Before we start talking about the chemical properties of oxides, we 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, you first need to understand whether the oxide of the metal or non-metal is 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) Non-metal oxidation state +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: acidic oxide Exception: Cl 2 O is an acid 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 metals 3) Metal oxidation state +5, +6, +7 Conclusion: acidic oxide

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

• To active basic oxides Let us refer oxides of alkali and alkaline earth metals (all elements of groups IA and IIA, except for hydrogen H, beryllium Be and magnesium Mg). For example, Na 2 O, CaO, Rb 2 O, SrO, etc.
• To inactive basic oxides we will assign all the main oxides that were 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 those reactions that do not enter into low-active ones.
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 chemical reactions can take place. However, it often takes a direct part in various transformations, in particular, some groups of oxides react with it.

What oxides react with water?

Of all oxides with water react only:
1) all active basic oxides (oxides of alkaline metals and alkaline earth metals);
2) all acidic oxides, except for silicon dioxide (SiO 2);

those. From the foregoing, it follows that with water exactly do not 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 of the test part of the exam.

Now let's see how, after all, certain oxides react with water, i.e. learn how to write the corresponding reaction equations.

Active basic oxides, reacting with water, form their corresponding hydroxides. Recall that the corresponding metal oxide is the 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, the corresponding hydroxides K + 1 OH and Ba + 2 (OH) 2 are formed:

K 2 O + H 2 O \u003d 2KOH– potassium hydroxide

BaO + H 2 O \u003d Ba (OH) 2– barium hydroxide

All hydroxides corresponding to active basic oxides (oxides of alkali metals and alkali earth metals) are alkalis. Alkalis are all water-soluble metal hydroxides, 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 acid oxides, they correspond not to basic, but to acidic hydroxides, more often called oxygenated acids. Recall that the corresponding acid 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 recall the main sulfur-containing acids studied in the school curriculum. These are hydrogen sulfide H 2 S, sulfurous H 2 SO 3 and sulfuric H 2 SO 4 acids. Hydrosulfide acid H 2 S, as you can easily 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 +6 oxidation state, as in oxide SO 3, contains only sulfuric acid H 2 SO 4. Therefore, it is she who will be formed in the reaction of SO 3 with water:

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

Similarly, 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, as in N 2 O 5 , equal to +5, and in nitrogenous - +3:

N +5 2 O 5 + H 2 O \u003d 2HN +5 O 3

Interaction of oxides with each other

First of all, it is necessary to clearly understand the fact that among salt-forming oxides (acidic, basic, amphoteric), reactions between oxides of the same class almost never occur, 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 between oxides belonging to different types is almost always possible, i.e. almost always flow reactions between:

1) basic oxide and acid oxide;

2) amphoteric oxide and acid oxide;

3) amphoteric oxide and basic oxide.

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

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

As a result of 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 original Me x O y) and the acid residue of the acid corresponding to the acid oxide.

For 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 acid oxide (P 2 O 5). In the second - amphoteric oxide (Al 2 O 3) and acid 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 acid 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 acid residue PO 4 3-, since the oxide P +5 2 O 5 corresponds to acid H 3 P +5 O 4 . Those. As a result of this interaction, sodium phosphate is formed:

3Na 2 O + P 2 O 5 \u003d 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 acid residue SO 4 2-, since the oxide S +6 O 3 corresponds to acid H 2 S +6 O 4 . Thus, as a result of this reaction, aluminum sulfate is obtained:

Al 2 O 3 + 3SO 3 \u003d 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 the acidic one. As a result of this interaction, a salt of a specific composition is formed, consisting of a metal cation that forms the initial basic oxide and an "acid residue" / anion, which includes the metal from the amphoteric oxide. The formula of such an "acid residue" / anion in general form 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 a general formula of the form Me + 2 O (ZnO, BeO, PbO) and x = 1 - for amphoteric oxides with the 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 as an example the interaction equations

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 an "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 containing a metal cation from the basic oxide is formed, i.e. Ba 2+ (from BaO) and "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 proceeds:

Me x O y + acid oxide,

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

However, two "finicky" acidic oxides should be remembered - carbon dioxide (CO 2) and sulfur dioxide (SO 2). Their “fastidiousness” lies in the fact that, despite the 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 earth metal). So, for example, Na 2 O and BaO, being active basic oxides, can react with them:

CO 2 + Na 2 O \u003d Na 2 CO 3

SO 2 + BaO = BaSO 3

While CuO and Al 2 O 3 oxides, 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. This forms salts and water:

FeO + H 2 SO 4 \u003d FeSO 4 + H 2 O

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

When does acid oxide react with acid?

When solving the part of the exam with answer options, you should conditionally assume that acid oxides do not react with either acid oxides or acids, except for 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 an excess of HF, the reaction equation has the form:

SiO 2 + 6HF \u003d H 2 + 2H 2 O,

and in case of lack of HF:

SiO 2 + 4HF \u003d SiF 4 + 2H 2 O

2) SO 2, being an acid oxide, easily reacts with hydrosulfide acid H 2 S according to the type co-proportionation:

S +4 O 2 + 2H 2 S -2 \u003d 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 of any concentration. In this case, the oxidation state of phosphorus increases from +3 to +5:

P2O3	+	2H2SO4	+	H2O	=t o=>	2SO2	+	2H3PO4
		(conc.)

