Sulfur (IV) oxide and sulfurous acid. SO2 - sulfur oxide (IV), sulfur dioxide, sulfur dioxide, sulfur dioxide
Sulfur oxide ( sulphur dioxide, sulfur dioxide, sulfur dioxide) is a colorless gas that under normal conditions has a sharp characteristic odor (similar to the smell of a burning match). Liquefied under pressure room temperature. Sulfur dioxide is soluble in water, resulting in the formation of unstable sulfuric acid. This substance is also soluble in sulfuric acid and ethanol. This is one of the main components that make up volcanic gases.
Sulphur dioxide
The production of SO2 - sulfur dioxide - industrially involves burning sulfur or roasting sulfides (pyrite is mainly used).
4FeS2 (pyrite) + 11O2 = 2Fe2O3 + 8SO2 (sulfur dioxide).
In a laboratory setting, sulfur dioxide can be produced by treating hydrosulfites and sulfites with strong acids. In this case, the resulting sulfurous acid immediately breaks down into water and sulfur dioxide. For example:
Na2SO3 + H2SO4 (sulfuric acid) = Na2SO4 + H2SO3 (sulfurous acid).
H2SO3 (sulfurous acid) = H2O (water) + SO2 (sulfur dioxide).
The third method of producing sulfur dioxide involves the action of concentrated sulfuric acid on low-active metals when heated. For example: Cu (copper) + 2H2SO4 (sulfuric acid) = CuSO4 (copper sulfate) + SO2 (sulfur dioxide) + 2H2O (water).
Chemical properties sulfur dioxide
The formula of sulfur dioxide is SO3. This substance belongs to acid oxides.
1. Sulfur dioxide dissolves in water, resulting in sulfurous acid. Under normal conditions, this reaction is reversible.
SO2 (sulfur dioxide) + H2O (water) = H2SO3 (sulfurous acid).
2. With alkalis, sulfur dioxide forms sulfites. For example: 2NaOH (sodium hydroxide) + SO2 (sulfur dioxide) = Na2SO3 (sodium sulfite) + H2O (water).
3. The chemical activity of sulfur dioxide is quite high. The reducing properties of sulfur dioxide are most pronounced. In such reactions, the oxidation state of sulfur increases. For example: 1) SO2 (sulfur dioxide) + Br2 (bromine) + 2H2O (water) = H2SO4 (sulfuric acid) + 2HBr (hydrogen bromide); 2) 2SO2 (sulfur dioxide) + O2 (oxygen) = 2SO3 (sulfite); 3) 5SO2 (sulfur dioxide) + 2KMnO4 (potassium permanganate) + 2H2O (water) = 2H2SO4 (sulfuric acid) + 2MnSO4 (manganese sulfate) + K2SO4 (potassium sulfate).
The last reaction is an example of a qualitative reaction to SO2 and SO3. The solution becomes purple in color.)
4. In the presence of strong reducing agents sulfur dioxide may exhibit oxidizing properties. For example, in order to extract sulfur from exhaust gases in the metallurgical industry, they use the reduction of sulfur dioxide with carbon monoxide (CO): SO2 (sulfur dioxide) + 2CO (carbon monoxide) = 2CO2 + S (sulfur).
Also, the oxidizing properties of this substance are used to obtain phosphorous acid: PH3 (phosphine) + SO2 (sulfur dioxide) = H3PO2 (phosphorous acid) + S (sulfur).
Where is sulfur dioxide used?
Sulfur dioxide is mainly used to produce sulfuric acid. It is also used in the production of low-alcohol drinks (wine and other mid-price drinks). Due to the property of this gas to kill various microorganisms, it is used to fumigate warehouses and vegetable storage. In addition, sulfur oxide is used to bleach wool, silk, and straw (those materials that cannot be bleached with chlorine). In laboratories, sulfur dioxide is used as a solvent and in order to obtain various salts of sulfur dioxide.
