc.)) → H 3 C–C(O)–O–CH 3 + H 2 O
H 3 C–CH 2 –OH (H 2 SO 4 (conc.)) → H 2 C =CH 2 + H 2 O
Sulfurous acid is an inorganic dibasic unstable acid of medium strength. An unstable compound, known only in aqueous solutions at a concentration of no more than six percent. When attempting to isolate pure sulfurous acid, it breaks down into sulfur oxide (SO2) and water (H2O). For example, when concentrated sulfuric acid (H2SO4) reacts with sodium sulfite (Na2SO3), sulfur oxide (SO2) is released instead of sulfurous acid. This is what it looks like this reaction:
Na2SO3 (sodium sulfite) + H2SO4 (sulfuric acid) = Na2SO4 (sodium sulfate) + SO2 (sulfur dioxide) + H2O (water)
Sulfurous acid solution
When storing it, it is necessary to exclude access to air. Otherwise, sulfurous acid, slowly absorbing oxygen (O2), will turn into sulfuric acid.
2H2SO3 (sulfuric acid) + O2 (oxygen) = 2H2SO4 (sulfuric acid)
Solutions of sulfurous acid have a rather specific odor (reminiscent of the odor remaining after lighting a match), the presence of which can be explained by the presence of sulfur oxide (SO2), which is not chemically bound with water.
Chemical properties sulfurous acid
1. H2SO3) can be used as a reducing agent or an oxidizing agent.
H2SO3 is a good reducing agent. With its help, it is possible to obtain hydrogen halides from free halogens. For example:
H2SO3 (sulfuric acid) + Cl2 (chlorine, gas) + H2O (water) = H2SO4 (sulfuric acid) + 2HCl ( hydrochloric acid)
But when interacting with strong reducing agents, this acid will act as an oxidizing agent. An example is the reaction of sulfurous acid with hydrogen sulfide:
H2SO3 (sulfuric acid) + 2H2S (hydrogen sulfide) = 3S (sulfur) + 3H2O (water)
2. The chemical compound we are considering forms two - sulfites (medium) and hydrosulfites (acidic). These salts are reducing agents, just like (H2SO3) sulfurous acid. When they are oxidized, salts of sulfuric acid are formed. When sulfites of active metals are calcined, sulfates and sulfides are formed. This is a self-oxidation-self-healing reaction. For example:
4Na2SO3 (sodium sulfite) = Na2S + 3Na2SO4 (sodium sulfate)
Sodium and potassium sulfites (Na2SO3 and K2SO3) are used in dyeing fabrics in the textile industry, in bleaching metals, and in photography. Calcium hydrosulfite (Ca(HSO3)2), which exists only in solution, is used to process wood material into a special sulfite pulp. It is then used to make paper.
Application of sulfurous acid
Sulfurous acid is used:
For bleaching wool, silk, wood pulp, paper and other similar substances that cannot withstand bleaching with stronger oxidizing agents (for example, chlorine);
As a preservative and antiseptic, for example, to prevent the fermentation of grain when producing starch or to prevent the fermentation process in wine barrels;
To preserve food, for example, when canning vegetables and fruits;
Processed into sulfite pulp, from which paper is then produced. In this case, a solution of calcium hydrosulfite (Ca(HSO3)2) is used, which dissolves lignin, a special substance that binds cellulose fibers.
Sulfurous acid: preparation
This acid can be produced by dissolving sulfur dioxide (SO2) in water (H2O). You will need concentrated sulfuric acid (H2SO4), copper (Cu) and a test tube. Algorithm of actions:
1. Carefully pour concentrated sulfuric acid into a test tube and then place a piece of copper in it. Heat up. The following reaction occurs:
Cu (copper) + 2H2SO4 (sulfuric acid) = CuSO4 (sulfur sulfate) + SO2 (sulfur dioxide) + H2O (water)
2. The flow of sulfur dioxide must be directed into a test tube with water. When it dissolves, it partially occurs with water, resulting in the formation of sulfurous acid:
SO2 (sulfur dioxide) + H2O (water) = H2SO3
So, by passing sulfur dioxide through water, you can get sulfurous acid. It is worth considering that this gas has an irritating effect on the shells respiratory tract, can cause inflammation, as well as loss of appetite. Inhaling it for a long time may cause loss of consciousness. This gas must be handled with extreme caution and care.