3 P2O3	+	4HNO 3	+	7 H2O	=t o=>	4NO	+	6 H3PO4
		(razb.)

2HNO 3	+	3SO2	+	2H2O	=t o=>	3H2SO4	+	2NO
(razb.)

Interaction of oxides with metal hydroxides

Acid oxides react with metal hydroxides, both basic and amphoteric. In this case, a salt is formed, consisting of a metal cation (from the initial metal hydroxide) and an acid residue of the acid corresponding to the acid oxide.

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

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

CO 2 + 2NaOH \u003d Na 2 CO 3 + H 2 O

CO 2 + NaOH = NaHCO 3

P 2 O 5 + 6KOH \u003d 2K 3 PO 4 + 3H 2 O

P 2 O 5 + 4KOH \u003d 2K 2 HPO 4 + H 2 O

P 2 O 5 + 2KOH + H 2 O \u003d 2KH 2 PO 4

The "finicky" oxides CO 2 and SO 2, whose activity, as already mentioned, is not enough for their reaction with low-activity basic and amphoteric oxides, nevertheless, react with most of the metal hydroxides corresponding to them. More precisely, carbon dioxide and sulfur dioxide interact with insoluble hydroxides in the form of their suspension in water. In this case, only basic about obvious salts, called hydroxocarbonates and hydroxosulfites, and the formation of medium (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, with metal hydroxides in the +3 oxidation state, for example, such as Al (OH) 3, Cr (OH) 3, etc., carbon dioxide and sulfur dioxide do not react at all.

It should also be noted the special inertness of silicon dioxide (SiO 2), which is most often found in nature in the form of ordinary sand. This oxide is acidic, however, among metal hydroxides, it is able to react 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 carrying out the reaction in aqueous solutions, soluble complex salts are formed:

ZnO + 2NaOH + H 2 O \u003d Na 2- sodium tetrahydroxozincate

BeO + 2NaOH + H 2 O \u003d Na 2- sodium tetrahydroxoberyllate

Al 2 O 3 + 2NaOH + 3H 2 O \u003d 2Na- sodium tetrahydroxoaluminate

Cr 2 O 3 + 6NaOH + 3H 2 O \u003d 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 MeO 2 x - type, 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 \u003d t o=> 2NaAlO 2 + H 2 O

Cr 2 O 3 + 2NaOH \u003d t o=> 2NaCrO 2 + H 2 O

Fe 2 O 3 + 2NaOH \u003d 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 their 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 this rule, which are often found 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 sulfurous (SO 2) and carbon dioxide (CO 2) gases from the latter, respectively. For example:

Al 2 O 3 + Na 2 CO 3 \u003d t o=> 2NaAlO 2 + CO 2

SiO 2 + K 2 SO 3 \u003d t o=> K 2 SiO 3 + SO 2

Also, the reactions of oxides with salts can be conditionally attributed to 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 \u003d 2NaHCO 3

CaCO 3 + CO 2 + H 2 O \u003d 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 \u003d K 2 SO 3 + CO 2

OVR involving oxides

Recovery of oxides of metals and non-metals

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

Recall that you can compare the activity of metals either using the activity series of metals, or, if one or two metals are not in the activity series at once, by their position relative to each other in the periodic table: the lower and to the left of the metal, the more active it is. It is also useful to remember that any metal from the SM and SHM family will always be more active than a metal that is not a representative of SHM or SHM.

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

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

During the process of aluminothermy, an enormous amount of heat is generated, and the temperature of the reaction mixture can reach more than 2000 o C.

Also, oxides of almost all metals that are in the activity series to the right of aluminum can be reduced to free metals with 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 + H 2 \u003d t o=> Fe + H 2 O

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

Fe 2 O 3 + CO =to=> 2FeO + CO 2

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

Oxides of active metals (alkaline, alkaline earth, magnesium and aluminum) with hydrogen and carbon monoxide do not react.

However, oxides of active metals react with carbon, but in a different way than oxides of less active metals.

Within the framework of the USE program, in order not to be confused, it should be considered that as a result of the reaction of active metal oxides (up to Al inclusive) with carbon, the formation of free alkaline metals, alkaline earth metals, Mg, and also Al is impossible. In such cases, the formation of metal carbide and carbon monoxide occurs. For example:

2Al 2 O 3 + 9C \u003d t o=> Al 4 C 3 + 6CO

CaO + 3C = t o=> CaC2 + CO

Non-metal oxides can often be reduced by metals to free non-metals. So, for example, oxides of carbon and silicon, when heated, react with alkali, alkaline earth metals and magnesium:

CO 2 + 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 Mg2Si:

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

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 exam, you first need to remember that oxides that can react with oxygen (of those that you can come across on the exam itself) can form only chemical elements from the list:

Oxides of any other chemical elements encountered in the real USE react with oxygen will not (!).