Physiological effects
Sulfur dioxide has strong toxic properties. Symptoms of poisoning are cough, runny nose, hoarseness, a peculiar taste in the mouth, and severe sore throat. When sulfur dioxide is inhaled in high concentrations, difficulty swallowing and choking, speech disturbance, nausea and vomiting occur, and acute pulmonary edema may develop.
MPC of sulfur dioxide:
- indoors - 10 mg/m³;
- average daily maximum one-time exposure in atmospheric air - 0.05 mg/m³.
Sensitivity to sulfur dioxide varies among individuals, plants, and animals. For example, among trees the most resistant are oak and birch, and the least resistant are spruce and pine.
Sulfur dioxide has a molecular structure similar to ozone. The sulfur atom at the center of the molecule is bonded to two oxygen atoms. This gaseous product of sulfur oxidation is colorless, emits a pungent odor, and when conditions change, it easily condenses into clear liquid. The substance is highly soluble in water and has antiseptic properties. IN large quantities SO 2 is obtained in the chemical industry, namely in the sulfuric acid production cycle. The gas is widely used for processing agricultural and food products, bleaching fabrics in the textile industry.
Systematic and trivial names of substances
It is necessary to understand the variety of terms related to the same compound. Official name of the connection, chemical composition which is reflected by the formula SO 2, is sulfur dioxide. IUPAC recommends the use of this term and its English equivalent- Sulfur dioxide. Textbooks for schools and universities often mention another name - sulfur (IV) oxide. The Roman numeral in parentheses indicates the valency of the S atom. Oxygen in this oxide is divalent, and the oxidation number of sulfur is +4. In the technical literature, outdated terms such as sulfur dioxide, sulfuric acid anhydride (a product of its dehydration) are used.
Composition and features of the molecular structure of SO 2
The SO 2 molecule is formed by one sulfur atom and two oxygen atoms. There is an angle of 120° between covalent bonds. In the sulfur atom, sp2 hybridization occurs—the clouds of one s and two p electrons are aligned in shape and energy. They are the ones who participate in the formation of a covalent bond between sulfur and oxygen. In the O–S pair, the distance between the atoms is 0.143 nm. Oxygen is a more electronegative element than sulfur, which means that the bonding pairs of electrons shift from the center to external corners. The entire molecule is also polarized, the negative pole is the O atoms, the positive pole is the S atom.
Some physical parameters of sulfur dioxide
Tetravalent sulfur oxide at normal levels environment retains a gaseous state of aggregation. The formula of sulfur dioxide allows you to determine its relative molecular and molar mass: Mr(SO 2) = 64.066, M = 64.066 g/mol (can be rounded to 64 g/mol). This gas is almost 2.3 times heavier than air (M(air) = 29 g/mol). Dioxide has a sharp, specific smell of burning sulfur, which is difficult to confuse with any other. It is unpleasant, irritates the mucous membranes of the eyes, and causes a cough. But sulfur (IV) oxide is not as poisonous as hydrogen sulfide.
Under pressure at room temperature, sulfur dioxide gas liquefies. At low temperatures the substance is in a solid state, melts at -72...-75.5 °C. With a further increase in temperature, liquid appears, and at -10.1 °C gas is formed again. SO 2 molecules are thermally stable; decomposition into atomic sulfur and molecular oxygen occurs at very high temperatures (about 2800 ºC).
Solubility and interaction with water
Sulfur dioxide, when dissolved in water, partially reacts with it to form a very weak sulfurous acid. At the moment of receipt, it immediately decomposes into anhydride and water: SO 2 + H 2 O ↔ H 2 SO 3. In fact, it is not sulfurous acid that is present in the solution, but hydrated SO 2 molecules. Dioxide gas reacts better with cool water, and its solubility decreases with increasing temperature. Under normal conditions, up to 40 volumes of gas can dissolve in 1 volume of water.
Sulfur dioxide in nature
Significant amounts of sulfur dioxide are released with volcanic gases and lava during eruptions. Many types of anthropogenic activities also lead to increased concentrations of SO 2 in the atmosphere.
