Before we start talking about Chemical properties oxides, you need to remember that all oxides are divided into 4 types, namely basic, acidic, amphoteric and non-salt-forming. In order to determine the type of any oxide, first of all you need to understand whether it is a metal or non-metal oxide in front of you, and then use the algorithm (you need to learn it!) presented in the following table:

Non-metal oxide	Metal oxide
1) Oxidation state of non-metal +1 or +2 Conclusion: non-salt-forming oxide Exception: Cl 2 O is not a non-salt-forming oxide	1) Metal oxidation state +1 or +2 Conclusion: metal oxide is basic Exception: BeO, ZnO and PbO are not basic oxides
2) The oxidation state is greater than or equal to +3 Conclusion: acid oxide Exception: Cl 2 O is an acidic oxide, despite the oxidation state of chlorine +1	2) Metal oxidation state +3 or +4 Conclusion: amphoteric oxide Exception: BeO, ZnO and PbO are amphoteric, despite the +2 oxidation state of the metals 3) Metal oxidation state +5, +6, +7 Conclusion: acid oxide

In addition to the types of oxides indicated above, we will also introduce two more subtypes of basic oxides, based on their chemical activity, namely active basic oxides And low-active basic oxides.

• TO active basic oxides We include oxides of alkali and alkaline earth metals (all elements of groups IA and IIA, except hydrogen H, beryllium Be and magnesium Mg). For example, Na 2 O, CaO, Rb 2 O, SrO, etc.
• TO low-active basic oxides we will include all the main oxides that are not included in the list active basic oxides. For example, FeO, CuO, CrO, etc.

It is logical to assume that active basic oxides often enter into reactions that low-active ones do not.
It should be noted that despite the fact that water is actually an oxide of a non-metal (H 2 O), its properties are usually considered in isolation from the properties of other oxides. This is due to its specifically huge distribution in the world around us, and therefore in most cases water is not a reagent, but a medium in which countless activities can take place. chemical reactions. However, it often takes a direct part in various transformations, in particular, some groups of oxides react with it.

Which oxides react with water?

Of all the oxides with water react only:
1) all active basic oxides (oxides of alkali metal and alkali metal);
2) all acid oxides, except silicon dioxide (SiO 2);

those. From the above it follows that with water exactly don't react:
1) all low-active basic oxides;
2) all amphoteric oxides;
3) non-salt-forming oxides (NO, N 2 O, CO, SiO).

The ability to determine which oxides can react with water even without the ability to write the corresponding reaction equations already allows you to get points for some questions in the test part of the Unified State Exam.

Now let's figure out how certain oxides react with water, i.e. Let's learn to write the corresponding reaction equations.

Active basic oxides, reacting with water, form their corresponding hydroxides. Recall that the corresponding metal oxide is a hydroxide that contains the metal in the same oxidation state as the oxide. So, for example, when the active basic oxides K +1 2 O and Ba +2 O react with water, their corresponding hydroxides K +1 OH and Ba +2 (OH) 2 are formed:

K2O + H2O = 2KOH– potassium hydroxide

BaO + H 2 O = Ba(OH) 2– barium hydroxide

All hydroxides corresponding to active basic oxides (alkaline metal and alkali metal oxides) belong to alkalis. Alkalis are all metal hydroxides that are highly soluble in water, as well as poorly soluble calcium hydroxide Ca(OH) 2 (as an exception).

The interaction of acidic oxides with water, as well as the reaction of active basic oxides with water, leads to the formation of the corresponding hydroxides. Only in the case of acidic oxides do they correspond not to basic ones, but to acidic hydroxides, more often called oxygen-containing acids. Let us recall that the corresponding acidic oxide is an oxygen-containing acid that contains an acid-forming element in the same oxidation state as in the oxide.

Thus, if we, for example, want to write down the equation for the interaction of acidic oxide SO 3 with water, first of all we must remember the basic ones studied within school curriculum, sulfur-containing acids. These are hydrogen sulfide H 2 S, sulfurous H 2 SO 3 and sulfuric H 2 SO 4 acids. Hydrogen sulfide acid H 2 S, as is easy to see, is not oxygen-containing, so its formation during the interaction of SO 3 with water can be immediately excluded. Of the acids H 2 SO 3 and H 2 SO 4, sulfur in the oxidation state +6, as in SO 3 oxide, contains only sulfuric acid H2SO4. Therefore, it is precisely this that will be formed in the reaction of SO 3 with water:

H 2 O + SO 3 = H 2 SO 4

Similarly, the oxide N 2 O 5, containing nitrogen in the oxidation state +5, reacting with water, forms nitric acid HNO 3, but in no case nitrous HNO 2, since in nitric acid the oxidation state of nitrogen is the same as in N 2 O 5 , is equal to +5, and in nitrogen - +3:

N +5 2 O 5 + H 2 O = 2HN +5 O 3

Interaction of oxides with each other

First of all, you need to clearly understand the fact that among salt-forming oxides (acidic, basic, amphoteric), reactions almost never occur between oxides of the same class, i.e. In the vast majority of cases, interaction is impossible:

1) basic oxide + basic oxide ≠

2) acid oxide + acid oxide ≠

3) amphoteric oxide + amphoteric oxide ≠

While interaction is almost always possible between oxides belonging to different types, i.e. almost always are leaking reactions between:

1) basic oxide and acidic oxide;

2) amphoteric oxide and acid oxide;

3) amphoteric oxide and basic oxide.

As a result of all such interactions, the product is always average (normal) salt.

Let us consider all these pairs of interactions in more detail.

As a result of the interaction:

Me x O y + acid oxide, where Me x O y – metal oxide (basic or amphoteric)

a salt is formed consisting of the metal cation Me (from the initial Me x O y) and the acid residue of the acid corresponding to the acid oxide.

As an example, let’s try to write down the interaction equations for the following pairs of reagents:

Na 2 O + P 2 O 5 And Al 2 O 3 + SO 3

In the first pair of reagents we see a basic oxide (Na 2 O) and an acidic oxide (P 2 O 5). In the second - amphoteric oxide (Al 2 O 3) and acidic oxide (SO 3).