For a more visual convenient memorization of the above list of elements, 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 the rest of the elements in the above list.

It should be clearly remembered that in total nitrogen is capable of forming five oxides, namely:

Of all nitrogen oxides, oxygen can react only NO. This reaction proceeds 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 above chemical elements react with oxygen (i.e. WITH,Si, P, S, Cu, Mn, Fe, Cr) — First of all, you need to remember them main 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, among the above, oxidation states will react with oxygen. In this case, the oxidation state of the element rises to the nearest positive value possible:

element	The ratio of its oxidesto oxygen
With	The minimum among the main positive oxidation states of carbon is +2 , and the closest positive to it is +4 . Thus, only CO reacts with oxygen from the oxides C +2 O and C +4 O 2. In this case, the reaction proceeds: 2C +2 O + O 2 = t o=> 2C+4O2 CO 2 + O 2 ≠- the reaction is impossible in principle, because +4 is the highest oxidation state of carbon.
Si	The minimum among the main positive oxidation states of silicon is +2, and the closest positive 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, only a part of the silicon atoms in the oxide Si + 2 O can be oxidized. 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 \u003d 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 is the highest oxidation state of silicon.
P	The minimum among the main positive oxidation states of phosphorus is +3, and the closest positive to it is +5. Thus, only P 2 O 3 reacts with oxygen from oxides P +3 2 O 3 and P +5 2 O 5 . In this case, the reaction of additional oxidation of phosphorus with oxygen proceeds 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 is the highest oxidation state of phosphorus.
S	The minimum among the main positive oxidation states of sulfur is +4, and the closest positive to it in value is +6. Thus, only SO 2 reacts with oxygen from oxides S +4 O 2 , S +6 O 3 . In this case, the reaction proceeds: 2S +4 O 2 + O 2 \u003d t o=> 2S +6 O 3 2S +6 O 3 + O 2 ≠- the reaction is impossible in principle, because +6 is the highest oxidation state of sulfur.
Cu	The minimum among the positive oxidation states of copper is +1, and the closest to it in value is the positive (and only) +2. Thus, only Cu 2 O reacts with oxygen from oxides Cu +1 2 O, Cu +2 O. In this case, the reaction proceeds: 2Cu +1 2 O + O 2 = t o=> 4Cu+2O CuO + O 2 ≠- the reaction is impossible in principle, because +2 is the highest oxidation state of copper.
Cr	The minimum among the main positive oxidation states of chromium is +2, and the closest positive to it in value 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 (out of possible) positive oxidation state, i.e. +3: 4Cr +2 O + O 2 \u003d 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 to it 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 neighboring (out of possible) positive oxidation state, t .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 proceed, despite the fact that there is manganese oxide Mn 2 O 7 containing Mn in a higher oxidation state than +4 and +6. This is due to the fact that the 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 is the highest oxidation state of manganese.
Fe	The minimum among the main positive oxidation states of iron is +2 , and the closest to it among the possible - +3 . Despite the fact that for iron there is an oxidation state of +6, the acid 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 \u003d t o=> 2Fe +3 2 O 3 or 6Fe +2 O + O 2 \u003d t o=> 2Fe +2,+3 3 O 4 mixed Fe oxide +2,+3 3 O 4 can be further oxidized to Fe +3 2O3: 4Fe +2 ,+3 3 O 4 + O 2 = t o=> 6Fe +3 2 O 3 Fe +3 2 O 3 + O 2 ≠ - the course of this reaction is impossible in principle, because oxides containing iron in an oxidation state higher than +3 do not exist.

Oxides are binary compounds of an element with oxygen in the oxidation state (-2). Oxides are characteristic compounds for chemical elements. It is no coincidence that D.I. Mendeleev, when compiling the periodic table, was guided by the stoichiometry of the higher oxide and combined elements with the same formula of the higher oxide into one group. The highest oxide is the oxide in which the element has attached the maximum possible number of oxygen atoms for it. In the higher oxide, the element is in its maximum (highest) oxidation state. Thus, the higher oxides of group VI elements, both non-metals S, Se, Te, and metals Cr, Mo, W, are described by the same formula EO 3 . All elements of the group show the greatest similarity precisely in the highest degree of oxidation. So, for example, all higher oxides of elements of group VI are acidic.

oxides- these are the most common compounds in metallurgical technologies.

Many metals are found in the earth's crust in the form of oxides.. From natural oxides, important metals such as Fe, Mn, Sn, Cr.

The table shows examples of natural oxides used to obtain metals.

Me	Oxide	Mineral
Fe	Fe 2 O 3 and Fe 3 O 4	hematite and magnetite
Mn	MnO2	pyrolusite
Cr	FeO . Cr2O3	chromite
Ti	TiO2 and FeO . TiO2	Rutile and ilmenite
sn	SnO 2	Cassiterite

Oxides are target compounds in a number of metallurgical technologies. Natural compounds are first converted into oxides, from which the metal is then reduced. For example, natural sulfides Zn, Ni, Co, Pb, Mo are burned, turning into oxides.

