Amphoteric hydroxides and oxides. Amphoteric compounds
Video tutorial 2: Amphoteric hydroxides. Experiments
Lecture: Characteristic Chemical properties bases and amphoteric hydroxides
Hydroxides and their classification
As you already know, bases are formed by metal atoms and a hydroxo group (OH -), therefore they are otherwise called hydroxides. There are several classifications of bases.
1. In relation to water they are divided into:
soluble,
insoluble.
Soluble bases include hydroxides of alkali and alkaline earth metals, which is why they are called alkalis. Ammonium hydroxide can also be included in this group, but unlike the former, it is a weaker electrolyte. Bases formed by other metals do not dissolve in water. Alkalis in an aqueous solution dissociate completely to metal cations and hydroxide anions - OH - ions. For example: NaOH → Na + + OH - .
2. Based on their interaction with other chemicals, hydroxides are divided into:
basic hydroxides,
acid hydroxides (oxygen-containing acids),
amphoteric hydroxides.
This division depends on the charge of the metal cation. When the charge of the cation is +1 or +2, the hydroxide will have basic properties. Amphoteric bases are considered to be hydroxides whose metal cations have a charge of +3 and +4.
But there are a number of exceptions:
La(OH) 3 , Bi(OH) 3 , Tl(OH) 3 – bases;
Be (OH) 2 , Sn (OH) 2 , Pb (OH) 2 , Zn (OH) 2 , Ge (OH) 2 are amphoteric bases.
Chemical properties of bases
Bases are capable of reacting with acids and acid oxides. During the interaction, salts and water are formed:
Ba(OH) 2 + CO 2 → BaCO 3 + H 2 O;
KOH + HCl → KCl + H 2 O.
Alkalis and ammonium hydroxide always react with salt solutions, only in the case of the formation of insoluble bases:
2KOH + FeCl 2 → 2KCl + Fe(OH) 2;
6NH 4 OH + Al 2 (SO 4) 3 → 2Al(OH) 3 + 3(NH 4) 2SO 4 .
The reaction of an acid with a base is called neutralization. During this reaction, acid cations H+ and base anions OH- form water molecules. After which, the solution medium becomes neutral. As a result, heat begins to be released. In solutions, this leads to gradual heating of the liquid. In the case of strong solutions, the heat is more than enough to cause the liquid to begin to boil. It must be remembered that the neutralization reaction occurs quite quickly.
Chemical properties of amphoteric hydroxides
Amphoteric bases react with both acids and alkalis. During the interaction, salt and water are formed. When undergoing any reaction with acids, amphoteric bases always exhibit the properties of typical bases:
Cr(OH) 3 + 3HCl → CrCl 3 + 3H 2 O.
During the reaction with alkalis, amphoteric bases are able to exhibit the properties of acids. In the process of fusion with alkalis, salt and water are formed.
There are hydroxides that react with both acids and bases, depending on the conditions. These compounds exhibiting a dual nature are called amphoteric hydroxides. They are formed by a metal cation and a hydroxide ion, like all bases. Only those hydroxides that contain the following metals have the ability to act as acids and bases: Be, Zn, Al, Pb, Sn, Ga, Cd, Fe, Cr(III), etc. As can be seen from Periodic table DI. Mendeleev, hydroxides with a dual nature form metals, which are closest to non-metals. It is believed that such elements are transitional forms, and the division into metals and non-metals is quite arbitrary.
Amphoteric hydroxides are solid, powdery, finely crystalline substances that are most often white in color, insoluble in water, and weakly conduct current (weak electrolytes). However, some of these bases can dissolve in acids and alkalis. Dissociation of “dual compounds” in aqueous solutions occurs according to the type of acids and bases. This is due to the fact that the holding force between metal and oxygen atoms (Me—O) and between oxygen and hydrogen atoms (O—H) is practically equal, i.e. Me - O - H. Therefore, these bonds will be broken simultaneously, and these substances will dissociate into H+ cations and OH- anions.
Amphoteric hydroxide - Be(OH) 2 - will help confirm the dual nature of these compounds. Let's consider the interaction of beryllium hydroxide with an acid and a base.
1. Be(OH) 2 + 2HCl -BeCl 2 +2H 2 O.
2. Be(OH) 2 + 2KOH - K 2 - potassium tetrahydroxoberyllate.
In the first case, a neutralization reaction takes place, the result of which is the formation of salt and water. In the second case, the reaction product will be. The neutralization reaction is typical for all hydroxides without exception, but interaction with their own kind is typical only for amphoteric ones. Such dual properties will also be exhibited by other amphoteric compounds - oxides and the metals themselves from which they are formed.
Other chemical properties of such hydroxides will be characteristic of all bases:
1. Thermal decomposition, reaction products - the corresponding oxide and water: Be(OH) 2 -BeO+H 2 O.
You also need to remember that there are substances with which amphoteric hydroxides do not interact, i.e. doesn't work, this:
- non-metals;
- metals;
- insoluble bases;
- amphoteric hydroxides.
- medium salts.
These compounds are obtained by precipitation of the corresponding salt solutions with alkali:
BeCl 2 + 2KOH - Be(OH) 2 + 2KCl.
Salts of some elements during this reaction form a hydrate, the properties of which almost completely correspond to those of hydroxides with a dual nature. The bases themselves with dual properties are included in the composition of minerals, in the form of which they are found in nature (bauxite, goethite, etc.).
Thus, amphoteric hydroxides are those that, depending on the nature of the substance that reacts with them, can act as bases or acids. Most often they correspond to amphoteric oxides containing the corresponding metal (ZnO-Zn(OH) 2; BeO - Be(OH) 2), etc.).
There are three main classes of inorganic chemical compounds: oxides, hydroxides and salts. The first are divided into two groups: non-salt-forming (these include carbon monoxide, nitrous oxide, nitrogen monoxide, etc.) and salt-forming, which, in turn, are basic, acidic and amphoteric. Hydroxides are divided into acids, bases and amphoteric. There are basic, acidic, medium and double salts. Amphoteric oxides and hydroxides will be described in more detail below.
What is amphotericity?
This is the ability of an inorganic chemical to exhibit both acidic and basic properties, depending on the reaction conditions. Substances that have this kind of feature may include oxides and hydroxides. Among the first are the oxide and dioxide of tin, beryllium, manganese, zinc, iron (II), (III). Amphoteric hydroxides are represented by the following substances: beryllium, aluminum, iron (II) hydroxide, iron and aluminum metahydroxide, titanium dihydroxide-oxide. The most common and frequently used of the compounds listed above are iron and aluminum oxide, as well as hydroxides of these metals.
Chemical properties of amphoteric oxides
Amphoteric oxides have both the properties of acidic and basic compounds. As acidic, they can interact with alkalis. In this type of reaction, salt and water are formed. They also react chemically with basic oxides. Displaying their basic properties, they interact with acids, resulting in the formation of salt and water, as well as with acidic oxides, due to which salt can be obtained.
Examples of reaction equations involving amphoteric oxides
AI 2 O 3 + 2KOH = 2KAIO 2 + H 2 O - this reaction shows the acidic properties of amphoteric oxides. 2АІ 2 О 3 + 6НСІ = 4АІСІ 3 + 3Н 2 О; АІ 2 О 3 + 3СО 2 = АІ2(СО 3) 3 - these equations serve as an example of the basic chemical properties of such oxides.