Sulfur dioxide is released into the air by metallurgical plants, where waste gases are not captured during ore roasting. Many fossil fuels contain sulfur, resulting in significant amounts of sulfur dioxide being released into atmospheric air when burning coal, oil, gas, and fuel obtained from them. Sulfur dioxide becomes toxic to humans at concentrations in the air above 0.03%. A person begins to experience shortness of breath, and symptoms resembling bronchitis and pneumonia may occur. Very high concentrations of sulfur dioxide in the atmosphere can lead to severe poisoning or death.
Sulfur dioxide - production in the laboratory and in industry
Laboratory methods:
- When sulfur is burned in a flask with oxygen or air, dioxide is obtained according to the formula: S + O 2 = SO 2.
- You can act on the salts of sulfurous acid with stronger inorganic acids, it is better to take hydrochloric acid, but you can use diluted sulfuric acid:
- Na 2 SO 3 + 2HCl = 2NaCl + H 2 SO 3;
- Na 2 SO 3 + H 2 SO 4 (diluted) = Na 2 SO 4 + H 2 SO 3;
- H 2 SO 3 = H 2 O + SO 2.
3. When copper reacts with concentrated sulfuric acid, it is not hydrogen that is released, but sulfur dioxide:
2H 2 SO 4 (conc.) + Cu = CuSO 4 + 2H 2 O + SO 2.
Modern methods industrial production sulfur dioxide:
- Oxidation of natural sulfur when it is burned in special furnaces: S + O 2 = SO 2.
- Firing iron pyrite (pyrite).
Basic chemical properties of sulfur dioxide
Sulfur dioxide is a chemically active compound. In redox processes, this substance often acts as a reducing agent. For example, when molecular bromine reacts with sulfur dioxide, the reaction products are sulfuric acid and hydrogen bromide. The oxidizing properties of SO 2 appear if this gas is passed through hydrogen sulfide water. As a result, sulfur is released, self-oxidation-self-reduction occurs: SO 2 + 2H 2 S = 3S + 2H 2 O.
Sulfur dioxide exhibits acidic properties. It corresponds to one of the weakest and most unstable acids - sulfurous. This connection is in pure form does not exist, the acidic properties of a sulfur dioxide solution can be detected using indicators (litmus turns pink). Sulfurous acid produces medium salts - sulfites and acidic salts - hydrosulfites. Among them there are stable compounds.
The process of oxidation of sulfur in dioxide to the hexavalent state in sulfuric acid anhydride is catalytic. The resulting substance dissolves energetically in water and reacts with H 2 O molecules. The reaction is exothermic, sulfuric acid is formed, or rather its hydrated form.
Practical uses of sulfur dioxide
The main method of industrial production of sulfuric acid, which requires elemental dioxide, has four stages:
- Obtaining sulfur dioxide by burning sulfur in special furnaces.
- Purification of the resulting sulfur dioxide from all kinds of impurities.
- Further oxidation to hexavalent sulfur in the presence of a catalyst.
- Absorption of sulfur trioxide by water.
Previously, almost all the sulfur dioxide needed to produce sulfuric acid was industrial scale, obtained by roasting pyrite as a by-product of steelmaking. New types of processing of metallurgical raw materials use less ore combustion. Therefore, the main starting material for sulfuric acid production in last years became natural sulfur. Significant global reserves of this raw material and its availability make it possible to organize large-scale processing.
Sulfur dioxide finds wide application not only in the chemical industry, but also in other sectors of the economy. Textile mills use this substance and the products of its chemical reaction to bleach silk and wool fabrics. This is a type of chlorine-free bleaching that does not destroy the fibers.
Sulfur dioxide has excellent disinfectant properties, which is used in the fight against fungi and bacteria. Sulfur dioxide is used to fumigate agricultural storage facilities, wine barrels and cellars. SO 2 is used in Food Industry as a preservative and antibacterial substance. They add it to syrups and soak fresh fruits in it. Sulfitization
Sugar beet juice decolorizes and disinfects raw materials. Canned vegetable purees and juices also contain sulfur dioxide as an antioxidant and preservative.