2ZnS + 3O 2 = 2 ZnO + 2SO 2

Natural hydroxides and carbonates undergo thermal decomposition leading to the formation of an oxide.

2MeOOH \u003d Me 2 O 3 + H 2 O

MeCO 3 \u003d MeO + CO 2

In addition, since metals, being in the environment, are oxidized by atmospheric oxygen, and at high temperatures, characteristic of many metallurgical industries, the oxidation of metals is enhanced, knowledge of the properties of the obtained oxides is required.

The above reasons explain why oxides are given special attention in discussions of metal chemistry.

Among the chemical elements of metals - 85, and many metals have more than one oxide, so the class of oxides includes a huge number of compounds, and this multiplicity makes reviewing their properties a difficult task. However, will try to identify:

• general properties inherent in all metal oxides,
• patterns in changes in their properties,
• reveal the chemical properties of the oxides most widely used in metallurgy,
• Let us present some of the important physical characteristics of metal oxides.

oxides metals differ in the stoichiometric ratio of metal and oxygen atoms. These stoichiometric ratios determine the degree of oxidation of the metal in the oxide.

The table lists the stoichiometric formulas of metal oxides depending on the degree of oxidation of the metal and indicates which metals are capable of forming oxides of a given stoichiometric type.

In addition to such oxides, which in the general case can be described by the formula MeO X / 2, where X is the oxidation state of the metal, there are also oxides containing the metal in different oxidation states, for example, Fe 3 O 4 , as well as the so-called mixed oxides, e.g. FeO . Cr2O3.

Not all metal oxides have a constant composition; oxides of variable composition are known, for example, TiOx, where x = 0.88 - 1.20; FeOx, where x = 1.04 - 1.12, etc.

S-metal oxides have only one oxide each. Metals of p- and d-blocks, as a rule, have several oxides, with the exception of Al, Ga, In and d-elements of groups 3 and 12.

Oxides like MeO and Me 2 O 3 form almost all d-metals of 4 periods. Most d-metals of periods 5 and 6 are characterized by oxides in which the metal is in high oxidation states³ 4. Oxides of the MeO type form only Cd, Hg and Pd; type Me 2 O 3 , in addition to Y and La, form Au, Rh; silver and gold form oxides of the Me 2 O type.

Stoichiometric types of metal oxides

Oxidation state	Oxide type	Metals forming an oxide
+1	Me 2 O	Metals 1 and 11 groups
+2	MeO	Alld-metals 4 periods(except Sc), all metals 2 and 12 groups, as well as Sn, Pb; Cd, Hg and Pd
+3	Me 2 O 3	Almost alld-metals 4 periods(except for Cu and Zn), all metals of groups 3 and 13, Au, Rh
+4	MeO 2	Metals 4 and 14 groups and many other d-metals: V, Nb, Ta; Cr, Mo, W; Mn, Tc, Re; Ru, Os; Ir, Pt
+5	Me 2 O 5	Metals5 and 15 groups
+6	MeO 3	Metals6 groups
+7	Me 2 O 7	Metals7 groups
+8	MeO 4	Os and Ru

Structure of oxides

The vast majority of metal oxides under normal conditions- they are crystalline solids. The exception is the acidic oxide Mn 2 O 7 (it is a dark green liquid). Only very few crystals of acid metal oxides have a molecular structure, these are acid oxides with a metal in a very high oxidation state: RuO 4, OsO4, Mn 2 O 7, Tc 2 O 7, Re 2 O 7.

In the most general form, the structure of many crystalline metal oxides can be represented as a regular three-dimensional arrangement of oxygen atoms in space; metal atoms are located in the voids between the oxygen atoms. Since oxygen is a very electronegative element, it pulls some of the valence electrons from the metal atom, converting it into a cation, and oxygen itself goes into an anionic form and increases in size due to the addition of foreign electrons. Large oxygen anions form a crystal lattice, and metal cations are located in the voids between them. Only in metal oxides that are in a small degree of oxidation and have a small electronegativity value, the bond in oxides can be considered as ionic. Practically ionic are oxides of alkali and alkaline earth metals. In most metal oxides, the chemical bond is intermediate between ionic and covalent. With an increase in the degree of oxidation of the metal, the contribution of the covalent component increases.

Crystal structures of metal oxides

Coordination numbers of metals in oxides

The metal in oxides is characterized not only by the degree of oxidation, but also by the coordination number, indicating how many oxygen atoms it coordinates.

Very common in metal oxides is the coordination number 6, in this case the metal cation is in the center of an octahedron formed by six oxygen atoms. Octahedrons are packed into a crystal lattice in such a way that the stoichiometric ratio of metal and oxygen atoms is maintained. So in the crystal lattice of calcium oxide, the coordination number of calcium is 6. Oxygen octahedrons with the Ca 2+ cation in the center are combined with each other in such a way that each oxygen is surrounded by six calcium atoms, i.e. oxygen belongs simultaneously to 6 calcium atoms. Such a crystal is said to have (6, 6) coordination. The first is the coordination number of the cation, and the second is the coordination number of the anion. Thus, the formula for CaO oxide should be written
CaO 6/6 ≡ CaO.
In TiO 2 oxide, the metal is also in an octahedral environment of oxygen atoms, some of the oxygen atoms are connected by opposite edges, and some by vertices. In a TiO 2 rutile crystal, coordination (6, 3) means that oxygen belongs to three titanium atoms. Titanium atoms form a rectangular parallelepiped in the crystal lattice of rutile.