Chemical properties of amphoteric hydroxides
They are capable of reacting chemically with both strong acids and alkalis, and some of them also react with weak acids. All of them, when exposed to high temperatures, decompose into oxide and water. When an amphoteric hydroxide reacts with an acid, salt and water are formed. All such hydroxides are insoluble in water, and therefore can only react with solutions of certain compounds, but not with dry substances.
Physical properties of amphoteric oxides, methods of their preparation and application
Ferum(II) oxide is perhaps the most common amphoteric oxide. There are quite a few ways to obtain it. It is widely used in industry. Other amphoteric oxides are also used in many industries: from metallurgy to the food industry.
Appearance, preparation and use of ferum (II) oxide
It is a black solid. Its crystal lattice is similar to that of table salt. It can be found in nature as the mineral wustite.
Given chemical compound obtained in four different ways. First— reduction of iron (III) oxide using carbon monoxide. In this case, by mixing the same amount of these two substances, you can get two parts of iron (II) oxide and one part of carbon dioxide. Second method preparation - the interaction of iron with its oxides, for example, ferum (III) oxide, without the formation of any by-products.
However, for such a reaction it is necessary to create conditions in the form of high temperature - 900-1000 degrees Celsius. Third way- a reaction between iron and oxygen, in this case only iron (II) oxide is formed. To implement this process heating of the starting substances will also be required. Fourth method obtained is ferrous oxalate. This reaction requires high temperatures as well as a vacuum. As a result, ferum (II) oxide, carbon dioxide and carbon monoxide are formed in a ratio of 1:1:1. From what was written above we can conclude that the simplest and not requiring special conditions is the first method of obtaining this substance. Iron (II) oxide is used to smelt cast iron; it is also one of the components of some dyes and is used in the process of blackening steel.
Iron(III) oxide
This is no less common amphoteric oxide than the one described above. At normal conditions it is a solid substance that is red-brown in color. In nature it can be found in the form of the mineral hematite, which is used in the manufacture of jewelry. In industry, this substance received wide application: it is used to color some building materials, such as brick, paving slabs, etc., in the manufacture of paints, including printing, and enamels. The substance in question also serves as a food coloring called E172. In the chemical industry it is used in the production of ammonia as a catalyst.
Aluminium oxide
Amphoteric oxides also include aluminum oxide in their list. This substance under normal conditions has a solid state. The color of this oxide is white. In nature, part of it can be found in the form of alumina, as well as sapphire and ruby. Mainly used in chemical industry as a catalyst. But it is also used in the manufacture of ceramics.
Zinc oxide
This chemical compound is also amphoteric. This colorless solid is insoluble in water. It is obtained mainly through the decomposition of various zinc compounds. For example, its nitrate. This releases zinc oxide, nitrogen dioxide and oxygen. You can also extract this substance through the decomposition of zinc carbonate. In this reaction, in addition to the desired compound, carbon dioxide is also released. It is also possible for zinc hydroxide to decompose into its oxide and water. In order to carry out all three of the above processes, exposure to high temperature is required. Zinc oxide is used in various industries, for example, in the chemical industry (as a catalyst) for the manufacture of glass, in medicine for the treatment of skin defects.
Beryllium oxide
It is obtained mainly by thermal decomposition of the hydroxide of this element. This also produces water. It looks like a solid, colorless substance. This oxide finds its application in various industries as a heat-resistant material.
Tin oxide
It has dark color, has a solid state under normal conditions. It can be obtained, like many other amphoteric oxides, through the decomposition of its hydroxide. As a result, the substance in question and water are formed. This also requires exposure to high temperature. This compound is used in the chemical industry as a reducing agent in redox reactions, and less commonly used as a catalyst.
Properties, preparation and application of amphoteric hydroxides
Amphoteric hydroxides are used no less widely than oxides. Due to their versatile chemical behavior, they are mainly used for the preparation of all kinds of compounds. Additionally, iron hydroxide (a colorless solid) is used in the manufacture of batteries; aluminum hydroxide - for water purification; beryllium hydroxide - to obtain oxide.
Amphotericity (duality of properties) of hydroxides and oxides of many elements is manifested in the formation of two types of salts. For example, for aluminum hydroxide and oxide:
a) 2Al(OH)3 + 3SO3 = Al2(SO4)3 + 3H2O
Al2O3 + 3H2SO4 = Al2(SO4)3 + 3H2O
b) 2Al(OH)3 + Na2O = 2NaAlO2 + 3H2O (in melt)
Al2O3 + 2NaOH(s) = 2NaAlO2 + H2O (in melt)
In reactions (a), Al(OH)3 and Al2O3 exhibit the properties of basic hydroxides and oxides, that is, like alkalis, they react with acids and acidic oxides, forming a salt in which aluminum is the Al3+ cation.
On the contrary, in reactions (b) Al(OH)3 and Al2O3 perform the function of acidic hydroxides and oxides, forming a salt in which the aluminum atom AlIII is part of the anion (acid residue) AlO2−.
The element aluminum itself exhibits the properties of a metal and a non-metal in these compounds. Therefore, aluminum is an amphoteric element.
A-group elements - Be, Ga, Ge, Sn, Pb, Sb, Bi and others, as well as most B-group elements - Cr, Mn, Fe, Zn, Cd and others also have similar properties.
For example, the amphoteric nature of zinc is proven by the following reactions:
a) Zn(OH)2 + N2O5 = Zn(NO3)2 + H2O
ZnO + 2HNO3 = Zn(NO3)2 + H2O
b) Zn(OH)2 + Na2O = Na2ZnO2 + H2O
ZnO + 2NaOH(s) = Na2ZnO2 + H2O
If an amphoteric element has several oxidation states in compounds, then amphoteric properties are most clearly manifested for an intermediate oxidation state.
For example, chromium has three known oxidation states: +II, +III and +VI. In the case of CrIII, the acidic and basic properties are expressed approximately equally, while in CrII there is a predominance of basic properties, and in CrVI - acidic properties:
CrII → CrO, Cr(OH)2 → CrSO4
CrIII → Cr2O3, Cr(OH)3 → Cr2(SO4)3 or KCrO2
CrVI → CrO3, H2CrO4 → K2CrO4
Very often, amphoteric hydroxides of elements in the +III oxidation state also exist in meta form, for example:
AlO(OH) - aluminum metahydroxide
FeO(OH) is iron metahydroxide (the ortho form "Fe(OH)3" does not exist).