As already mentioned, as a result of the interaction of a basic/amphoteric oxide with an acidic one, a salt is formed, consisting of a metal cation (from the original basic/amphoteric oxide) and an acidic residue of the acid corresponding to the original acidic oxide.

Thus, the interaction of Na 2 O and P 2 O 5 should form a salt consisting of Na + cations (from Na 2 O) and the acidic residue PO 4 3-, since the oxide P +5 2 O 5 corresponds to acid H 3 P +5 O4. Those. As a result of this interaction, sodium phosphate is formed:

3Na 2 O + P 2 O 5 = 2Na 3 PO 4- sodium phosphate

In turn, the interaction of Al 2 O 3 and SO 3 should form a salt consisting of Al 3+ cations (from Al 2 O 3) and the acidic residue SO 4 2-, since the oxide S +6 O 3 corresponds to acid H 2 S +6 O4. Thus, as a result of this reaction, aluminum sulfate is obtained:

Al 2 O 3 + 3SO 3 = Al 2 (SO 4) 3- aluminum sulfate

More specific is the interaction between amphoteric and basic oxides. These reactions are carried out at high temperatures, and their occurrence is possible due to the fact that the amphoteric oxide actually takes on the role of an acidic one. As a result of this interaction, a salt of a specific composition is formed, consisting of a metal cation forming the original basic oxide and an “acid residue”/anion, which includes the metal from the amphoteric oxide. The formula of such an “acid residue”/anion is general view can be written as MeO 2 x -, where Me is a metal from an amphoteric oxide, and x = 2 in the case of amphoteric oxides with general formula type Me +2 O (ZnO, BeO, PbO) and x = 1 – for amphoteric oxides with a general formula of the form Me +3 2 O 3 (for example, Al 2 O 3, Cr 2 O 3 and Fe 2 O 3).

Let's try to write down the interaction equations as an example

ZnO + Na 2 O And Al 2 O 3 + BaO

In the first case, ZnO is an amphoteric oxide with the general formula Me +2 O, and Na 2 O is a typical basic oxide. According to the above, as a result of their interaction, a salt should be formed, consisting of a metal cation forming a basic oxide, i.e. in our case, Na + (from Na 2 O) and the “acid residue”/anion with the formula ZnO 2 2-, since the amphoteric oxide has a general formula of the form Me + 2 O. Thus, the formula of the resulting salt, subject to the condition of electrical neutrality of one of its structural units (“molecules”) will look like Na 2 ZnO 2:

ZnO + Na 2 O = t o=> Na 2 ZnO 2

In the case of an interacting pair of reagents Al 2 O 3 and BaO, the first substance is an amphoteric oxide with the general formula of the form Me + 3 2 O 3, and the second is a typical basic oxide. In this case, a salt is formed containing a metal cation from the main oxide, i.e. Ba 2+ (from BaO) and the “acid residue”/anion AlO 2 - . Those. the formula of the resulting salt, subject to the condition of electrical neutrality of one of its structural units (“molecules”), will have the form Ba(AlO 2) 2, and the interaction equation itself will be written as:

Al 2 O 3 + BaO = t o=> Ba(AlO 2) 2

As we wrote above, the reaction almost always occurs:

Me x O y + acid oxide,

where Me x O y is either a basic or amphoteric metal oxide.

However, there are two “finicky” acid oxides to remember: carbon dioxide(CO 2) and sulfur dioxide (SO 2). Their “fastidiousness” lies in the fact that despite their obvious acidic properties, the activity of CO 2 and SO 2 is not enough for their interaction with low-active basic and amphoteric oxides. Of the metal oxides, they react only with active basic oxides(oxides of alkali metal and alkali metal). For example, Na 2 O and BaO, being active basic oxides, can react with them:

CO 2 + Na 2 O = Na 2 CO 3

SO 2 + BaO = BaSO 3

While the oxides CuO and Al 2 O 3, which are not related to active basic oxides, do not react with CO 2 and SO 2:

CO 2 + CuO ≠

CO 2 + Al 2 O 3 ≠

SO 2 + CuO ≠

SO 2 + Al 2 O 3 ≠

Interaction of oxides with acids

Basic and amphoteric oxides react with acids. In this case, salts and water are formed:

FeO + H 2 SO 4 = FeSO 4 + H 2 O

Non-salt-forming oxides do not react with acids at all, and acidic oxides do not react with acids in most cases.

When does an acidic oxide react with an acid?

Deciding part of the Unified State Exam with answer options, you should assume that acidic oxides do not react with either acidic oxides or acids, except in the following cases:

1) silicon dioxide, being an acidic oxide, reacts with hydrofluoric acid, dissolving in it. In particular, thanks to this reaction, glass can be dissolved in hydrofluoric acid. In the case of excess HF, the reaction equation has the form:

SiO 2 + 6HF = H 2 + 2H 2 O,

and in case of HF deficiency:

SiO 2 + 4HF = SiF 4 + 2H 2 O

2) SO 2, being an acidic oxide, easily reacts with hydrosulfide acid H 2 S like co-proportionation:

S +4 O 2 + 2H 2 S -2 = 3S 0 + 2H 2 O

3) Phosphorus (III) oxide P 2 O 3 can react with oxidizing acids, which include concentrated sulfuric acid and Nitric acid any concentration. In this case, the oxidation state of phosphorus increases from +3 to +5:

P2O3	+	2H2SO4	+	H2O	=t o=>	2SO 2	+	2H3PO4
		(conc.)

3 P2O3	+	4HNO3	+	7 H2O	=t o=>	4NO	+	6 H3PO4
		(detailed)

2HNO3	+	3SO 2	+	2H2O	=t o=>	3H2SO4	+	2NO
(detailed)

Interaction of oxides with metal hydroxides

Acidic oxides react with metal hydroxides, both basic and amphoteric. This produces a salt consisting of a metal cation (from the original metal hydroxide) and an acid residue corresponding to the acid oxide.