The crystal structures of oxides are quite diverse. Metals can be located not only in an octahedral environment of oxygen atoms, but also in a tetrahedral environment, for example, in the oxide BeO ≡ BeO 4|4. In PbO oxide, which also has crystal coordination (4.4), lead is at the top of a tetragonal prism, at the base of which there are oxygen atoms.

Metal atoms can be in different environments of oxygen atoms, for example, in octahedral and tetrahedral voids, and the metal is in different oxidation states., as for example, in magnetite Fe 3 O 4 ≡ FeO. Fe2O3.

Defects in crystal lattices explain the variability in the composition of some oxides.

The concept of spatial structures makes it possible to understand the reasons for the formation of mixed oxides. In the voids between the oxygen atoms, there can be atoms of not one metal, but two different ones., such as,
in chromite FeO . Cr2O3.

Rutile structure

Some physical properties of metal oxides

The vast majority of oxides at ordinary temperatures are solids. They have a lower density than metals.

Many metal oxides are refractory substances. This makes it possible to use refractory oxides as refractory materials for metallurgical furnaces.

CaO oxide is produced on an industrial scale in the amount of 109 million tons/year. It is used for lining furnaces. Oxides of BeO and MgO are also used as refractories. MgO oxide is one of the few refractories that is very resistant to the action of molten alkalis.

Sometimes the refractoriness of oxides creates problems in obtaining metals by electrolysis from their melts. So Al 2 O 3 oxide, which has a melting point of about 2000 ° C, has to be mixed with Na 3 cryolite in order to reduce the melting temperature to ~ 1000 ° C, and an electric current is passed through this melt.

Refractory are oxides of d-metals 5 and 6 periods Y 2 O 3 (2430), La 2 O 3 (2280), ZrO 2 (2700), HfO 2 (2080), Ta 2 O 5 (1870), Nb 2 O 5 (1490), as well as many oxides of period 4 d-metals (see table). All oxides of group 2 s-metals, as well as Al 2 O 3, Ga 2 O 3, SnO, SnO 2, PbO, have high melting points (see table).

Low melting points (about C) usually have acidic oxides: RuO 4 (25), OsO 4 (41); Te 2 O 7 (120), Re 2 O 7 (302), ReO 3 (160), CrO 3 (197). But some acid oxides have rather high melting points (o C): MoO 3 (801) WO 3 (1473), V 2 O 5 (680).

Some of the basic oxides of the d-elements that complete the series are fragile, melt at low temperatures, or decompose when heated. Decompose when heated HgO (400 o C), Au 2 O 3 (155), Au 2 O, Ag 2 O (200), PtO 2 (400).

When heated above 400 ° C, all alkali metal oxides also decompose with the formation of metal and peroxide. Oxide Li 2 O is more stable and decomposes at temperatures above 1000 o C.

The table below shows some characteristics of period 4 d-metals, as well as s- and p-metals.

Characteristics of s- and p-metal oxides

Me	Oxide	Colour	T pl., оС	Acid-base character
s-metals
Li	Li2O	white	All oxides decompose at T > 400 o C, Li 2 O at T > 1000 o C	All alkali metal oxides are basic, soluble in water
Na	Na2O	white
K	K2O	yellow
Rb	Rb2O	yellow
Cs	Cs2O	orange
Be	BeO	white	2580	amphoteric
mg	MgO	white	2850	basic
Ca	CaO	white	2614	Basic, limited solubility in water
Sr	SrO	white	2430
Ba	BaO	white	1923
p-metals
Al	Al2O3	white	2050	amphoteric
Ga	Ga2O3	yellow	1795	amphoteric
In	In 2 O 3	yellow	1910	amphoteric
Tl	Tl2O3	brown	716	amphoteric
	Tl2O	black	303	basic
sn	SNO	Navy blue	1040	amphoteric
	SnO 2	white	1630	amphoteric
Pb	PbO	red	Turns yellow at T > 490 o C	amphoteric
	PbO	yellow	1580	amphoteric
	Pb3O4	red	Diff.
	PbO2	black	Diff. At 300 o C	amphoteric

Chemical properties(see link)