Amphoteric hydroxides are practically insoluble in water; the most convenient way to obtain them is precipitation from an aqueous solution using a weak base - ammonia hydrate:
Al(NO3)3 + 3(NH3 H2O) = Al(OH)3↓ + 3NH4NO3 (20 °C)
Al(NO3)3 + 3(NH3 H2O) = AlO(OH)↓ + 3NH4NO3 + H2O (80 °C)
If an excess of alkalis is used in an exchange reaction of this type, aluminum hydroxide will not precipitate, since aluminum, due to its amphoteric nature, transforms into an anion:
Al(OH)3(s) + OH− = −
Examples of molecular equations for reactions of this type:
Al(NO3)3 + 4NaOH(excess) = Na + 3NaNO3
ZnSO4 + 4NaOH(excess) = Na2 + Na2SO4
The resulting salts are classified as complex compounds (complex salts): they include complex anions − and 2−. The names of these salts are:
Na - sodium tetrahydroxyaluminate
Na2 - sodium tetrahydroxozincate
The reaction products of aluminum or zinc oxides with solid alkali are called differently:
NaAlO2 - sodium dioxoaluminate(III)
Na2ZnO2 - sodium dioxocincate(II)
Acidification of solutions of complex salts of this type leads to the destruction of complex anions:
− → Al(OH)3 → Al3+
For example: 2Na + CO2 = 2Al(OH)3↓ + NaHCO3
For many amphoteric elements, the exact formulas of hydroxides are unknown, since hydrated oxides, for example MnO2 nH2O, Sb2O5 nH2O, precipitate from an aqueous solution instead of hydroxides.
Amphoteric elements in free form interact with both typical acids and alkalis:
2Al + 3H2SO4(dil.) = Al2(SO4)3 + H2
2Al + 6H2O + 4NaOH(conc.) = 2Na + 3H2
In both reactions, salts are formed, and the element in question is part of the cation in one case, and part of the anion in the second.
Aluminum halides under normal conditions - colorless crystalline
substances. Among the aluminum halides, AlF3 differs greatly in properties
from their counterparts. It is refractory, slightly soluble in water, chemically
inactive. The main method for producing AlF3 is based on the action of anhydrous HF
on Al2O3 or Al:
Al2O3 + 6HF = 2AlF3 + 3H2O
Aluminum compounds with chlorine, bromine and iodine are fusible, very
reactive and highly soluble not only in water, but also in many
organic solvents. Interaction aluminum halides with water
accompanied by a significant release of heat. All of them are in aqueous solution
highly hydrolyzed, but unlike typical acid halides
non-metals, their hydrolysis is incomplete and reversible. Being noticeably volatile already
under normal conditions, AlCl3, AlBr3 and AlI3 smoke in humid air
(due to hydrolysis). They can be obtained by direct interaction
simple substances.
Complex halides(halogenometalates) contain complex anions, in which halogen atoms are ligands, for example. Potassium hexachloroplatinate(IV) K2, sodium heptafluorotantalate(V) Na, lithium hexafluoroarsenate(V) Li. max. thermal Fluoro-, oxofluoro- and chlorometalates are resistant. By the nature of the bonds, ionic compounds are close to complex halides. with cations NF4+, N2F3+, C1F2+, XeF+, etc.
Many halides are characterized by association and polymerization in the liquid and gas phases with the formation of bridging bonds. max. Group I and II metal halides, AlC13, Sb and transition metal pentafluorides, and MOF4 oxofluorides are prone to this. Halides with a metal-metal bond are known, for example. Hg2Cl2.
Fluorides differ significantly in their properties from other halides. However, in simple halides these differences are less pronounced than in the halogens themselves, and in complex halides they are less pronounced than in simple ones.
Many covalent halides (especially fluorides) have strong Lewis acids, for example. AsF5, SbF5, BF3, A1C13. Fluorides are part of superacids. Higher halides are reduced by metals and H2, for example:
Metal halides of groups V-VIII, except Cr and Mn, are reduced by H2 to metals, for example: WF6 + 3H2 -> W + 6HF
Many covalent and ionic metal halides interact with each other to form complex halides, for example: KC1 + TaC15 -> K[TaC16]
Lighter halogens can displace heavier halides. Oxygen can oxidize halides, releasing C12, Br2 and I2. One of the characteristic reactions of covalent halides is mutual. with water (hydrolysis) or its vapor when heated. (pyrohydrolysis), leading to the formation of oxides, oxy- or
oxohalides, hydroxides and hydrogen halides. The exceptions are CF4, CC14 and SF6, which are resistant to water vapor at high temperatures.
Halides are obtained directly from elements, interaction. hydrogen halides or hydrogen halides with elements, oxides, hydroxides or salts, as well as exchange solutions.
Halides are widely used in technology as starting materials for the production of halogens, alkaline and alkali-earth. metals, as components of glasses, etc. inorganic. materials; they are in between. products in the production of rare and certain non-ferrous metals, U, Si, Ge, etc.
In nature, halides form separate classes minerals, which include fluorides (for example, the minerals fluorite, cryolite) and chlorides (sylvite, carnallite). Bromine and iodine are included in the composition of certain minerals in the form of isomorphic impurities. Significant quantities of halides are contained in the water of seas and oceans, in salt and underground brines. Some halides, e.g. NaCl, K.C1, CaC12 are part of living organisms.
Cryolite(from ancient Greek κρύος - frost + λίθος - stone) - a rare mineral from the class of natural fluorides, sodium hexafluoroaluminate Na3. Crystallizes in the monoclinic system; cuboid crystals and twin plates are rare. Usually forms colorless, white or gray crystalline clusters with a glassy luster, often containing quartz, siderite, pyrite, galena, chalcopyrite, columbite, cassiterite. Coloring with impurities of organic substances is possible.
Methods have now been developed obtaining artificial cryolite. It is artificially obtained by the interaction of aluminum fluoride with sodium fluoride, as well as the action of hydrofluoric acid on aluminum hydroxide in the presence of soda. It is used in the process of electrolytic production of aluminum, in the production of hydrofluoric acid, glass and enamels.
Alum. Alum is the group name for double salts of the composition ME(SO4)2. 12H2O, where M is potassium K, rubidium Rb, cesium Cs, ammonium NH4, and E is aluminum Al, chromium Cr, iron Fe and other elements in the oxidation state (+III), which give triply charged cations during the dissociation of salts.
Alum is highly soluble in water; its aqueous solutions have an astringent, sour taste and an acidic reaction due to hydrolysis, for example:
3+ + H2O<<здесь знак обратимости >> 2+ + H3O+
When heated, alum first melts in the water it contains, and then this water is lost, forming anhydrous salts. Further heating converts the alum into a mixture of metal oxides. Aluminum-potassium alum can be obtained by modifying the production process of purified aluminum sulfate. First, kaolin is cooked with sulfuric acid. After neutralization of sulfuric acid is completed, sodium sulfate is added to the reactor to obtain sodium alum. The latter, due to their high solubility, are in solution. After diluting the solution to a density of 1.33 g/cm3, it is separated from the silica precipitate, cooled and mixed with a saturated solution of potassium chloride. In this case, aluminum-potassium alum, which is poorly soluble at low temperatures, precipitates. After separation of the crystals of aluminum-potassium alum, soluble impurities remain in the mother solution - iron compounds and sodium chloride 89.
During hydrolysis hydrated aluminum ions lose protons, forming successive hydro-oxo complexes. When the last neutral complex loses water, insoluble hydroxide A1(OH)3 is formed.
Complex ions[A1(H20)5OH]2+ and [A1(H20)4(OH)2]+ remain in solution, while A1(OH)3 hydroxide precipitates immediately after its formation. Precipitation occurs at pH values > 3. Completely before the formation of aluminum hydroxide hydrolysis occurs under the condition of neutralization of the resulting protons, for example with alkali.