Sulfur is an element of the sixth group of the third period of the Mendeleev periodic table. Therefore, the structure of the sulfur atom is depicted as follows:
The structure of the sulfur atom indicates that it is a non-metal, i.e., the sulfur atom is capable of both receiving electrons and giving away electrons:
Task 15.1. Create formulas for sulfur compounds containing sulfur atoms with given oxidation states.
Simple substance " sulfur» - hard brittle mineral yellow color, insoluble in water. In nature, both native sulfur and its compounds are found: sulfides, sulfates. Sulfur, as an active non-metal, easily reacts with hydrogen, oxygen, and almost all metals and non-metals:
Task 15.2. Name the compounds obtained. Determine what properties (oxidizing agent or reducing agent) sulfur exhibits in these reactions.
As a typical non-metal, the simple substance sulfur can be both an oxidizing agent and a reducing agent:
Sometimes these properties appear in one reaction:
Since the oxidizing atom and the reducing atom are the same, they can be “added,” i.e., both processes require three sulfur atom.
Task 15.3. Set up the remaining coefficients in this equation.
Sulfur can react with acids - strong oxidizing agents:
Thus, being an active non-metal, sulfur forms many compounds. Let's consider the properties of hydrogen sulfide, sulfur oxides and their derivatives.
Hydrogen sulfide
H 2 S is hydrogen sulfide, a highly poisonous gas with a nasty rotten egg smell. It would be more correct to say that when egg whites rot, they decompose, releasing hydrogen sulfide.
Task 15.4. Based on the oxidation state of the sulfur atom in hydrogen sulfide, predict what properties this atom will exhibit in redox reactions.
Since hydrogen sulfide is a reducing agent (sulfur atom has lowest oxidation state), it oxidizes easily. Air oxygen oxidizes hydrogen sulfide even at room temperature:
Hydrogen sulfide burns:
Hydrogen sulfide is slightly soluble in water, and its solution exhibits the properties very weak acid (hydrogen sulfide H2S). It forms salts sulfides:
Question. How can you get hydrogen sulfide if you have sulfide?
Hydrogen sulfide is produced in laboratories by acting on sulfides stronger (than H2S) acids, for example:
Sulfur dioxide and sulfurous acid
SO 2- sulfur dioxide with a pungent suffocating odor. Poisonous. Dissolves in water to form sulfurous acid:
This acid is of medium strength, but very unstable, existing only in solutions. Therefore, when acting on its salts - sulf it s- other acids can produce sulfur dioxide:
When the resulting solution is boiled, this acid decomposes completely.
Task 15.5. Determine the degree of oxidation of sulfur in sulfur dioxide, sulfurous acid, sodium sulfite.
Since the oxidation state +4 for sulfur is intermediate, all of the listed compounds can be both oxidizing and reducing agents:
For example:
Task 15.6. Arrange the coefficients in these schemes using the electronic balance method. Indicate what properties a sulfur atom with oxidation state +4 exhibits in each of the reactions.
The reducing properties of sulfur dioxide are used in practice. So, when restored, some lose color organic compounds Therefore, sulfur oxide IV and sulfites are used in bleaching. Sodium sulfite, dissolved in water, slows down the corrosion of pipes, as it easily absorbs oxygen from water, and it is oxygen that is the “culprit” of corrosion:
Oxidizing in the presence of a catalyst, sulfur dioxide turns into sulfuric anhydride SO 3:
Sulfuric anhydride and sulfuric acid
Sulfuric anhydride SO 3- colorless liquid that reacts violently with water:
Sulfuric acid H2SO4- a strong acid that concentrated form actively absorbs moisture from the air (this property is used when drying various gases) and from some complex substances:
The physical properties of sulfur directly depend on the allotropic modification. For example, the most famous modification of sulfur is rhombic, S₈. It is quite fragile crystalline substance yellow color.
Structure of the rhombic sulfur molecule S₈
In addition to the rhombic one, there are many other modifications, the number of which, according to different sources, reaches three dozen.
Chemical properties of the element
At normal temperature The chemical activity of sulfur is quite small. But when heated, sulfur often interacts with all simple substances, metals and non-metals.
S + O₂ → SO₂
Sulfur is the most important element in life and animals, it is widely used in industry, ranging from medicine to pyrotechnic devices.