Characteristics of d-metal oxides 4 periods

	Oxide	Colour	r, g/cm3	T pl., оС	- ΔGo, kJ/mol	- ΔHo, kJ/mol	Prevailing Acid-base character
sc	Sc2O3	white	3,9	2450	1637	1908	basic
Ti	TiO	brown	4,9	1780, p	490	526	basic
	Ti2O3	Violet	4,6	1830	1434	1518	basic
	TiO2	white	4,2	1870	945	944	amphoteric
V	VO	grey	5,8	1830	389	432	basic
	V 2 O 3	black	4,9	1970	1161	1219	basic
	VO2	blue	4,3	1545	1429	713	amphoteric
	V 2 O 5	orange	3,4	680	1054	1552	acid
Cr	Cr2O3	green	5,2	2335p	536	1141	amphoteric
	CrO3	red	2,8	197p	513	590	acid
Mn	MNO	Grey-green	5,2	1842	385	385	basic
	Mn2O3	brown	4,5	1000p	958	958	basic
	Mn3O4	brown	4,7	1560p	1388	1388
	MnO2	brown	5,0	535p	521	521	amphoteric
	Mn2O7	green	2,4	6.55p		726	acid
Fe	FeO	Black	5,7	1400	265	265	basic
	Fe 3 O 4	black	5,2	1540p	1117	1117
	Fe2O3	brown	5,3	1565 p	822	822	basic
co	COO	Grey-green	5,7	1830	213	239	basic
	Co 3 O 4	black	6,1	900p	754	887
Ni	NiO	Grey-green	7,4	1955	239	240	basic
Cu	Cu2O	orange	6,0	1242	151	173	basic
	CuO	black	6,4	800p	134	162	basic
Zn	ZnO	white	5,7	1975	348	351	amphoteric

Chemical properties(see link)

The acid-base character of oxides depends on the degree of oxidation of the metal and on the nature of the metal.

The lower the oxidation state, the stronger the basic properties.If the metal is in the oxidation state X £ 4 , then its oxide is either basic or amphoteric.

The higher the degree of oxidation, the more pronounced the acidic properties.. If the metal is in the oxidation state X ≥ 5 , then its hydroxide is acidic.

In addition to acidic and basic oxides, there are amphoteric oxides that simultaneously exhibit both acidic and basic properties..

All p-metal oxides are amphoteric, exceptTl 2 O.

Froms-metals, only Be has an amphoteric oxide.

Among d-metals, oxides are amphoteric ZnO, Cr 2 O 3, Fe 2 O 3, Au 2 O 3, and almost all metal oxides in the oxidation state+4 except for the basic ZrO 2 and HfO 2 .

Most oxides, including Cr 2 O 3 , Fe 2 O 3 and metal dioxides, exhibit amphotericity only when fused with alkalis. ZnO, VO 2 , Au 2 O 3 interact with alkali solutions.

For oxides, in addition to acid-base interactions, i.e., reactions between basic oxides and acids and acid oxides, as well as reactions of acid and amphoteric oxides with alkalis, redox reactions are also characteristic.

Redox properties of metal oxides

Since in any oxides the metal is in an oxidized state, all oxides, without exception, are capable of exhibiting oxidizing properties.

The most common reactions in pyrometallurgy- these are redox interactions between metal oxides and various reducing agents, leading to the production of a metal.

Examples

2Fe 2 O 3 + 3C \u003d 4Fe + 3CO 2

Fe 3 O 4 + 2C \u003d 3Fe + 2CO 2

MnO 2 + 2C \u003d Mn + 2CO

SnO 2 + C \u003d Sn + 2CO 2

ZnO + C = Zn + CO

Cr 2 O 3 + 2Al \u003d 2Cr + Al 2 O 3

WO 3 + 3H 2 \u003d W + 3H 2 O

If the metal has several oxidation states, then with a sufficient increase in temperature, it becomes possible to decompose the oxide with the release of oxygen.

4CuO \u003d 2Cu 2 O + O 2

3PbO 2 \u003d Pb 3 O 4 + O 2,

2Pb 3 O 4 \u003d O 2 + 6PbO

Some oxides, especially noble metal oxides, can decompose to form metal when heated.

2Ag 2 O \u003d 4Ag + O 2

2Au 2 O 3 \u003d 4Au + 3O 2

The strong oxidizing properties of some oxides are used in practice. For example,

The oxidizing properties of PbO 2 oxide are used in lead batteries, in which an electric current is obtained due to a chemical reaction between PbO 2 and metallic lead.

PbO 2 + Pb + 2H 2 SO 4 \u003d 2PbSO 4 + 2H 2 O

The oxidizing properties of MnO 2 are also used to generate electric current in galvanic cells (electric batteries).

2MnO 2 + Zn + 2NH 4 Cl = + 2MnOOH

The strong oxidizing properties of some oxides lead to their peculiar interaction with acids. So the oxides PbO 2 and MnO 2 are reduced when dissolved in concentrated hydrochloric acid.

MnO 2 + 4HCl \u003d MnCl 2 + Cl 2 + 2H 2 O
If a metal forms several oxides, then metal oxides in a lower oxidation state can oxidize, i.e., exhibit reducing properties.

Particularly strong reducing properties are exhibited by metal oxides in low and unstable oxidation states, such as, for example. TiO, VO, CrO. When dissolved in water, they are oxidized, restoring water. Their reactions with water are similar to the reactions of a metal with water.

2TiO + 2H 2 O = 2TiOOH + H 2 .