Deep hydrolysis aluminum sulfate salts are widely used for purification of drinking and Wastewater. Hydronium released during hydrolysis reacts with the bicarbonates H30+ + HC03 = CO2 + 2H20 usually contained in water. In this case final products hydrolysis are colloidal aluminum hydroxide and carbon dioxide.
When the aluminum hydroxide sol coagulates, a voluminous gelatinous sediment is obtained, which captures suspended particles and bacteria and carries them to the bottom of the settling tank. The consumption of aluminum sulfate required for water purification depends on the composition and amount of contaminants in the water. Doses of aluminum sulfate for the purification of natural waters and for post-treatment of wastewater range from 3 to 15 mg/l according to A1203, and for the physicochemical treatment of urban wastewater they reach 30-50 mg/l according to A1203. The consumption of aluminum sulfate should ensure the formation of a sufficiently large mass of flakes, which is necessary to remove contaminants in it from the water. The pH value of the solution should be reduced to 6.5-7.6, which corresponds to the minimum solubility of aluminum hydroxide in water. At higher or lower pH values, some aluminum remains in the water in a dissolved state. In waters with low alkalinity, when the bicarbonate content is insufficient to neutralize the released acid, the hydrolysis process does not reach completion due to a strong decrease in pH. To increase alkalinity, complete the hydrolysis process and reduce the content of dissolved aluminum in water, lime and soda are introduced into the water simultaneously with the coagulant.
If the protons accumulated during hydrolysis are not neutralized, the hydrolysis process slows down, which leads to the onset of hydrolytic equilibrium, which can be characterized by the degree and constant of hydrolysis. Hydrolysis solutions of aluminum sulfate, which is the reaction of replacing sulfate ions in Al2(804)3 with OH ions formed due to the dissociation of water, can be represented in general form by the equation
2A13+ + (3 - -|-) EOG + aOH" + ad^ACONTSBOZH --^EOG + ad,
where a is the degree and basicity of substitution.
This equation shows that the concentration of OH- ions in the solution, i.e., the degree of dissociation of water, has a decisive influence on the shift to the right. As is known, for salts with a weak base and a strong acid, the degree of hydrolysis k is related to the hydrolysis constant A-, the salt concentration (c, mol"l), the ionic product of water kw and the dissociation constant of the base kw by the following relationship:
/g = UkTs = UkiLs.
If A-, changes little with temperature, then ksh increases significantly, which causes a significant increase in the degree of hydrolysis with increasing temperature.
N. I. Eremin, based on the experimental data obtained, derived equations for the dependence of the degree of hydrolysis of a solution on temperature and concentration
for aluminum sulfate:
1ек = - 2.23 + 0.05с + 0.0036т7 + 18 УЦс, for ammonium alum:
18 L = -1.19 +0.29s+ 0.0016G + 18ugSh for potassium alum:
\yok= - 1.17 + 0.29s + 0.00167 + 18 UPs,
for sodium alum:
18k = - 1.18 + 0.29s + 0.0016t7 + \е UPs.
As can be seen from these equations, the effect of concentration on the degree of hydrolysis is more significant for alum than for aluminum sulfate.
Bor. Obtaining boron. Chemical properties. Diagonal similarity of boron to silicon. Boron hydrides. Diboran. Features of the chemical bond in the diborane molecule. Boron halides. Oxygen compounds of boron. Boron oxide and boric acids. Borax. Preparation of boric acid. Borosilicate glasses. Bornoethyl ether.
Bor- element of the thirteenth group (according to the outdated classification - the main subgroup of the third group), the second period of the periodic table chemical elements with atomic number 5. Denoted by the symbol B (Latin: Borum). In the free state, boron is a colorless, gray or red crystalline or dark amorphous substance. More than 10 allotropic modifications of boron are known, the formation and mutual transitions of which are determined by the temperature at which the boron was obtained
Receipt. The purest boron is obtained by pyrolysis of borohydrides. This boron is used for the production of semiconductor materials and fine chemical syntheses.
Metallothermy method (usually reduction with magnesium or sodium):
Thermal decomposition of boron bromide vapor on a hot (1000-1200 °C) tungsten wire in the presence of hydrogen (Van Arkel method):
Physical properties. An extremely hard substance (second only to diamond, boron nitride (borazon), boron carbide, boron-carbon-silicon alloy, scandium-titanium carbide). Possesses fragility and semiconductor properties (wide-gap
semiconductor). Boron has the highest tensile strength of 5.7 GPa
In nature, boron is found in the form of two isotopes 10B (20%) and 11B (80%)[.
10B has a very high thermal neutron absorption cross section, so 10B in boric acid is used in nuclear reactors to regulate reactivity.
Chemical properties. Boron ions color the flame green.
In many physical and chemical properties, the nonmetal boron resembles silicon.
Chemically, boron is quite inert and at room temperature interacts only with fluorine:
When heated, boron reacts with other halogens to form trihalides, with nitrogen it forms boron nitride BN, with phosphorus - phosphide BP, with carbon - carbides of various compositions (B4C, B12C3, B13C2). When heated in an oxygen atmosphere or in air, boron burns with a large release of heat, forming the oxide B2O3:
Boron does not interact directly with hydrogen, although a fairly large number of borohydrides (boranes) of various compositions are known, obtained by treating borides of alkali or alkaline earth metals with acid:
When heated strongly, boron exhibits restorative properties. It is capable, for example, of reducing silicon or phosphorus from their oxides:
This property of boron can be explained by the very high strength of chemical bonds in boron oxide B2O3.
In the absence of oxidizing agents, boron is resistant to alkali solutions. In hot nitric and sulfuric acids and in aqua regia, boron dissolves to form boric acid
Boron oxide is a typical acidic oxide. It reacts with water to form boric acid:
When boric acid interacts with alkalis, salts are formed not of boric acid itself - borates (containing the BO33− anion), but tetraborates, for example:
Bor- semiconductor, diagonal similarity to silicon:
1) Both are refractory, solid, semiconductors. B – gray-black, Si – gray.
I1(B)=8.298 eV; I1(Si)=8.151 eV. Both do not tend to form cations.
2) Both are chemically inert (although boron does dissolve in hot oxidizing acids. Both dissolve in alkalis.
2B + KOH + 2H2O ® 2KBO2 + 3H2
Si + 2KOH + H2O®K2SiO3+ 2H2
3) At high temperatures they react with metals, forming borides and silicides - Ca3B2; Mg2Si - refractory, electrically conductive compounds.
Oxygen compounds of boron. B2O3 is an acidic oxide (SiO2 too) - both are polymeric, glassy, only B2O3 forms flat networks, and SiO2 forms three-dimensional structures. The difference between them is that boron oxide hydrates easily, while sand (SiO2) is known not to.
H3BO3 - orthoboric acid.