SO 3 + 2NaOH = Na 2 SO 4 + H 2 O

Acidic oxides, which correspond to polybasic acids, can form both normal and acid salts with alkalis:

CO 2 + 2NaOH = Na 2 CO 3 + H 2 O

CO 2 + NaOH = NaHCO 3

P 2 O 5 + 6KOH = 2K 3 PO 4 + 3H 2 O

P 2 O 5 + 4KOH = 2K 2 HPO 4 + H 2 O

P 2 O 5 + 2KOH + H 2 O = 2KH 2 PO 4

“Finicky” oxides CO 2 and SO 2, the activity of which, as already mentioned, is not enough for their reaction with low-active basic and amphoteric oxides, nevertheless, react with most of the corresponding metal hydroxides. More precisely, carbon dioxide and sulfur dioxide react with insoluble hydroxides in the form of their suspension in water. In this case, only the basic O natural salts called hydroxycarbonates and hydroxosulfites, and the formation of intermediate (normal) salts is impossible:

2Zn(OH) 2 + CO 2 = (ZnOH) 2 CO 3 + H 2 O(in solution)

2Cu(OH) 2 + CO 2 = (CuOH) 2 CO 3 + H 2 O(in solution)

However, carbon dioxide and sulfur dioxide do not react at all with metal hydroxides in the oxidation state +3, for example, such as Al(OH) 3, Cr(OH) 3, etc.

It should also be noted that silicon dioxide (SiO 2) is particularly inert, most often found in nature in the form of ordinary sand. This oxide is acidic, but among metal hydroxides it is capable of reacting only with concentrated (50-60%) solutions of alkalis, as well as with pure (solid) alkalis during fusion. In this case, silicates are formed:

2NaOH + SiO 2 = t o=> Na 2 SiO 3 + H 2 O

Amphoteric oxides from metal hydroxides react only with alkalis (hydroxides of alkali and alkaline earth metals). In this case, when the reaction is carried out in aqueous solutions, soluble complex salts are formed:

ZnO + 2NaOH + H 2 O = Na 2- sodium tetrahydroxozincate

BeO + 2NaOH + H 2 O = Na 2- sodium tetrahydroxoberyllate

Al 2 O 3 + 2NaOH + 3H 2 O = 2Na- sodium tetrahydroxyaluminate

Cr 2 O 3 + 6NaOH + 3H 2 O = 2Na 3- sodium hexahydroxochromate (III)

And when these same amphoteric oxides are fused with alkalis, salts are obtained consisting of an alkali or alkaline earth metal cation and an anion of the type MeO 2 x -, where x= 2 in the case of amphoteric oxide type Me +2 O and x= 1 for an amphoteric oxide of the form Me 2 +2 O 3:

ZnO + 2NaOH = t o=> Na 2 ZnO 2 + H 2 O

BeO + 2NaOH = t o=> Na 2 BeO 2 + H 2 O

Al 2 O 3 + 2NaOH = t o=> 2NaAlO 2 + H 2 O

Cr 2 O 3 + 2NaOH = t o=> 2NaCrO 2 + H 2 O

Fe 2 O 3 + 2NaOH = t o=> 2NaFeO 2 + H 2 O

It should be noted that salts obtained by fusing amphoteric oxides with solid alkalis can be easily obtained from solutions of the corresponding complex salts by evaporation and subsequent calcination:

Na 2 = t o=> Na 2 ZnO 2 + 2H 2 O

Na = t o=> NaAlO 2 + 2H 2 O

Interaction of oxides with medium salts

Most often, medium salts do not react with oxides.

However, you should learn the following exceptions to of this rule, which often appear on the exam.

One of these exceptions is that amphoteric oxides, as well as silicon dioxide (SiO 2), when fused with sulfites and carbonates, displace sulfur dioxide (SO 2) and carbon dioxide (CO 2) gases from the latter, respectively. For example:

Al 2 O 3 + Na 2 CO 3 = t o=> 2NaAlO 2 + CO 2

SiO 2 + K 2 SO 3 = t o=> K 2 SiO 3 + SO 2

Also, reactions of oxides with salts can conditionally include the interaction of sulfur dioxide and carbon dioxide with aqueous solutions or suspensions of the corresponding salts - sulfites and carbonates, leading to the formation of acid salts:

Na 2 CO 3 + CO 2 + H 2 O = 2NaHCO 3

CaCO 3 + CO 2 + H 2 O = Ca(HCO 3) 2

Also, sulfur dioxide, when passed through aqueous solutions or suspensions of carbonates, displaces carbon dioxide from them due to the fact that sulfurous acid is a stronger and more stable acid than carbonic acid:

K 2 CO 3 + SO 2 = K 2 SO 3 + CO 2

ORR involving oxides

Reduction of metal and non-metal oxides

Just as metals can react with solutions of salts of less active metals, displacing the latter in free form, metal oxides when heated are also able to react with more active metals.

Let us recall that the activity of metals can be compared either using the activity series of metals, or, if one or two metals are not in the activity series, by their position relative to each other in the periodic table: the lower and to the left the metal, the more active it is. It is also useful to remember that any metal from the AHM and ALP family will always be more active than a metal that is not a representative of ALM or ALP.

In particular, the aluminothermy method, used in industry to obtain such difficult-to-reduce metals as chromium and vanadium, is based on the interaction of a metal with the oxide of a less active metal:

Cr 2 O 3 + 2Al = t o=> Al 2 O 3 + 2Cr

During the aluminothermic process, a colossal amount of heat is generated, and the temperature of the reaction mixture can reach more than 2000 o C.

Also, the oxides of almost all metals located in the activity series to the right of aluminum can be reduced to free metals by hydrogen (H 2), carbon (C) and carbon monoxide (CO) when heated. For example:

Fe 2 O 3 + 3CO = t o=> 2Fe + 3CO 2

CuO+C= t o=> Cu + CO

FeO + H2 = t o=> Fe + H 2 O

It should be noted that if the metal can have several states of oxidation, if there is a lack of the reducing agent used, incomplete reduction of the oxides is also possible. For example:

Fe 2 O 3 + CO =t o=> 2FeO + CO 2

4CuO + C = t o=> 2Cu 2 O + CO 2

Oxides of active metals (alkali, alkaline earth, magnesium and aluminum) with hydrogen and carbon monoxide don't react.