Sulfuric acid
Sulfuric acid has the formula H₂SO₄ and is the strongest dibasic acid. Previously, this substance was called oil of vitriol because the concentrated acid has a thick, oily consistency.
Sulfuric acid mixes easily with water, but such solutions must be prepared with caution: concentrated acid you need to carefully pour it into the water, and in no case vice versa.
Sulfuric acid is a caustic substance and can dissolve some. Therefore, it is often used in ore mining. Acid leaves severe burns on the skin, so it is extremely important to follow safety precautions when working with it.
Obtaining "oil of vitriol"
The industry uses a contact method for producing SO₂ (sulfur dioxide) through the oxidation of sulfur dioxide, which is formed during the combustion of sulfur. Next, sulfur trioxide SO₃ is obtained from sulfur dioxide, which is then dissolved in the most concentrated sulfuric acid. The resulting solution is called oleum. To obtain “oil of vitriol,” oleum is diluted with water.
Chemical properties of sulfuric acid
When interacting with metals, as well as carbon and sulfur, concentrated sulfuric acid oxidizes them:
Сu + 2H₂SO₄ (conc.) → CuSO₄ + SO₂ + 2H₂O.
C(graphite) + 2H₂SO₄ (conc., horizontal) → CO₂ + 2SO₂ + 2H₂O
S + 2H₂SO₄ (conc.) → 3SO₂ + 2H₂O
Dilute acid is capable of reacting with all metals that are to the left of hydrogen in the voltage series:
Fe + H₂SO₄ (dil.) → FeSO₄ + H₂
Zn + H₂SO₄ (dil.) → ZnSO₄ + H₂
In reactions with bases, dilute H₂SO₄ forms sulfates and hydrosulfates:
H₂SO₄ + NaOH → NaHSO₄ + H₂O;
H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O.
This acid can also react with basic oxides, resulting in sulfates:
CaO + H₂SO₄ → CaSO₄↓ + H₂O.
Hydrogen sulfide (H₂S) is a colorless gas with a rotten egg odor. It is denser than hydrogen. Hydrogen sulfide is deadly poisonous to humans and animals. Even a small amount of it in the air causes dizziness and nausea, but the worst thing is that after inhaling it for a long time, this smell is no longer felt. However, for hydrogen sulfide poisoning, there is a simple antidote: you should wrap a piece of bleach in a handkerchief, then moisten it, and sniff the package for a while. Hydrogen sulfide is produced by reacting sulfur with hydrogen at a temperature of 350 °C:
H₂ + S → H₂S
This is a redox reaction: during it, the oxidation states of the elements participating in it change.
In laboratory conditions, hydrogen sulfide is produced by treating iron sulfide with sulfuric or hydrochloric acid:
FeS + 2HCl → FeCl₂ + H₂S
This is an exchange reaction: in it, the interacting substances exchange their ions. This process is usually performed using a Kipp apparatus.
Kipp apparatus
Properties of hydrogen sulfide
When hydrogen sulfide burns, sulfur oxide 4 and water vapor are formed:
2H₂S + 3О₂ → 2Н₂О + 2SO₂
H₂S burns with a bluish flame, and if you hold an inverted beaker over it, clear condensate (water) will appear on its walls.
However, with a slight decrease in temperature, this reaction proceeds somewhat differently: a yellowish coating of free sulfur will appear on the walls of the pre-cooled glass:
2H₂S + O₂ → 2H₂O + 2S
The industrial method for producing sulfur is based on this reaction.
When a pre-prepared gaseous mixture of hydrogen sulfide and oxygen is ignited, an explosion occurs.
The reaction of hydrogen sulfide and sulfur(IV) oxide also produces free sulfur:
2H₂S + SO₂ → 2H₂O + 3S
Hydrogen sulfide is soluble in water, and three volumes of this gas can dissolve in one volume of water, forming weak and unstable hydrosulfide acid (H₂S). This acid is also called hydrogen sulfide water. As you can see, the formulas of hydrogen sulfide gas and hydrogen sulfide acid are written the same way.
If a solution of lead salt is added to hydrosulfide acid, a black precipitate of lead sulfide will form:
H₂S + Pb(NO₃)₂ → PbS + 2HNO₃
This is a qualitative reaction for the detection of hydrogen sulfide. It also demonstrates the ability of hydrosulfide acid to enter into exchange reactions with salt solutions. Thus, any soluble lead salt is a reagent for hydrogen sulfide. Some other metal sulfides also have a characteristic color, for example: zinc sulfide ZnS - white, cadmium sulfide CdS - yellow, copper sulfide CuS - black, antimony sulfide Sb₂S₃ - red.