Hydrogen sulfide – H2S
Sulfur compounds -2, +4, +6. Qualitative reactions to sulfides, sulfites, sulfates.
Receipt upon interaction:
1. hydrogen with sulfur at t – 300 0
2. when acting on mineral acid sulfides:
Na 2 S+2HCl =2 NaCl+H 2 S
a colorless gas with the smell of rotten eggs, poisonous, heavier than air, and dissolving in water to form weak hydrogen sulfide acid.
Chemical properties
Acid-base properties
1. A solution of hydrogen sulfide in water - hydrosulfide acid - is a weak dibasic acid, therefore it dissociates stepwise:
H 2 S ↔ HS - + H +
HS - ↔ H - + S 2-
2. Hydrogen sulfide acid has general properties acids, reacts with metals, basic oxides, bases, salts:
H 2 S + Ca = CaS + H 2
H 2 S + CaO = CaS + H 2 O
H 2 S + 2NaOH = Na 2 S + 2H 2 O
H 2 S + CuSO 4 = CuS↓ + H 2 SO 4
All acid salts - hydrosulfides - are highly soluble in water. Normal salts - sulfides - dissolve in water in different ways: sulfides of alkali and alkaline earth metals are highly soluble, sulfides of other metals are insoluble in water, and sulfides of copper, lead, mercury and some other heavy metals are not soluble even in acids (except nitric acid)
CuS+4HNO 3 =Cu(NO 3) 2 +3S+2NO+2H 2 O
Soluble sulfides undergo hydrolysis - at the anion.
Na 2 S ↔ 2Na + + S 2-
S 2- +HOH ↔HS - +OH -
Na 2 S + H 2 O ↔ NaHS + NaOH
A qualitative reaction to hydrosulfide acid and its soluble salts (i.e., to the sulfide ion S 2-) is their interaction with soluble lead salts, which results in the formation of a black PbS precipitate
Na 2 S + Pb(NO 3) 2 = 2NaNO 3 + PbS↓
Pb 2+ + S 2- = PbS↓
Shows only restorative properties, because the sulfur atom has the lowest oxidation state -2
1. with oxygen
a) with a disadvantage
2H 2 S -2 +O 2 0 = S 0 +2H 2 O -2
b) with excess oxygen
2H 2 S+3O 2 =2SO 2 +2H 2 O
2. with halogens (bromine water discoloration)
H 2 S -2 +Br 2 =S 0 +2HBr -1
3. with conc. HNO3
H 2 S+2HNO 3 (k) = S+2NO 2 +2H 2 O
b) with strong oxidizing agents (KMnO 4, K 2 CrO 4 in an acidic environment)
2KMnO 4 +3H 2 SO 4 +5H 2 S = 5S+2MnSO 4 +K 2 SO 4 +8H 2 O
c) hydrosulfide acid is oxidized not only by strong oxidizing agents, but also by weaker ones, for example, iron (III) salts, sulfurous acid, etc.
2FeCl 3 + H 2 S = 2FeCl 2 + S + 2HCl
H 2 SO 3 + 2H 2 S = 3S + 3H 2 O
Receipt
1. combustion of sulfur in oxygen.
2. combustion of hydrogen sulfide in excess O 2
2H 2 S+3O 2 = 2SO 2 +2H 2 O
3. sulfide oxidation
2CuS+3O2 = 2SO2 +2CuO
4. interaction of sulfites with acids
Na 2 SO 3 +H 2 SO 4 =Na 2 SO 4 +SO 2 +H 2 O
5. interaction of metals in the activity series after (H 2) with conc. H2SO4
Cu+2H 2 SO 4 = CuSO 4 + SO 2 +2H 2 O
Physical properties
Gas, colorless, with a suffocating odor of burnt sulfur, poisonous, more than 2 times heavier than air, highly soluble in water (at room temperature, about 40 volumes of gas dissolve in one volume).
Chemical properties:
Acid-base properties
SO 2 is a typical acidic oxide.