However, oxides of active metals react with carbon, but differently than oxides of less active metals.

Within the framework of the Unified State Examination program, in order not to be confused, it should be assumed that as a result of the reaction of oxides of active metals (up to Al inclusive) with carbon, the formation of free alkali metal, alkali metal, Mg, and Al is impossible. In such cases, metal carbide is formed and carbon monoxide. For example:

2Al 2 O 3 + 9C = t o=> Al 4 C 3 + 6CO

CaO + 3C = t o=> CaC 2 + CO

Oxides of nonmetals can often be reduced by metals to free nonmetals. For example, when heated, oxides of carbon and silicon react with alkali, alkaline earth metals and magnesium:

CO2 + 2Mg = t o=> 2MgO + C

SiO2 + 2Mg = t o=>Si + 2MgO

With an excess of magnesium, the latter interaction can also lead to the formation magnesium silicide Mg 2 Si:

SiO2 + 4Mg = t o=> Mg 2 Si + 2 MgO

Nitrogen oxides can be reduced relatively easily even with less active metals, such as zinc or copper:

Zn + 2NO = t o=> ZnO + N 2

NO 2 + 2Cu = t o=> 2CuO + N 2

Interaction of oxides with oxygen

In order to be able to answer the question of whether any oxide reacts with oxygen (O 2) in the tasks of the real Unified State Examination, you first need to remember that oxides that can react with oxygen (from those that you may come across in the exam itself) can form only chemical elements from the list:

Found in real Unified State Exam oxides of any other chemical elements react with oxygen will not (!).

For a more visual and convenient memorization of the list of elements listed above, in my opinion, the following illustration is convenient:

All chemical elements capable of forming oxides that react with oxygen (from those encountered in the exam)

First of all, among the listed elements, nitrogen N should be considered, because the ratio of its oxides to oxygen differs markedly from the oxides of other elements in the above list.

It should be clearly remembered that nitrogen can form five oxides in total, namely:

Of all the nitrogen oxides that can react with oxygen only NO. This reaction occurs very easily when NO is mixed with both pure oxygen and air. In this case, a rapid change in the color of the gas from colorless (NO) to brown (NO 2) is observed:

2NO	+	O2	=	2NO 2
colorless				brown

In order to answer the question: does any oxide of any other of the chemical elements listed above react with oxygen (i.e. WITH,Si, P, S, Cu, Mn, Fe, Cr) — First of all, you need to remember them basic oxidation state (CO). Here they are :

Next, you need to remember the fact that of the possible oxides of the above chemical elements, only those that contain the element in the minimum oxidation state among those indicated above will react with oxygen. In this case, the oxidation state of the element increases to the nearest positive value from the possible:

element	The ratio of its oxidesto oxygen
WITH	The minimum among the main positive oxidation states of carbon is equal to +2 , and the closest positive one is +4 . Thus, only CO reacts with oxygen from the oxides C +2 O and C +4 O 2. In this case the reaction occurs: 2C +2 O + O 2 = t o=> 2C +4 O 2 CO 2 + O 2 ≠- the reaction is impossible in principle, because +4 – the highest degree of carbon oxidation.
Si	The minimum among the main positive oxidation states of silicon is +2, and the closest positive one to it is +4. Thus, only SiO reacts with oxygen from the oxides Si +2 O and Si +4 O 2. Due to some features of the oxides SiO and SiO 2, oxidation of only part of the silicon atoms in the oxide Si + 2 O is possible. as a result of its interaction with oxygen, a mixed oxide is formed containing both silicon in the +2 oxidation state and silicon in the +4 oxidation state, namely Si 2 O 3 (Si +2 O·Si +4 O 2): 4Si +2 O + O 2 = t o=> 2Si +2 ,+4 2 O 3 (Si +2 O·Si +4 O 2) SiO 2 + O 2 ≠- the reaction is impossible in principle, because +4 – the highest oxidation state of silicon.
P	The minimum among the main positive oxidation states of phosphorus is +3, and the closest positive one to it is +5. Thus, only P 2 O 3 reacts with oxygen from the oxides P +3 2 O 3 and P +5 2 O 5. In this case, the reaction of additional oxidation of phosphorus with oxygen occurs from the oxidation state +3 to the oxidation state +5: P +3 2 O 3 + O 2 = t o=> P +5 2 O 5 P +5 2 O 5 + O 2 ≠- the reaction is impossible in principle, because +5 – the highest oxidation state of phosphorus.
S	The minimum among the main positive oxidation states of sulfur is +4, and the closest positive oxidation state to it is +6. Thus, only SO 2 reacts with oxygen from the oxides S +4 O 2 and S +6 O 3 . In this case the reaction occurs: 2S +4 O 2 + O 2 = t o=> 2S +6 O 3 2S +6 O 3 + O 2 ≠- the reaction is impossible in principle, because +6 – the highest degree of sulfur oxidation.
Cu	The minimum among positive oxidation states of copper is +1, and the closest value to it is positive (and the only one) +2. Thus, only Cu 2 O reacts with oxygen from the oxides Cu +1 2 O, Cu +2 O. In this case, the reaction occurs: 2Cu +1 2 O + O 2 = t o=> 4Cu +2 O CuO + O 2 ≠- the reaction is impossible in principle, because +2 – the highest oxidation state of copper.
Cr	The minimum among the main positive oxidation states of chromium is +2, and the positive one closest to it is +3. Thus, only CrO reacts with oxygen from the oxides Cr +2 O, Cr +3 2 O 3 and Cr +6 O 3, while being oxidized by oxygen to the next (possible) positive oxidation state, i.e. +3: 4Cr +2 O + O 2 = t o=> 2Cr +3 2 O 3 Cr +3 2 O 3 + O 2 ≠- the reaction does not proceed, despite the fact that chromium oxide exists and in an oxidation state greater than +3 (Cr +6 O 3). The impossibility of this reaction occurring is due to the fact that the heating required for its hypothetical implementation greatly exceeds the decomposition temperature of CrO 3 oxide. Cr +6 O 3 + O 2 ≠ — this reaction cannot proceed in principle, because +6 is the highest oxidation state of chromium.
Mn	The minimum among the main positive oxidation states of manganese is +2, and the closest positive one is +4. Thus, of the possible oxides Mn +2 O, Mn +4 O 2, Mn +6 O 3 and Mn +7 2 O 7, only MnO reacts with oxygen, while being oxidized by oxygen to the next (possible) positive oxidation state, i.e. .e. +4: 2Mn +2 O + O 2 = t o=> 2Mn +4 O 2 while: Mn +4 O 2 + O 2 ≠ And Mn +6 O 3 + O 2 ≠- reactions do not occur, despite the fact that there is manganese oxide Mn 2 O 7 containing Mn in an oxidation state greater than +4 and +6. This is due to the fact that required for further hypothetical oxidation of Mn oxides +4 O2 and Mn +6 O 3 heating significantly exceeds the decomposition temperature of the resulting oxides MnO 3 and Mn 2 O 7. Mn +7 2 O 7 + O 2 ≠- this reaction is impossible in principle, because +7 – the highest oxidation state of manganese.
Fe	The minimum among the main positive oxidation states of iron is equal to +2 , and the closest one among the possible ones is +3 . Despite the fact that for iron there is an oxidation state of +6, the acidic oxide FeO 3, however, as well as the corresponding “iron” acid does not exist. Thus, of the iron oxides, only those oxides that contain Fe in the +2 oxidation state can react with oxygen. It's either Fe oxide +2 O, or mixed iron oxide Fe +2 ,+3 3 O 4 (iron scale): 4Fe +2 O + O 2 = t o=> 2Fe +3 2 O 3 or 6Fe +2 O + O 2 = t o=> 2Fe +2,+3 3 O 4 mixed oxide Fe +2,+3 3 O 4 can be oxidized to Fe +3 2 O 3: 4Fe +2,+3 3 O 4 + O 2 = t o=> 6Fe +3 2 O 3 Fe +3 2 O 3 + O 2 ≠ - this reaction is impossible in principle, because There are no oxides containing iron in an oxidation state higher than +3.