H3BO3«HBO2+H2Ometaboric acid (100оС)
4HBO2 «H2B4O7 + H2Otetraboric acid (140оС) - weak, both Kd
H2B4O7 «2B2O3 + H2O are almost the same - no acid salts
Orthoboric acid is weak, sometimes its dissociation is written
B(OH)3 + H2O « B(OH)4 + H+
Forms esters with alcohols: H3BO3+3CH3OH®B(OCH3)3+3H2O
Properties. Boron is known in amorphous (brown) and crystalline (black) forms, mp. 2300°C, bp. 3700°C, p = 2.34 g/cm3. The crystal lattice of boron is very strong, which is reflected in its high hardness, low entropy, and high melting point. Boron semiconductor. The non-metallicity of boron corresponds to its position in the periodic table - between beryllium and carbon and diagonally next to silicon. Therefore, boron shows similarities not only with aluminum, but also with silicon. From its position it also follows that compounds of boron with nitrogen should be similar in electronic structure and properties to carbon.
2ВН3(g) - В2Н6(g);
delta G= - 126 kJ
3NaBH4+4BF3 ->2B2H6 + 3NaBF4
6H2 (g) + 2ВС13 (g) ->В2Н6 (g) + 6НCl (g)
Diboran B2H6 is an energetic reducing agent; it spontaneously ignites in air
В2Н6+3О2 =>В2О3+ЗН2О
Reacts with water to release hydrogen;
В2Н6+6Н2О =>. 2H3VO3+6H2
In the ether medium, B2H6 reacts with lithium hydride, forming borohydride
B2H6+2LiH => 2LiBH4
More often than Li, they use Na, obtained by the reaction -
4NaH + B(OCH3)3 => Na + 3NaOCH3
B2O3 + ZS => 2B + ZSO
2B2O3+P4O10 => 4BPO4
H3BO3+H2O => [B(OH)4] + H
When neutralizing H3BO3, no orthoborates , containing the (BO3)3- ion, and the result is tetraborates, metaborates or salts of other polyboric acids:
4H3BO3 + 2NaOH => Na2BO4 + 7H2O H3BO3 + NaOH => NaBO2 + 2H2O
Boron oxide B2O3 - boric anhydride, colorless, rather refractory glassy or crystalline substance with a bitter taste, dielectric.
Glassy boron oxide has a layered structure (the distance between layers is 0.185 nm); in the layers, boron atoms are located inside equilateral triangles BO3 (d BO = 0.145 nm). This modification melts in the temperature range 325-450 °C and has high hardness. It is obtained by heating boron in air at 700 °C or by dehydrating orthoboric acid. Crystalline B2O3, which is obtained by careful elimination of water from metaboric acid HBO2, exists in two modifications - with a hexagonal crystal lattice, which turns into a monoclinic lattice at 400 °C and 2200 MPa.
In industry Borax is obtained from natural borates by fusing with soda . When natural boron minerals are treated with sulfuric acid, boric acid . From boric acid H3BO3, oxide B2O3 is obtained by calcination, and then it or borax is reduced with active metals (magnesium or sodium) to free boron:
B2O3 + 3Mg = 3MgO + 2B,
2Na2B4O7 + 3Na = B + 7NaBO2.
In this case, it forms in the form of a gray powder. amorphous boron. High-purity crystalline boron can be obtained by recrystallization, but in industry it is more often obtained by electrolysis of molten fluoroborates or thermal decomposition of boron bromide vapor BBr3 on tantalum wire heated to 1000-1500 °C in the presence of hydrogen:
2BBr3 + 3H2 = 2B + 6HBr
It is also possible to use cracking of borohydrides:
B4H10 = 4B + 5H2.
Boric acid(orthoboric acid) is a weak acid with the chemical formula H3BO3. A colorless, odorless, crystalline substance in the form of flakes, it has a layered triclinic lattice in which acid molecules are connected by hydrogen bonds into flat layers, the layers are connected to each other by intermolecular bonds (d = 0.318 nm).
Metaboric acid(HBO2) also occurs as colorless crystals. It exists in three modifications - the most stable γ-HBO2 with a cubic lattice, β-HBO2 with a monoclinic lattice and α-HBO2 with an orthorhombic lattice.
When heated orthoboric acid loses water and first transforms into metaboric acid, then into tetraboric acid H2B4O7. With further heating it dehydrates to boric anhydride.
Boric acid exhibits very weak acidic properties. It is relatively slightly soluble in water. Its acidic properties are due not to the abstraction of the H+ proton, but to the addition of a hydroxyl anion:
Ka = 5.8·10−10 mol/l; pKa = 9.24.
It is easily displaced from solutions of its salts by most other acids. Its salts, called borates, are usually produced from various polyboric acids, most often tetraboric H2B4O7, which is a much stronger acid than orthoboric acid. B(OH)3 shows very weak signs of amphotericity, forming low-stable boron hydrogen sulfate B(HSO4)3.
When orthoboric acid is neutralized with alkalis in aqueous solutions, orthoborates containing the (BO3)3− ion are not formed, since orthoborates are hydrolyzed almost completely due to the too low formation constant [B(OH)4]−. Tetraborates, metaborates or salts of other polyboric acids are formed in solution:
With an excess of alkali they can be converted into metaborates:
Meta- and tetraborates are hydrolyzed, but to a lesser extent (reactions opposite to those given).
In acidified aqueous solutions of borates, the following equilibria are established:
The most common boric acid salt is sodium tetraborate decahydrate Na2B4O7 10H2O (technical name borax).
When heated, boric acid dissolves metal oxides, forming salts.
With alcohols in the presence of concentrated sulfuric acid it forms esters:
Formation of boron methyl ether B(OCH3)3 is a qualitative reaction to H3BO3 and boric acid salts; when ignited, boron methyl ether burns with a beautiful bright green flame.
Borosilicate glass- glass of regular composition, in which the alkaline components in the feedstock are replaced with boron oxide (B2O3). This achieves increased chemical resistance and a low coefficient of thermal expansion - up to 3.3·10−6 at 20 °C for the best samples. In borosilicate glass it is very small, only smaller in quartz glass(almost 10 times). This prevents the glass from cracking during sudden temperature changes. This explains its use as a fire-fighting material and in other cases where thermal resistance is required.
Usage In everyday life, for making cookware for open fires, teapots. It is used as a material for laboratory glassware, as well as for the chemical industry and other industries, for example, as a heat exchanger material for thermal power plants. Also used for making cheap guitar slides. Borosilicate glass can also be used to make pipettes for ICSI, blastomere biopsy, which is carried out for preimplantation genetic diagnosis using biopsy cells as genetic material. There are 3 pipette options with internal diameters from 4 µm to 7.5 µm. The pipette length ranges from 60 to 75 mm and has a bevel angle of 30°. Pipettes are intended for single use.
general characteristics elements of the IVA subgroup. The structure of atoms. Oxidation states. Prevalence and forms of occurrence in nature. Allotropic modifications of carbon. Physical and chemical properties. Varieties of black graphite: coke, charcoal, soot.
General characteristics of group IVA elements The elements of the main subgroup of group IV include C, Si, Ge, Sn, Pv. The electronic formula of the outer valence level is nS2np2, that is, they have 4 valence electrons and these are p elements, therefore they are in the main subgroup of group IV. ││││ │↓│ np nS In the ground state of an atom, two electrons are paired and two are unpaired. The outermost electron shell of carbon has 2 electrons, silicon has 8, and Ge, Sn, Pv have 18 electrons each. Therefore, Ge, Sn, Pv are combined into the germanium subgroup (these are complete electronic analogues). In this subgroup of p-elements, as in other subgroups of p-elements, the properties of the atoms of the elements change periodically.