1.with alkalis, forming two types of salts: sulfites and hydrosulfites
2KOH+SO2 = K2SO3 +H2O
KOH+SO 2 = KHSO 3 +H 2 O
2.with basic oxides
K 2 O+SO 2 = K 2 SO 3
3. weak sulfurous acid is formed with water
H 2 O+SO 2 = H 2 SO 3
Sulfurous acid exists only in solution and is a weak acid.
has all the general properties of acids.
4. qualitative reaction to sulfite - ion - SO 3 2 - action of mineral acids
Na 2 SO 3 +2HCl= 2Na 2 Cl+SO 2 +H 2 O smell of burnt sulfur
Redox properties
In ORR it can be both an oxidizing agent and a reducing agent, because the sulfur atom in SO 2 has an intermediate oxidation state of +4.
As an oxidizing agent:
SO 2 + 2H 2 S = 3S + 2H 2 S
As a reducing agent:
2SO 2 +O 2 = 2SO 3
Cl 2 +SO 2 +2H 2 O = H 2 SO 4 +2HCl
2KMnO 4 +5SO 2 +2H 2 O = K 2 SO 4 +2H 2 SO 4 +2MnSO 4
Sulfur oxide (VI) SO 3 (sulfuric anhydride)
Receipt:
Oxidation of sulfur dioxide
2SO 2 + O 2 = 2SO 3 ( t 0 , kat)
Physical properties
A colorless liquid, at temperatures below 17 0 C it turns into a white crystalline mass. Thermally unstable compound, completely decomposes at 700 0 C. It is highly soluble in water and anhydrous sulfuric acid and reacts with it to form oleum
SO 3 + H 2 SO 4 = H 2 S 2 O 7
Chemical properties
Acid-base properties
Typical acid oxide.
1.with alkalis, forming two types of salts: sulfates and hydrosulfates
2KOH+SO 3 = K 2 SO 4 +H 2 O
KOH+SO 3 = KHSO 4 +H 2 O
2.with basic oxides
CaO+SO 2 = CaSO 4
3. with water
H 2 O + SO 3 = H 2 SO 4
Redox properties
Sulfur oxide (VI) is a strong oxidizing agent, usually reduced to SO 2
3SO 3 + H 2 S = 4SO 2 + H 2 O
Sulfuric acid H 2 SO 4
Preparation of sulfuric acid
In industry, acid is produced by contact method:
1. pyrite firing
4FeS 2 +11O 2 = 2Fe 2 O 3 + 8SO 2
2. oxidation of SO 2 to SO 3
2SO 2 + O 2 = 2SO 3 ( t 0 , kat)
3. dissolution of SO 3 in sulfuric acid
n SO 3 + H 2 SO 4 = H 2 SO 4 ∙ n SO 3 (oleum)
H2SO4∙ n SO 3 + H 2 O = H 2 SO 4
Physical properties
H 2 SO 4 is a heavy oily liquid, odorless and colorless, hygroscopic. It mixes with water in any ratio; when concentrated sulfuric acid is dissolved in water, a large amount of heat is released, so it must be carefully poured into water, and not vice versa (first water, then acid, otherwise big trouble will happen)
A solution of sulfuric acid in water with a H 2 SO 4 content of less than 70% is usually called dilute sulfuric acid, more than 70% - concentrated.
Chemical properties
Acid-base
Dilute sulfuric acid reveals everything characteristic properties strong acids. Dissociates in aqueous solution:
H 2 SO 4 ↔ 2H + + SO 4 2-
1. with basic oxides
MgO + H 2 SO 4 = MgSO 4 + H 2 O
2. with grounds
2NaOH +H 2 SO 4 = Na 2 SO 4 + 2H 2 O
3. with salts
BaCl 2 + H 2 SO 4 = BaSO 4 ↓ + 2HCl
Ba 2+ + SO 4 2- = BaSO 4 ↓ (white precipitate)
Qualitative reaction to sulfate ion SO 4 2-
Thanks to more high temperature boiling, compared to other acids, sulfuric acid, when heated, displaces them from salts:
NaCl + H 2 SO 4 = HCl + NaHSO 4
Redox properties
In dilute H 2 SO 4 the oxidizing agents are H + ions, and in concentrated H 2 SO 4 the oxidizing agents are SO 4 2 sulfate ions.