Oxides- these are complex inorganic compounds consisting of two elements, one of which is oxygen (in oxidation state -2).

For example, Na 2 O, B 2 O 3, Cl 2 O 7 are classified as oxides. All of these substances contain oxygen and one more element. The substances Na 2 O 2 , H 2 SO 4 , and HCl are not oxides: in the first, the oxidation state of oxygen is -1, in the second there are not two, but three elements, and the third does not contain oxygen at all.

If you don't understand the meaning of the term oxidation number, that's okay. First, you can refer to the corresponding article on this site. Secondly, even without understanding this term, you can continue reading. You can temporarily forget about mentioning the oxidation state.

Oxides of almost all currently known elements have been obtained, except for some noble gases and “exotic” transuranium elements. Moreover, many elements form several oxides (for nitrogen, for example, six are known).

Nomenclature of oxides

We must learn to name oxides. It's very simple.

Example 1. Name the following compounds: Li 2 O, Al 2 O 3, N 2 O 5, N 2 O 3.

Li 2 O - lithium oxide,
Al 2 O 3 - aluminum oxide,
N 2 O 5 - nitric oxide (V),
N 2 O 3 - nitric oxide (III).

pay attention to important point: If the valence of an element is constant, we do NOT mention it in the name of the oxide. If the valence changes, be sure to indicate it in parentheses! Lithium and aluminum have constant valency, nitrogen has a variable valency; It is for this reason that the names of nitrogen oxides are supplemented with Roman numerals symbolizing valency.

Exercise 1. Name the oxides: Na 2 O, P 2 O 3, BaO, V 2 O 5, Fe 2 O 3, GeO 2, Rb 2 O. Do not forget that there are elements with both constant and variable valency.

Another important point: it is more correct to call the substance F 2 O not “fluorine oxide”, but “oxygen fluoride”!

Physical properties of oxides

Physical properties are very diverse. This is due, in particular, to the fact that oxides can exhibit different types chemical bond. Melting and boiling points vary widely. At normal conditions oxides can be in the solid state (CaO, Fe 2 O 3, SiO 2, B 2 O 3), liquid state (N 2 O 3, H 2 O), in the form of gases (N 2 O, SO 2, NO, CO ).

Various colors: MgO and Na 2 O white, CuO - black, N 2 O 3 - blue, CrO 3 - red, etc.

Melts of oxides with ionic type connections are good electricity, covalent oxides generally have low electrical conductivity.

Oxides classification

All oxides existing in nature can be divided into 4 classes: basic, acidic, amphoteric and non-salt-forming. Sometimes the first three classes are combined into the group of salt-forming oxides, but for us this is not important now. The chemical properties of oxides from different classes differ greatly, so the issue of classification is very important for further study of this topic!

Let's start with non-salt-forming oxides. They need to be remembered: NO, SiO, CO, N 2 O. Just learn these four formulas!

To move forward we must remember that in nature there are two types simple substances- metals and nonmetals (sometimes another group of semimetals or metalloids is distinguished). If you have a clear understanding of which elements are metals, continue reading this article. If you have the slightest doubt, refer to the material "Metals and non-metals" on that website.

So, let me tell you that all amphoteric oxides are metal oxides, but not all metal oxides are amphoteric. I will list the most important of them: BeO, ZnO, Al 2 O 3, Cr 2 O 3, SnO. The list is not complete, but you should definitely remember the listed formulas! In most amphoteric oxides, the metal exhibits an oxidation state of +2 or +3 (but there are exceptions).

In the next part of the article we will continue to talk about classification; Let's discuss acidic and basic oxides.

Oxides are complex substances consisting of two elements, one of which is oxygen. Oxides can be salt-forming or non-salt-forming: one type of salt-forming oxides is basic oxides. How do they differ from other species, and what are their chemical properties?

Salt-forming oxides are divided into basic, acidic and amphoteric oxides. If basic oxides correspond to bases, then acidic oxides correspond to acids, and amphoteric oxides correspond to amphoteric formations. Amphoteric oxides are those compounds that, depending on conditions, can exhibit either basic or acidic properties.

Rice. 1. Classification of oxides.