By the way, hydrogen sulfide is an unstable gas and, when heated, almost completely decomposes into hydrogen and free sulfur:
H₂S → H₂ + S
Hydrogen sulfide interacts intensively with aqueous solutions halogens:
H₂S + 4Cl₂ + 4H₂O→ H₂SO₄ + 8HCl
Hydrogen sulfide in nature and human activity
Hydrogen sulfide is part of volcanic gases, natural gas and gases associated with oil fields. There is a lot of it in natural mineral waters, for example, in the Black Sea it lies at a depth of 150 meters and below.
Hydrogen sulfide is used:
- in medicine (treatment with hydrogen sulfide baths and mineral waters);
- in industry (production of sulfur, sulfuric acid and sulfides);
- in analytical chemistry (for the precipitation of heavy metal sulfides, which are usually insoluble);
- in organic synthesis (to produce sulfur analogues of organic alcohols (mercaptans) and thiophene (sulfur-containing aromatic hydrocarbon). Another recently emerging area in science is hydrogen sulfide energy. The production of energy from hydrogen sulfide deposits from the bottom of the Black Sea is being seriously studied.
The nature of redox reactions of sulfur and hydrogen
The reaction of hydrogen sulfide formation is redox:
Н₂⁰ + S⁰→ H₂⁺S²⁻
The process of interaction of sulfur with hydrogen is easily explained by the structure of their atoms. Hydrogen occupies first place in the periodic table, therefore, its charge atomic nucleus is equal to (+1), and 1 electron is circling around the nucleus of the atom. Hydrogen easily gives up its electron to atoms of other elements, turning into a positively charged hydrogen ion - a proton:
Н⁰ -1е⁻= Н⁺
Sulfur is in position sixteen in the periodic table. This means that the charge of the nucleus of its atom is (+16), and the number of electrons in each atom is also 16e⁻. The location of sulfur in the third period suggests that its sixteen electrons swirl around the atomic nucleus, forming 3 layers, the last of which contains 6 valence electrons. The number of valence electrons of sulfur corresponds to the number of group VI in which it is located in the periodic table.
So, sulfur can donate all six valence electrons, as in the case of the formation of sulfur(VI) oxide:
2S⁰ + 3O2⁰ → 2S⁺⁶O₃⁻²
In addition, as a result of the oxidation of sulfur, 4e⁻ can be given up by its atom to another element to form sulfur(IV) oxide:
S⁰ + O2⁰ → S⁺4 O2⁻²
Sulfur can also donate two electrons to form sulfur(II) chloride:
S⁰ + Cl2⁰ → S⁺² Cl2⁻
In all three of the above reactions, sulfur donates electrons. Consequently, it is oxidized, but at the same time acts as a reducing agent for oxygen atoms O and chlorine Cl. However, in the case of the formation of H2S, oxidation is the lot of hydrogen atoms, since they are the ones who lose electrons, restoring the external energy level of sulfur from six electrons to eight. As a result, each hydrogen atom in its molecule becomes a proton:
Н2⁰-2е⁻ → 2Н⁺,
and the sulfur molecule, on the contrary, being reduced, turns into a negatively charged anion (S⁻²): S⁰ + 2е⁻ → S⁻²
Thus, in chemical reaction In the formation of hydrogen sulfide, the oxidizing agent is sulfur.
From the point of view of the manifestation of sulfur in various oxidation states, another interesting interaction between sulfur(IV) oxide and hydrogen sulfide is the reaction to produce free sulfur:
2H₂⁺S-²+ S⁺⁴О₂-²→ 2H₂⁺O-²+ 3S⁰
As can be seen from the reaction equation, both the oxidizing agent and the reducing agent in it are sulfur ions. Two sulfur anions (2-) donate two of their electrons to the sulfur atom in the sulfur(II) oxide molecule, as a result of which all three sulfur atoms are reduced to free sulfur.
2S-² - 4е⁻→ 2S⁰ - reducing agent, oxidizes;
S⁺⁴ + 4е⁻→ S⁰ - oxidizing agent, reduced.