These are complex substances consisting of two chemical elements, one of which is oxygen with an oxidation state (-2). General formula of oxides: EmOn, where m- number of element atoms E, a n is the number of oxygen atoms. Oxides can be solid (sand SiO 2 , varieties of quartz), liquid (hydrogen oxide H 2 O), gaseous (carbon oxides: carbon dioxide CO 2 and carbon monoxide CO gases).

The nomenclature of chemical compounds has evolved with the accumulation of factual material. At first, while the number of known compounds was small, widely used trivial names, not reflecting the composition, structure and properties of a substance, - minium Pb 3 O 4, litharge RIO, magnesia MgO iron oxide Fe 3 O 4, laughing gas N 2 O, white arsenic As 2 O 3 The trivial nomenclature was replaced by semi-systematic nomenclature - indications of the number of oxygen atoms in the compound were included in the name: nitrous- for lower oxide- for higher degrees of oxidation; anhydride- for acidic oxides.

At present, the transition to modern nomenclature is almost complete. According to international nomenclature, in the title oxide, the valency of the element should be indicated; for example, SO 2 - sulfur (IV) oxide, SO 3 - sulfur (VI) oxide, CrO - chromium (II) oxide, Cr 2 O 3 - chromium (III) oxide, CrO 3 - chromium (VI) oxide.

According to their chemical properties, oxides are divided into salt-forming and non-salt-forming.

Types of oxides

Non-salt-forming such oxides are called that do not interact with either alkalis or acids and do not form salts. There are few of them, the composition includes non-metals.

Salt-forming Oxides are called those that react with acids or bases and form salt and water.

Among salt-forming oxides distinguish between oxides basic, acidic, amphoteric.

Basic oxides are oxides that correspond to bases. For example: CuO corresponds to the base Cu (OH) 2, Na 2 O - the base of NaOH, Cu 2 O - CuOH, etc.

Oxides in the periodic table

Typical reactions of basic oxides

1. Basic oxide + acid \u003d salt + water (exchange reaction):

2. Basic oxide + acid oxide = salt (compound reaction):

3. Basic oxide + water = alkali (compound reaction):

Acid oxides are those oxides to which acids correspond. These are non-metal oxides: N 2 O 5 corresponds to HNO 3, SO 3 - H 2 SO 4, CO 2 - H 2 CO 3, P 2 O 5 - H 4 PO 4 as well as metal oxides with a high value of oxidation states: Cr 2 + 6 O 3 corresponds to H 2 CrO 4 , Mn 2 +7 O 7 - HMnO 4 .

Typical reactions of acid oxides

1. Acid oxide + base \u003d salt + water (exchange reaction):

2. Acid oxide + basic oxide salt (compound reaction):

3. Acid oxide + water = acid (compound reaction):

Such a reaction is possible only if the acid oxide is soluble in water.

amphoteric called oxides, which, depending on the conditions, exhibit basic or acidic properties. These are ZnO, Al 2 O 3, Cr 2 O 3, V 2 O 5.

Amphoteric oxides do not combine directly with water.

Typical reactions of amphoteric oxides

1. Amphoteric oxide + acid \u003d salt + water (exchange reaction):

2. Amphoteric oxide + base \u003d salt + water or complex compound:

basic oxides. To main refer typical metal oxides, they correspond to hydroxides with the properties of bases.

Obtaining basic oxides

Oxidation of metals when heated in an oxygen atmosphere.

2Mg + O 2 \u003d 2MgO

2Cu + O 2 \u003d 2CuO

The method is not applicable for the production of alkali metal oxides. In reaction with oxygen, alkali metals usually give peroxides, so Na 2 O, K 2 O oxides are difficult to access.

Sulfide roasting

2CuS + 3O 2 = 2CuO + 2SO 2

4FeS 2 + 110 2 = 2Fe 2 O 3 + 8SO 2

The method is not applicable to active metal sulfides that oxidize to sulfates.

Decomposition of hydroxides

Cu(OH) 2 \u003d CuO + H 2 O

Thisthe method cannot be used to obtain oxides of alkali metals.

Decomposition of salts of oxygen-containing acids.

VaCO 3 \u003d BaO + CO 2

2Pb (NO 3) 2 \u003d 2PbO + 4N0 2 + O 2

4FeSO 4 \u003d 2Fe 2 O 3 + 4SO 2 + O 2

Decomposition is easily carried out for nitrates and carbonates, including basic salts.

2 CO 3 \u003d 2ZnO + CO 2 + H 2 O

Obtaining acid oxides

Acid oxides are represented by oxides of non-metals or transition metals in high oxidation states. They can be obtained by methods similar to those for basic oxides, for example:

1. 4P + 5O 2 \u003d 2P 2 O 5
2. 2ZnS + 3O 2 = 2ZnO + 2SO 2
3. K 2 Cr 2 O 7 + H 2 SO 4 \u003d 2CrO 3 ↓ + K 2 SO 4 + H 2 O
4. Na 2 SiO 3 + 2HCl \u003d 2NaCl + SiO 2 ↓ + H 2 O

Oxides, their classification and properties are the basis of such an important science as chemistry. They begin to study in the first year of study of chemistry. In such exact sciences as mathematics, physics and chemistry, all the material is interconnected, which is why the failure to assimilate the material entails a misunderstanding of new topics. Therefore, it is very important to understand the topic of oxides and fully navigate it. We will try to talk about this in more detail today.