The physical properties of oxides are very diverse. They can be either gases (CO 2), solids (Fe 2 O 3) or liquid substances (H 2 O).

However, most basic oxides are solids of various colors.

oxides in which elements exhibit their highest activity are called higher oxides. The order of increase in the acidic properties of higher oxides of the corresponding elements in periods from left to right is explained by a gradual increase in the positive charge of the ions of these elements.

Chemical properties of basic oxides

Basic oxides are the oxides to which bases correspond. For example, the basic oxides K 2 O, CaO correspond to the bases KOH, Ca(OH) 2.

Rice. 2. Basic oxides and their corresponding reasons.

Basic oxides are formed by typical metals, as well as metals of variable valence in the lowest oxidation state (for example, CaO, FeO), react with acids and acid oxides, forming salts:

CaO (basic oxide) + CO 2 (acid oxide) = CaCO 3 (salt)

FeO (basic oxide)+H 2 SO 4 (acid)=FeSO 4 (salt)+2H 2 O (water)

Basic oxides also react with amphoteric oxides, resulting in the formation of a salt, for example:

Only oxides of alkali and alkaline earth metals react with water:

BaO (basic oxide)+H 2 O (water)=Ba(OH) 2 (alkali earth metal base)

Many basic oxides tend to be reduced to substances consisting of atoms of one chemical element:

3CuO+2NH 3 =3Cu+3H 2 O+N 2

When heated, only oxides of mercury and noble metals decompose:

Rice. 3. Mercury oxide.

List of main oxides:

Oxide name	Chemical formula	Properties
Calcium oxide	CaO	quicklime, white crystalline substance
Magnesium oxide	MgO	white substance, slightly soluble in water
Barium oxide	BaO	colorless crystals with a cubic lattice
Copper oxide II	CuO	black substance practically insoluble in water
	HgO	red or yellow-orange solid
Potassium oxide	K2O	colorless or pale yellow substance
Sodium oxide	Na2O	substance consisting of colorless crystals
Lithium oxide	Li2O	a substance consisting of colorless crystals that have a cubic lattice structure

Non-salt-forming (indifferent, indifferent) oxides CO, SiO, N 2 0, NO.

Salt-forming oxides:

Basic. Oxides whose hydrates are bases. Metal oxides with oxidation states +1 and +2 (less often +3). Examples: Na 2 O - sodium oxide, CaO - calcium oxide, CuO - copper (II) oxide, CoO - cobalt (II) oxide, Bi 2 O 3 - bismuth (III) oxide, Mn 2 O 3 - manganese (III) oxide ).

Amphoteric. Oxides whose hydrates are amphoteric hydroxides. Metal oxides with oxidation states +3 and +4 (less often +2). Examples: Al 2 O 3 - aluminum oxide, Cr 2 O 3 - chromium (III) oxide, SnO 2 - tin (IV) oxide, MnO 2 - manganese (IV) oxide, ZnO - zinc oxide, BeO - beryllium oxide.

Acidic. Oxides whose hydrates are oxygen-containing acids. Non-metal oxides. Examples: P 2 O 3 - phosphorus (III) oxide, CO 2 - carbon oxide (IV), N 2 O 5 - nitrogen oxide (V), SO 3 - sulfur oxide (VI), Cl 2 O 7 - chlorine oxide ( VII). Metal oxides with oxidation states +5, +6 and +7. Examples: Sb 2 O 5 - antimony (V) oxide. CrOz - chromium (VI) oxide, MnOz - manganese (VI) oxide, Mn 2 O 7 - manganese (VII) oxide.

Change in the nature of oxides with increasing oxidation state of the metal

Physical properties

Oxides are solid, liquid and gaseous, of different colors. For example: copper (II) oxide CuO is black, calcium oxide CaO is white - solids. Sulfur oxide (VI) SO 3 is a colorless volatile liquid, and carbon monoxide (IV) CO 2 is a colorless gas under ordinary conditions.

State of aggregation

CaO, CuO, Li 2 O and other basic oxides; ZnO, Al 2 O 3, Cr 2 O 3 and other amphoteric oxides; SiO 2, P 2 O 5, CrO 3 and other acid oxides.

SO 3, Cl 2 O 7, Mn 2 O 7, etc.

Gaseous:

CO 2, SO 2, N 2 O, NO, NO 2, etc.

Solubility in water

Soluble:

a) basic oxides of alkali and alkaline earth metals;

b) almost all acid oxides (exception: SiO 2).

Insoluble:

a) all other basic oxides;

b) all amphoteric oxides

Chemical properties

1. Acid-base properties

Common properties of basic, acidic and amphoteric oxides are acid-base interactions, which are illustrated by the following diagram:

(only for oxides of alkali and alkaline earth metals) (except SiO 2).

Amphoteric oxides, having the properties of both basic and acidic oxides, interact with strong acids and alkalis:

2. Redox properties

If an element has a variable oxidation state (s.o.), then its oxides with low s. O. can exhibit reducing properties, and oxides with high c. O. - oxidative.

Examples of reactions in which oxides act as reducing agents:

Oxidation of oxides with low c. O. to oxides with high c. O. elements.

2C +2 O + O 2 = 2C +4 O 2

2S +4 O 2 + O 2 = 2S +6 O 3

2N +2 O + O 2 = 2N +4 O 2

Carbon (II) monoxide reduces metals from their oxides and hydrogen from water.

C +2 O + FeO = Fe + 2C +4 O 2

C +2 O + H 2 O = H 2 + 2C +4 O 2

Examples of reactions in which oxides act as oxidizing agents:

Reduction of oxides with high o. elements to oxides with low c. O. or to simple substances.

C +4 O 2 + C = 2C +2 O

2S +6 O 3 + H 2 S = 4S +4 O 2 + H 2 O

C +4 O 2 + Mg = C 0 + 2MgO

Cr +3 2 O 3 + 2Al = 2Cr 0 + 2Al 2 O 3

Cu +2 O + H 2 = Cu 0 + H 2 O

The use of oxides of low-active metals for the oxidation of organic substances.