Thus, from top to bottom in the subgroup, the atomic radius increases, so the ionization energy decreases, so the ability to donate electrons increases, and the tendency to supplement the outer electron shell to the octet decreases sharply, so from C to Pb the reducing properties and metallic properties increase, and the nonmetallic properties decrease . Carbon and silicon are typical non-metals; Ge already exhibits metallic properties and appearance it is similar to metal, although it is a semiconductor. Tin already has metallic properties that predominate, while lead is a typical metal. Having 4 valence electrons, atoms in their compounds can exhibit oxidation states from minimum (-4) to maximum (+4), and they are characterized by even S.O.: -4, 0, +2, +4; S.O. = -4 is typical for C and Si with metals. The nature of the connection with other elements. Carbon forms only covalent bonds, silicon also predominantly forms covalent bonds. For tin and lead, especially in S.O. = +2, the ionic nature of the bond is more typical (for example, Рв(NO3)2). Covalency is determined by the valence structure of an atom. The carbon atom has 4 valence orbitals and the maximum covalency is 4. For other elements, the covalence can be more than four, since there is a valence d-sublevel (for example, H2). Hybridization. The type of hybridization is determined by the type and number of valence orbitals. Carbon has only S- and p-valence orbitals, so there can be Sp (carbyne, CO2, CS2), Sp2 (graphite, benzene, COCl2), Sp3 hybridization (CH4, diamond, CCl4). For silicon, the most characteristic Sp3 is hybridization (SiO2, SiCl4), but it has a valence d-sublevel, so there is also Sp3d2 hybridization, for example, H2. Group IV of PSE is the middle of D.I. Mendeleev’s table. A sharp change in properties from non-metals to metals is clearly visible here. Let us separately consider carbon, then silicon, then elements of the germanium subgroup.
Atom(from the Greek atomos - indivisible) - a single-nuclear, indivisible particle of a chemical element, a carrier of the properties of a substance. Substances are made up of atoms. The atom itself consists of a positively charged nucleus and a negatively charged electron cloud. In general, the atom is electrically neutral. The size of an atom is entirely determined by the size of its electron cloud, since the size of the nucleus is negligible compared to the size of the electron cloud. The nucleus consists of Z positively charged protons (the charge of a proton corresponds to +1 in arbitrary units) and N neutrons, which do not carry a charge (protons and neutrons are called nucleons). Thus, the charge of the nucleus is determined only by the number of protons and is equal to the ordinal number of the element in the periodic table. The positive charge of the nucleus is compensated by negatively charged electrons (electron charge -1 in arbitrary units), which form an electron cloud. The number of electrons is equal to the number of protons. The masses of protons and neutrons are equal (1 and 1 amu, respectively). The mass of an atom is determined by the mass of its nucleus, since the mass of an electron is approximately 1850 times less than the mass of a proton and neutron and is rarely taken into account in calculations. The number of neutrons can be determined by the difference between the mass of an atom and the number of protons (N=A-Z). A type of atom of a chemical element with a nucleus consisting of strictly a certain number protons (Z) and neutrons (N) is called a nuclide.
Since almost all the mass is concentrated in the nucleus of an atom, but its dimensions are negligible compared to the total volume of the atom, the nucleus is conventionally accepted material point resting at the center of the atom, and the atom itself is considered as a system of electrons. During a chemical reaction, the nucleus of an atom is not affected (except nuclear reactions), as well as internal electronic levels, and only electrons of the outer electron shell participate. For this reason, it is necessary to know the properties of the electron and the rules for the formation of electron shells of atoms.
Oxidation state(oxidation number, formal charge) - an auxiliary conventional value for recording the processes of oxidation, reduction and redox reactions. It indicates the oxidation state of an individual atom of a molecule and is only a convenient method of accounting for electron transfer: it is not the true charge of an atom in the molecule (see #Conventions).
Ideas about the oxidation state of elements form the basis and are used in classification chemical substances, description of their properties, drawing up formulas of compounds and their international titles(nomenclature). But it is especially widely used in the study of redox reactions.
The concept of oxidation state is often used in inorganic chemistry instead of the concept of valency.
The oxidation state of an atom is equal to the numerical value of the electrical charge assigned to the atom, assuming that the bonding electron pairs are completely biased towards more electronegative atoms (that is, assuming that the compound consists only of ions).
The oxidation number corresponds to the number of electrons that must be added to a positive ion to reduce it to a neutral atom, or subtracted from a negative ion to oxidize it to a neutral atom:
Al3+ + 3e− → Al
S2− → S + 2e− (S2− − 2e− → S)
Carbon- substance with the most [source not specified 1528 days] a large number allotropic modifications (more than 8 have already been discovered).
Allotropic modifications of carbon in their properties they differ most radically from each other, from soft to hard, opaque to transparent, abrasive to lubricating, inexpensive to expensive. These allotropes include amorphous carbon allotropes (coal, soot), nanofoam, crystalline allotropes - nanotube, diamond, fullerenes, graphite, lonsdaleite and ceraphite.
Classification of carbon allotropes according to the nature of the chemical bond between atoms:
Diamond (cube)
Lonsdaleite (hexagonal diamond)
Fullerenes (C20+)
Nanotubes
Nanofibers
Astralens
Glassy carbon
Colossal nanotubes
Mixed sp3/sp2 forms:
Amorphous carbon
Carbon nanobuds
Carbon nanofoam
Other shapes: C1 - C2 - C3 - C8
Carbon(chemical symbol - C, lat. Carboneum) - chemical element of the fourteenth group (according to the outdated classification - the main subgroup of the fourth
group), 2nd period of the periodic table of chemical elements. serial number 6, atomic mass - 12.0107.
Physical properties.
Carbon exists in a variety of allotropes with very diverse physical properties. The variety of modifications is due to the ability of carbon to form chemical bonds of different types.
Bases, amphoteric hydroxides
Bases are complex substances consisting of metal atoms and one or more hydroxyl groups (-OH). General formula Me +y (OH) y, where y is the number of hydroxo groups equal to the oxidation state of the metal Me. The table shows the classification of bases.
Properties of alkalis, hydroxides of alkali and alkaline earth metals
1. Aqueous solutions of alkalis are soapy to the touch and change the color of indicators: litmus - blue, phenolphthalein - crimson.
2. Aqueous solutions dissociate:
3. Interact with acids, entering into an exchange reaction:
Polyacid bases can give medium and basic salts:
4. React with acidic oxides, forming medium and acidic salts depending on the basicity of the acid corresponding to this oxide:
5. Interact with amphoteric oxides and hydroxides:
a) fusion:
b) in solutions:
6. Interact with water-soluble salts if a precipitate or gas is formed:
Insoluble bases (Cr(OH) 2, Mn(OH) 2, etc.) interact with acids and decompose when heated:
Amphoteric hydroxides
Amphoteric compounds are compounds that, depending on conditions, can be both donors of hydrogen cations and exhibit acidic properties, and their acceptors, i.e., exhibit basic properties.