Metals in the activity series up to hydrogen dissolve in dilute sulfuric acid, sulfates are formed and hydrogen is released
Zn + H 2 SO 4 = ZnSO 4 + H 2
Concentrated sulfuric acid is a vigorous oxidizing agent, especially when heated. It oxidizes many metals, non-metals, inorganic and organic substances.
H 2 SO 4 (k) oxidizing agent S +6
With more active metals, sulfuric acid can be reduced to a variety of products depending on concentration
Zn + 2H 2 SO 4 = ZnSO 4 + SO 2 + 2H 2 O
3Zn + 4H 2 SO 4 = 3ZnSO 4 + S + 4H 2 O
4Zn + 5H 2 SO 4 = 4ZnSO 4 + H 2 S + 4H 2 O
Concentrated sulfuric acid oxidizes some non-metals (sulfur, carbon, phosphorus, etc.), reducing to sulfur oxide (IV)
S + 2H 2 SO 4 = 3SO 2 + 2H 2 O
C + 2H 2 SO 4 = 2SO 2 + CO 2 + 2H 2 O
Interaction with some complex substances
H 2 SO 4 + 8HI = 4I 2 + H 2 S + 4 H 2 O
H 2 SO 4 + 2HBr = Br 2 + SO 2 + 2H 2 O
Sulfuric acid salts
2 types of salts: sulfates and hydrosulfates
Salts of sulfuric acid have all the general properties of salts. Their relationship to heat is special. Sulfates of active metals (Na, K, Ba) do not decompose even when heated above 1000 0 C, salts of less active metals (Al, Fe, Cu) decompose even with slight heating
Sulfur dioxide is a colorless gas with a pungent odor. The molecule has angular shape.
- Melting point - -75.46 °C,
- Boiling point - -10.6 °C,
- Gas density - 2.92655 g/l.
Easily liquefies into a colorless, highly mobile liquid at a temperature of 25 ° C and a pressure of about 0.5 MPa.
For the liquid form, the density is 1.4619 g/cm 3 (at - 10 ° C).
Solid sulfur dioxide - colorless crystals, orthorhombic system.
Sulfur dioxide dissociates noticeably only around 2800 °C.
Dissociation of liquid sulfur dioxide proceeds according to the following scheme:
2SO 2 ↔ SO 2+ + SO 3 2-
Three-dimensional model of a molecule
The solubility of sulfur dioxide in water depends on temperature:
- at 0 °C 22.8 g of sulfur dioxide dissolves in 100 g of water,
- at 20 °C - 11.5 g,
- at 90 °C - 2.1 g.
An aqueous solution of sulfur dioxide is sulfurous acid H 2 SO 3.
Sulfur dioxide is soluble in ethanol, H 2 SO 4, oleum, CH 3 COOH. Liquid sulfur dioxide is mixed in any ratio with SO 3. CHCl 3, CS 2, diethyl ether.
Liquid sulfur dioxide dissolves chlorides. Metal iodides and thiocyanates do not dissolve.
Salts dissolved in liquid sulfur dioxide dissociate.
Sulfur dioxide is capable of being reduced to sulfur and oxidized to hexavalent sulfur compounds.
Sulfur dioxide is toxic. At a concentration of 0.03-0.05 mg/l, it irritates the mucous membranes, respiratory organs, and eyes.
Basic industrial method obtaining sulfur dioxide - from sulfur pyrite FeS 2 by burning it and further processing with weak cold H 2 SO 4.
In addition, sulfur dioxide can be produced by burning sulfur, and also as a by-product of roasting copper and zinc sulfide ores.
Sulfide sulfur is available to plants only after converting to the sulfate form. Most of the sulfur is present in the soil as organic compounds, not absorbed by plants. Only after mineralization organic matter and the transition of sulfur into sulfate form, organic sulfur becomes available to plants.