Some oxides in which the element has an intermediate c. o., capable of disproportionation;

For example:

2NO 2 + 2NaOH = NaNO 2 + NaNO 3 + H 2 O

Methods of obtaining

1. Interaction of simple substances - metals and non-metals - with oxygen:

4Li + O 2 = 2Li 2 O;

2Cu + O 2 = 2CuO;

4P + 5O 2 = 2P 2 O 5

2. Dehydration of insoluble bases, amphoteric hydroxides and some acids:

Cu(OH) 2 = CuO + H 2 O

2Al(OH) 3 = Al 2 O 3 + 3H 2 O

H 2 SO 3 = SO 2 + H 2 O

H 2 SiO 3 = SiO 2 + H 2 O

3. Decomposition of some salts:

2Cu(NO 3) 2 = 2CuO + 4NO 2 + O 2

CaCO 3 = CaO + CO 2

(CuOH) 2 CO 3 = 2CuO + CO 2 + H 2 O

4. Oxidation complex substances oxygen:

CH 4 + 2O 2 = CO 2 + H 2 O

4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2

4NH 3 + 5O 2 = 4NO + 6H 2 O

5. Reduction of oxidizing acids with metals and non-metals:

Cu + H 2 SO 4 (conc) = CuSO 4 + SO 2 + 2H 2 O

10HNO 3 (conc) + 4Ca = 4Ca(NO 3) 2 + N 2 O + 5H 2 O

2HNO 3 (diluted) + S = H 2 SO 4 + 2NO

6. Interconversions of oxides during redox reactions (see redox properties of oxides).

1. Metal + Non-metal. Inert gases do not enter into this interaction. The higher the electronegativity of a nonmetal, the more a large number metals it will react. For example, fluorine reacts with all metals, and hydrogen only with active ones. The further to the left a metal is in the metal activity series, the more nonmetals it can react with. For example, gold reacts only with fluorine, lithium - with all non-metals.

2. Non-metal + non-metal. In this case, a more electronegative nonmetal acts as an oxidizing agent, and a less electronegative nonmetal acts as a reducing agent. Nonmetals with similar electronegativity interact poorly with each other, for example, the interaction of phosphorus with hydrogen and silicon with hydrogen is practically impossible, since the equilibrium of these reactions is shifted towards the formation of simple substances. Helium, neon and argon do not react with non-metals; other inert gases can react with fluorine under harsh conditions.
Oxygen does not interact with chlorine, bromine and iodine. Oxygen can react with fluorine at low temperatures.

3. Metal + acid oxide. The metal reduces the nonmetal from the oxide. The excess metal can then react with the resulting nonmetal. For example:

2 Mg + SiO 2 = 2 MgO + Si (with magnesium deficiency)

2 Mg + SiO 2 = 2 MgO + Mg 2 Si (with excess magnesium)

4. Metal + acid. Metals located in the voltage series to the left of hydrogen react with acids to release hydrogen.

The exception is oxidizing acids (concentrated sulfur and any nitric acid), which can react with metals that are in the voltage series to the right of hydrogen; in the reactions, hydrogen is not released, but water and the acid reduction product are obtained.

It is necessary to pay attention to the fact that when a metal reacts with an excess of a polybasic acid, an acid salt can be obtained: Mg +2 H 3 PO 4 = Mg (H 2 PO 4 ) 2 + H 2 .

If the product of the interaction between an acid and a metal is an insoluble salt, then the metal is passivated, since the surface of the metal is protected by the insoluble salt from the action of the acid. For example, the effect of dilute sulfuric acid on lead, barium or calcium.

5. Metal + salt. In solution This reaction involves metals that are in the voltage series to the right of magnesium, including magnesium itself, but to the left of the metal salt. If the metal is more active than magnesium, then it reacts not with salt, but with water to form an alkali, which subsequently reacts with salt. In this case, the original salt and the resulting salt must be soluble. The insoluble product passivates the metal.

However, there are exceptions to this rule:

2FeCl 3 + Cu = CuCl 2 + 2FeCl 2;

2FeCl 3 + Fe = 3FeCl 2. Since iron has an intermediate oxidation state, its salt in the highest oxidation state is easily reduced to a salt in the intermediate oxidation state, oxidizing even less active metals.

In meltsa number of metal stresses are not effective. Determining whether a reaction between a salt and a metal is possible can only be done using thermodynamic calculations. For example, sodium can displace potassium from a potassium chloride melt, since potassium is more volatile: Na + KCl = NaCl + K (this reaction is determined by the entropy factor). On the other hand, aluminum was obtained by displacement from sodium chloride: 3 Na + AlCl 3 = 3 NaCl + Al . This process is exothermic and is determined by the enthalpy factor.

It is possible that the salt decomposes when heated, and the products of its decomposition can react with the metal, for example, aluminum nitrate and iron. Aluminum nitrate decomposes when heated into aluminum oxide, nitric oxide ( IV ) and oxygen, oxygen and nitric oxide will oxidize iron:

10Fe + 2Al(NO 3) 3 = 5Fe 2 O 3 + Al 2 O 3 + 3N 2

6. Metal + basic oxide. Just as in molten salts, the possibility of these reactions is determined thermodynamically. Aluminum, magnesium and sodium are often used as reducing agents. For example: 8 Al + 3 Fe 3 O 4 = 4 Al 2 O 3 + 9 Fe exothermic reaction, enthalpy factor);2 Al + 3 Rb 2 O = 6 Rb + Al 2 O 3 (volatile rubidium, enthalpy factor).

8. Non-metal + base. As a rule, the reaction occurs between a non-metal and an alkali. Not all non-metals can react with alkalis: you need to remember that halogens (in different ways depending on temperature), sulfur (when heated), silicon, phosphorus enter into this interaction.

KOH + Cl 2 = KClO + KCl + H 2 O (in the cold)

6 KOH + 3 Cl 2 = KClO 3 + 5 KCl + 3 H 2 O (in hot solution)

6KOH + 3S = K 2 SO 3 + 2K 2 S + 3H 2 O

2KOH + Si + H 2 O = K 2 SiO 3 + 2H 2

3KOH + 4P + 3H 2 O = PH 3 + 3KPH 2 O 2

1) non-metal – reducing agent (hydrogen, carbon):

CO 2 + C = 2CO;

2NO 2 + 4H 2 = 4H 2 O + N 2;

SiO 2 + C = CO 2 + Si. If the resulting non-metal can react with the metal used as a reducing agent, then the reaction will go further (with an excess of carbon) SiO 2 + 2 C = CO 2 + Si C

2) non-metal – oxidizing agent (oxygen, ozone, halogens):

2С O + O 2 = 2СО 2.