Chemical properties of amphoteric compounds
1. Interacting with strong acids, they exhibit basic properties:
Zn(OH) 2 + 2HCl = ZnCl 2 + 2H 2 O
2. Interacting with alkalis - strong bases, they exhibit acidic properties:
Zn(OH) 2 + 2NaOH = Na 2 ( complex salt)
Al(OH) 3 + NaOH = Na ( complex salt)
Complex compounds are those in which at least one covalent bond is formed by a donor-acceptor mechanism.
The general method for preparing bases is based on exchange reactions, with the help of which both insoluble and soluble bases can be obtained.
CuSO 4 + 2KOH = Cu(OH) 2 ↓ + K 2 SO 4
K 2 CO 3 + Ba(OH) 2 = 2 KOH + BaCO 3 ↓
When soluble bases are obtained by this method, an insoluble salt precipitates.
When preparing water-insoluble bases with amphoteric properties, excess alkali should be avoided, since dissolution of the amphoteric base may occur, for example:
AlCl 3 + 4KOH = K[Al(OH) 4 ] + 3KCl
In such cases, ammonium hydroxide is used to obtain hydroxides, in which amphoteric hydroxides do not dissolve:
AlCl 3 + 3NH 3 + ZH 2 O = Al(OH) 3 ↓ + 3NH 4 Cl
Silver and mercury hydroxides decompose so easily that when trying to obtain them by exchange reaction, instead of hydroxides, oxides precipitate:
2AgNO 3 + 2KOH = Ag 2 O↓ + H 2 O + 2KNO 3
In industry, alkalis are usually obtained by electrolysis of aqueous solutions of chlorides.
2NaCl + 2H 2 O → ϟ → 2NaOH + H 2 + Cl 2
Alkalis can also be obtained by reacting alkali and alkaline earth metals or their oxides with water.
2Li + 2H 2 O = 2LiOH + H 2
SrO + H 2 O = Sr(OH) 2
Acids
Acids are complex substances whose molecules consist of hydrogen atoms that can be replaced by metal atoms and acidic residues. Under normal conditions, acids can be solid (phosphoric H 3 PO 4; silicon H 2 SiO 3) and liquid (in its pure form, sulfuric acid H 2 SO 4 will be a liquid).
Gases such as hydrogen chloride HCl, hydrogen bromide HBr, hydrogen sulfide H 2 S form the corresponding acids in aqueous solutions. The number of hydrogen ions formed by each acid molecule during dissociation determines the charge of the acid residue (anion) and the basicity of the acid.
According to protolytic theory of acids and bases, proposed simultaneously by the Danish chemist Brønsted and the English chemist Lowry, an acid is a substance splitting off with this reaction protons, A basis- a substance that can accept protons.
acid → base + H +
Based on such ideas, it is clear basic properties of ammonia, which, due to the presence of a lone electron pair at the nitrogen atom, effectively accepts a proton when interacting with acids, forming an ammonium ion through a donor-acceptor bond.
HNO 3 + NH 3 ⇆ NH 4 + + NO 3 —
acid base acid base
More general definition acids and bases proposed by the American chemist G. Lewis. He suggested that acid-base interactions are completely do not necessarily occur with the transfer of protones. In the determination of acids and bases according to Lewis, the main role is in chemical reactions is given electron pairs
Cations, anions, or neutral molecules that can accept one or more pairs of electrons are called Lewis acids.
For example, aluminum fluoride AlF 3 is an acid, since it is able to accept an electron pair when interacting with ammonia.
AlF 3 + :NH 3 ⇆ :
Cations, anions, or neutral molecules capable of donating electron pairs are called Lewis bases (ammonia is a base).
Lewis's definition covers all acid-base processes that were considered by previously proposed theories. The table compares the definitions of acids and bases currently used.
Nomenclature of acids
Since there are different definitions of acids, their classification and nomenclature are rather arbitrary.
According to the number of hydrogen atoms capable of elimination in an aqueous solution, acids are divided into monobasic(e.g. HF, HNO 2), dibasic(H 2 CO 3, H 2 SO 4) and tribasic(H 3 PO 4).
According to the composition of the acid, they are divided into oxygen-free(HCl, H 2 S) and oxygen-containing(HClO 4, HNO 3).
Usually names of oxygen-containing acids are derived from the name of the non-metal with the addition of the endings -kai, -vaya, if the oxidation state of the non-metal is equal to the group number. As the oxidation state decreases, the suffixes change (in order of decreasing oxidation state of the metal): -opaque, rusty, -ovish:
If we consider the polarity of the hydrogen-nonmetal bond within a period, we can easily relate the polarity of this bond to the position of the element in the Periodic Table. From metal atoms, which easily lose valence electrons, hydrogen atoms accept these electrons, forming a stable two-electron shell like the shell of a helium atom, and give ionic metal hydrides.
In hydrogen compounds of elements of groups III-IV of the Periodic Table, boron, aluminum, carbon, and silicon form covalent, weakly polar bonds with hydrogen atoms that are not prone to dissociation. For elements of groups V-VII of the Periodic Table, within a period, the polarity of the nonmetal-hydrogen bond increases with the charge of the atom, but the distribution of charges in the resulting dipole is different than in hydrogen compounds of elements that tend to donate electrons. Non-metal atoms, which require several electrons to complete the electron shell, attract (polarize) a pair of bonding electrons the more strongly, the greater the nuclear charge. Therefore, in the series CH 4 - NH 3 - H 2 O - HF or SiH 4 - PH 3 - H 2 S - HCl, bonds with hydrogen atoms, while remaining covalent, become more polar in nature, and the hydrogen atom in the element-hydrogen bond dipole becomes more electropositive. If polar molecules find themselves in a polar solvent, a process of electrolytic dissociation can occur.
Let us discuss the behavior of oxygen-containing acids in aqueous solutions. These acids have an H-O-E bond and, naturally, the polarity of the H-O bond is affected by O-E connection. Therefore, these acids, as a rule, dissociate more easily than water.
H 2 SO 3 + H 2 O ⇆ H 3 O + + HSO 3
HNO 3 + H 2 O ⇆ H 3 O + + NO 3
Let's look at a few examples properties of oxygen-containing acids, formed by elements that are capable of exhibiting different degrees of oxidation. It is known that hypochlorous acid HClO very weak chlorous acid HClO 2 also weak, but stronger than hypochlorous, hypochlorous acid HClO 3 strong. Perchloric acid HClO 4 is one of the strongest inorganic acids.
For acid-type dissociation (with elimination of the H ion), a rupture is required O-N connections. How can we explain the decrease in the strength of this bond in the series HClO - HClO 2 - HClO 3 - HClO 4? In this series, the number of oxygen atoms associated with the central chlorine atom increases. Each time a new oxygen-chlorine bond is formed, electron density is drawn from the chlorine atom, and therefore from the O-Cl single bond. As a result, the electron density partially leaves the O-H bond, which is weakened as a result.
This pattern - strengthening of acidic properties with increasing degree of oxidation of the central atom - characteristic not only of chlorine, but also of other elements. For example, nitric acid HNO 3, in which the oxidation state of nitrogen is +5, is stronger than nitrous acid HNO 2 (the oxidation state of nitrogen is +3); sulfuric acid H 2 SO 4 (S +6) is stronger than sulfurous acid H 2 SO 3 (S +4).