The chemical industry does not produce fertilizers with the main active ingredient sulfur dioxide. However, it is found as an impurity in many fertilizers. These include phosphogypsum, simple superphosphate, ammonium sulfate, potassium sulfate, potassium magnesia, gypsum, oil shale ash, manure, peat and many others.
Absorption of sulfur dioxide by plants
Sulfur enters plants through the roots in the form SO 4 2- and leaves in the form of sulfur dioxide. At the same time, the absorption of sulfur from the atmosphere provides up to 80% of the plants’ needs for this element. In this regard, near industrial centers, where the atmosphere is rich in sulfur dioxide, plants are well supplied with sulfur. In remote areas, the amount of sulfur dioxide in precipitation and the atmosphere is greatly reduced and the nutrition of plants with sulfur depends on its presence in the soil.
Sulfur(IV) oxide has acidic properties, which manifest themselves in reactions with substances that exhibit basic properties. Acidic properties appear when interacting with water. This produces a solution of sulfurous acid:
The oxidation degree of sulfur in sulfur dioxide gas (+4) determines the reducing and oxidizing properties of sulfur dioxide gas:
vo-tel: S+4 – 2e => S+6
ok-tel: S+4 + 4e => S0
Reductive properties are manifested in reactions with strong oxidizing agents: oxygen, halogens, nitric acid, potassium permanganate and others. For example:
2SO2 + O2 = 2SO3
S+4 – 2e => S+6 2
O20 + 4e => 2O-2 1
With strong reducing agents, the gas exhibits oxidizing properties. For example, if you mix sulfur dioxide and hydrogen sulfide, they interact under normal conditions:
2H2S + SO2 = 3S + 2H2O
S-2 – 2e => S0 2
S+4 + 4e => S0 1
Sulfurous acid exists only in solution. It is unstable and decomposes into sulfur dioxide and water. Sulfurous acid is not a strong acid. It is an acid of medium strength and dissociates stepwise. When alkali is added to sulfurous acid, salts are formed. Sulfurous acid produces two series of salts: medium - sulfites and acidic - hydrosulfites.
Sulfur(VI) oxide
Sulfur trioxide exhibits acidic properties. It reacts violently with water, releasing a large amount of heat. This reaction is used to produce the most important product of the chemical industry - sulfuric acid.
SO3 + H2O = H2SO4
Since sulfur in sulfur trioxide has the highest oxidation state, sulfur(VI) oxide exhibits oxidizing properties. For example, it oxidizes halides, nonmetals with low electronegativity:
2SO3 + C = 2SO2 + CO2
S+6 + 2e => S+4 2
C0 – 4e => C+4 2
Sulfuric acid reacts three types: acid-base, ion exchange, redox. It also actively interacts with organic substances.
Acid-base reactions
Sulfuric acid exhibits acidic properties in reactions with bases and basic oxides. These reactions are best carried out with dilute sulfuric acid. Since sulfuric acid is dibasic, it can form both intermediate salts (sulfates) and acidic ones (hydrogen sulfates).
Ion exchange reactions
Sulfuric acid is characterized by ion exchange reactions. At the same time, it interacts with salt solutions, forming a precipitate, a weak acid, or releasing gas. These reactions occur at a faster rate if you take 45% or even more dilute sulfuric acid. Gas evolution occurs in reactions with salts of unstable acids, which decompose to form gases (carbonic, sulfur dioxide, hydrogen sulfide) or to form volatile acids such as hydrochloric acid.
Redox reactions
Sulfuric acid manifests its properties most clearly in redox reactions, since in its composition sulfur has the highest oxidation state of +6. The oxidizing properties of sulfuric acid can be detected in a reaction, for example, with copper.
There are two oxidizing elements in a sulfuric acid molecule: a sulfur atom with CO. +6 and hydrogen ions H+. Copper cannot be oxidized by hydrogen to the +1 oxidation state, but sulfur can. This is the reason for the oxidation of such an inactive metal as copper by sulfuric acid.