C O + Cl 2 = CO Cl 2.

2 NO + O 2 = 2 N O 2.

10. Acidic oxide + basic oxide . The reaction occurs if the resulting salt exists in principle. For example, aluminum oxide can react with sulfuric anhydride to form aluminum sulfate, but cannot react with carbon dioxide because the corresponding salt does not exist.

11. Water + basic oxide . The reaction is possible if an alkali is formed, that is, a soluble base (or slightly soluble, in the case of calcium). If the base is insoluble or slightly soluble, then the reverse reaction of decomposition of the base into oxide and water occurs.

12. Basic oxide + acid . The reaction is possible if the resulting salt exists. If the resulting salt is insoluble, the reaction may be passivated due to the blocking of acid access to the oxide surface. In case of excess polybasic acid, the formation of an acid salt is possible.

13. Acid oxide + base. Typically, the reaction occurs between an alkali and an acidic oxide. If an acid oxide corresponds to a polybasic acid, an acid salt can be obtained: CO 2 + KOH = KHCO 3.

Acidic oxides, corresponding to strong acids, can also react with insoluble bases.

Sometimes oxides corresponding to weak acids react with insoluble bases, which can result in an average or basic salt (usually less soluble substance): 2 Mg (OH) 2 + CO 2 = (MgOH) 2 CO 3 + H 2 O.

14. Acid oxide + salt. The reaction can take place in a melt or in solution. In the melt, the less volatile oxide displaces the more volatile oxide from the salt. In solution, the oxide corresponding to the stronger acid displaces the oxide corresponding to the weaker acid. For example, Na 2 CO 3 + SiO 2 = Na 2 SiO 3 + CO 2 , in the forward direction, this reaction occurs in the melt, carbon dioxide is more volatile than silicon oxide; in the opposite direction, the reaction occurs in solution, carbonic acid is stronger than silicic acid, and silicon oxide precipitates.

It is possible to combine an acidic oxide with its own salt, for example, dichromate can be obtained from chromate, and disulfate from sulfate, and disulfite from sulfite:

Na 2 SO 3 + SO 2 = Na 2 S 2 O 5

To do this, you need to take a crystalline salt and pure oxide, or a saturated salt solution and an excess of acidic oxide.

In solution, salts can react with their own acid oxides to form acid salts: Na 2 SO 3 + H 2 O + SO 2 = 2 NaHSO 3

15. Water + acid oxide . The reaction is possible if a soluble or slightly soluble acid is formed. If the acid is insoluble or slightly soluble, then a reverse reaction occurs, the decomposition of the acid into oxide and water. For example, sulfuric acid is characterized by a reaction of production from oxide and water, the decomposition reaction practically does not occur, silicic acid cannot be obtained from water and oxide, but it easily decomposes into these components, but carbon and sulfurous acid can participate in both direct and reverse reactions.

16. Base + acid. A reaction occurs if at least one of the reactants is soluble. Depending on the ratio of the reagents, medium, acidic and basic salts can be obtained.

17. Base + salt. The reaction occurs if both starting substances are soluble, and at least one non-electrolyte or weak electrolyte (precipitate, gas, water) is obtained as a product.

18. Salt + acid. As a rule, a reaction occurs if both starting substances are soluble, and at least one non-electrolyte or weak electrolyte (precipitate, gas, water) is obtained as a product.

A strong acid can react with insoluble salts of weak acids (carbonates, sulfides, sulfites, nitrites), and a gaseous product is released.

Reactions between concentrated acids and crystalline salts are possible if this results in a more volatile acid: for example, hydrogen chloride can be obtained by the action of concentrated sulfuric acid on crystalline sodium chloride, hydrogen bromide and hydrogen iodide - by the action of orthophosphoric acid on the corresponding salts. You can act with an acid on your own salt to produce an acidic salt, for example: BaSO 4 + H 2 SO 4 = Ba (HSO 4 ) 2 .

19. Salt + salt.As a rule, a reaction occurs if both starting substances are soluble, and at least one non-electrolyte or weak electrolyte is obtained as a product.

1) salt does not exist because irreversibly hydrolyzes . These are the majority of carbonates, sulfites, sulfides, silicates of trivalent metals, as well as some salts of divalent metals and ammonium. Trivalent metal salts are hydrolyzed to the corresponding base and acid, and divalent metal salts are hydrolyzed to less soluble basic salts.

Let's look at examples:

2 FeCl 3 + 3 Na 2 CO 3 = Fe 2 ( CO 3 ) 3 + 6 NaCl (1)

Fe 2 (CO 3) 3+ 6H 2 O = 2Fe(OH) 3 + 3 H2CO3

H 2 CO 3 decomposes into water and carbon dioxide, the water in the left and right parts is reduced and the result is: Fe 2 ( CO 3 ) 3 + 3 H 2 O = 2 Fe (OH) 3 + 3 CO 2 (2)

If we now combine (1) and (2) equations and reduce iron carbonate, we obtain a total equation reflecting the interaction of ferric chloride ( III ) and sodium carbonate: 2 FeCl 3 + 3 Na 2 CO 3 + 3 H 2 O = 2 Fe (OH) 3 + 3 CO 2 + 6 NaCl

CuSO 4 + Na 2 CO 3 = CuCO 3 + Na 2 SO 4 (1)

The underlined salt does not exist due to irreversible hydrolysis:

2CuCO3+ H 2 O = (CuOH) 2 CO 3 +CO 2 (2)

If we now combine (1) and (2) equations and reduce copper carbonate, we obtain a total equation reflecting the interaction of sulfate ( II ) and sodium carbonate:

2CuSO 4 + 2Na 2 CO 3 + H 2 O = (CuOH) 2 CO 3 + CO 2 + 2Na 2 SO 4