Obtaining acids
1. Oxygen-free acids can be obtained by direct combination of non-metals with hydrogen.
H 2 + Cl 2 → 2HCl,
H 2 + S ⇆ H 2 S
2. Some oxygen-containing acids can be obtained interaction of acid oxides with water.
3. Both oxygen-free and oxygen-containing acids can be obtained by metabolic reactions between salts and other acids.
BaBr 2 + H 2 SO 4 = BaSO 4 ↓ + 2НВr
CuSO 4 + H 2 S = H 2 SO 4 + CuS↓
FeS + H 2 SO 4 (pa zb) = H 2 S + FeSO 4
NaCl (T) + H 2 SO 4 (conc) = HCl + NaHSO 4
AgNO 3 + HCl = AgCl↓ + HNO 3
CaCO 3 + 2HBr = CaBr 2 + CO 2 + H 2 O
4. Some acids can be obtained using redox reactions.
H 2 O 2 + SO 2 = H 2 SO 4
3P + 5HNO 3 + 2H 2 O = ZN 3 PO 4 + 5NO 2
Sour taste, effect on indicators, electrical conductivity, interaction with metals, basic and amphoteric oxides, bases and salts, formation of esters with alcohols - these properties are common to inorganic and organic acids.
can be divided into two types of reactions:
1) are common For acids reactions are associated with the formation of hydronium ion H 3 O + in aqueous solutions;
2) specific(i.e. characteristic) reactions specific acids.
The hydrogen ion can enter into redox reaction, reducing to hydrogen, as well as in a compound reaction with negatively charged or neutral particles having lone pairs of electrons, i.e. acid-base reactions.
TO general properties acids include reactions of acids with metals in the voltage series up to hydrogen, for example:
Zn + 2Н + = Zn 2+ + Н 2
Acid-base reactions include reactions with basic oxides and bases, as well as with intermediate, basic, and sometimes acidic salts.
2 CO 3 + 4HBr = 2CuBr 2 + CO 2 + 3H 2 O
Mg(HCO 3) 2 + 2HCl = MgCl 2 + 2CO 2 + 2H 2 O
2KHSO 3 + H 2 SO 4 = K 2 SO 4 + 2SO 2 + 2H 2 O
Note that polybasic acids dissociate stepwise, and at each subsequent step the dissociation is more difficult, therefore, with an excess of acid, acidic salts are most often formed, rather than average ones.
Ca 3 (PO 4) 2 + 4H 3 PO 4 = 3Ca (H 2 PO 4) 2
Na 2 S + H 3 PO 4 = Na 2 HPO 4 + H 2 S
NaOH + H 3 PO 4 = NaH 2 PO 4 + H 2 O
KOH + H 2 S = KHS + H 2 O
At first glance, the formation of acid salts may seem surprising monobasic hydrofluoric acid. However, this fact can be explained. Unlike all other hydrohalic acids, hydrofluoric acid in solutions is partially polymerized (due to the formation of hydrogen bonds) and various particles (HF) X may be present in it, namely H 2 F 2, H 3 F 3, etc.
A special case of acid-base equilibrium - reactions of acids and bases with indicators that change their color depending on the acidity of the solution. Indicators are used in qualitative analysis to detect acids and bases in solutions.
The most commonly used indicators are litmus(V neutral environment purple, V sour - red, V alkaline - blue), methyl orange(V sour environment red, V neutral - orange, V alkaline - yellow), phenolphthalein(V highly alkaline environment raspberry red, V neutral and acidic - colorless).
Specific properties different acids can be of two types: firstly, reactions leading to the formation insoluble salts, and secondly, redox transformations. If the reactions associated with the presence of the H + ion are common to all acids (qualitative reactions for detecting acids), specific reactions are used as qualitative reactions for individual acids:
Ag + + Cl - = AgCl (white precipitate)
Ba 2+ + SO 4 2- = BaSO 4 (white precipitate)
3Ag + + PO 4 3 - = Ag 3 PO 4 (yellow precipitate)
Some specific reactions of acids are due to their redox properties.
Anoxic acids in an aqueous solution can only be oxidized.
2KMnO 4 + 16HCl = 5Сl 2 + 2КСl + 2МnСl 2 + 8Н 2 O
H 2 S + Br 2 = S + 2НВг
Oxygen-containing acids can be oxidized only if the central atom in them is in a lower or intermediate oxidation state, as, for example, in sulfurous acid:
H 2 SO 3 + Cl 2 + H 2 O = H 2 SO 4 + 2HCl
Many oxygen-containing acids, in which the central atom has the maximum oxidation state (S +6, N +5, Cr +6), exhibit the properties of strong oxidizing agents. Concentrated H 2 SO 4 is a strong oxidizing agent.
Cu + 2H 2 SO 4 (conc) = CuSO 4 + SO 2 + 2H 2 O
Pb + 4HNO 3 = Pb(NO 3) 2 + 2NO 2 + 2H 2 O
C + 2H 2 SO 4 (conc) = CO 2 + 2SO 2 + 2H 2 O
It should be remembered that:
- Acid solutions react with metals that are to the left of hydrogen in the electrochemical voltage series, subject to a number of conditions, the most important of which is the formation of a soluble salt as a result of the reaction. The interaction of HNO 3 and H 2 SO 4 (conc.) with metals proceeds differently.
Concentrated sulfuric acid in the cold passivates aluminum, iron, and chromium.
- In water, acids dissociate into hydrogen cations and anions of acid residues, for example:
- Inorganic and organic acids react with basic and amphoteric oxides, provided that a soluble salt is formed:
- Both acids react with bases. Polybasic acids can form both intermediate and acid salts (these are neutralization reactions):
- The reaction between acids and salts occurs only if a precipitate or gas is formed:
The interaction of H 3 PO 4 with limestone will stop due to the formation of the last insoluble precipitate of Ca 3 (PO 4) 2 on the surface.
The peculiarities of the properties of nitric HNO 3 and concentrated sulfuric H 2 SO 4 (conc.) acids are due to the fact that when they interact with simple substances (metals and non-metals), the oxidizing agents will not be H + cations, but nitrate and sulfate ions. It is logical to expect that as a result of such reactions, not hydrogen H2 is formed, but other substances are obtained: necessarily salt and water, as well as one of the products of the reduction of nitrate or sulfate ions, depending on the concentration of acids, the position of the metal in the voltage series and reaction conditions (temperature, degree of metal grinding, etc.).
These features of the chemical behavior of HNO 3 and H 2 SO 4 (conc.) clearly illustrate the thesis of the theory chemical structure about the mutual influence of atoms in the molecules of substances.
The concepts of volatility and stability (stability) are often confused. Volatile acids are acids whose molecules easily pass into a gaseous state, that is, evaporate. For example, hydrochloric acid is a volatile but stable acid. It is impossible to judge the volatility of unstable acids. For example, non-volatile, insoluble silicic acid decomposes into water and SiO 2. Aqueous solutions of hydrochloric, nitric, sulfuric, phosphoric and a number of other acids are colorless. Water solution chromic acid H 2 CrO 4 is yellow in color, manganese acid HMnO 4 is crimson.
Reference material for taking the test:
Mendeleev table
Solubility table