What are oxides?

Oxides, their classification and properties - this is what needs to be understood paramount. So what are oxides? Do you remember this from the school curriculum?

Oxides (or oxides) are binary compounds, which include atoms of an electronegative element (less electronegative than oxygen) and oxygen with an oxidation state of -2.

Oxides are incredibly common substances on our planet. Examples of an oxide compound are water, rust, some dyes, sand, and even carbon dioxide.

Oxide formation

Oxides can be obtained in a variety of ways. The formation of oxides is also studied by such a science as chemistry. Oxides, their classification and properties - that's what scientists need to know in order to understand how this or that oxide was formed. For example, they can be obtained by direct connection of an oxygen atom (or atoms) with a chemical element - this is the interaction of chemical elements. However, there is also an indirect formation of oxides, this is when oxides are formed by the decomposition of acids, salts or bases.

Classification of oxides

Oxides and their classification depend on how they were formed. According to their classification, oxides are divided into only two groups, the first of which is salt-forming, and the second is non-salt-forming. So, let's take a closer look at both groups.

Salt-forming oxides are a fairly large group, which is divided into amphoteric, acidic and basic oxides. As a result of any chemical reaction, salt-forming oxides form salts. As a rule, the composition of salt-forming oxides includes elements of metals and non-metals, which, as a result of a chemical reaction with water, form acids, but when interacting with bases, they form the corresponding acids and salts.

Non-salt-forming oxides are oxides that do not form salts as a result of a chemical reaction. Examples of such oxides are carbon.

Amphoteric oxides

Oxides, their classification and properties are very important concepts in chemistry. Salt-forming compounds include amphoteric oxides.

Amphoteric oxides are oxides that can exhibit basic or acidic properties, depending on the conditions of chemical reactions (show amphotericity). Such oxides are formed (copper, silver, gold, iron, ruthenium, tungsten, rutherfordium, titanium, yttrium, and many others). Amphoteric oxides react with strong acids, and as a result of a chemical reaction they form salts of these acids.

Acid oxides

Or anhydrides are such oxides that, in chemical reactions, exhibit and also form oxygen-containing acids. Anhydrides are always formed by typical non-metals, as well as some transitional chemical elements.

Oxides, their classification and chemical properties are important concepts. For example, acidic oxides have completely different chemical properties from amphoteric ones. For example, when an anhydride interacts with water, the corresponding acid is formed (the exception is SiO2 - Anhydrides interact with alkalis, and as a result of such reactions, water and soda are released. When interacting with, a salt is formed.

Basic oxides

Basic (from the word "base") oxides are oxides of the chemical elements of metals with oxidation states of +1 or +2. These include alkali, alkaline earth metals, as well as the chemical element magnesium. Basic oxides differ from others in that they are able to react with acids.

Basic oxides interact with acids, in contrast to acid oxides, as well as with alkalis, water, and other oxides. As a result of these reactions, as a rule, salts are formed.

Properties of oxides

If you carefully study the reactions of various oxides, you can independently draw conclusions about what chemical properties the oxides are endowed with. The common chemical property of absolutely all oxides is the redox process.

Nevertheless, all oxides are different from each other. The classification and properties of oxides are two related topics.

Non-salt-forming oxides and their chemical properties

Non-salt-forming oxides are a group of oxides that exhibit neither acidic, nor basic, nor amphoteric properties. As a result of chemical reactions with non-salt-forming oxides, no salts are formed. Previously, such oxides were called not non-salt-forming, but indifferent and indifferent, but such names do not correspond to the properties of non-salt-forming oxides. According to their properties, these oxides are quite capable of chemical reactions. But there are very few non-salt-forming oxides; they are formed by monovalent and divalent non-metals.

Salt-forming oxides can be obtained from non-salt-forming oxides as a result of a chemical reaction.

Nomenclature

Almost all oxides are usually called like this: the word "oxide", followed by the name of the chemical element in the genitive case. For example, Al2O3 is aluminum oxide. In chemical language, this oxide is read like this: aluminum 2 o 3. Some chemical elements, such as copper, can have several degrees of oxidation, respectively, the oxides will also be different. Then CuO oxide is copper (two) oxide, that is, with an oxidation degree of 2, and Cu2O oxide is copper (three) oxide, which has an oxidation degree of 3.

But there are other names of oxides, which are distinguished by the number of oxygen atoms in the compound. A monoxide or monoxide is an oxide that contains only one oxygen atom. Dioxides are those oxides that contain two oxygen atoms, as indicated by the prefix "di". Trioxides are those oxides that already contain three oxygen atoms. Names such as monoxide, dioxide, and trioxide are already obsolete, but are often found in textbooks, books, and other manuals.

There are also so-called trivial names of oxides, that is, those that have developed historically. For example, CO is the oxide or monoxide of carbon, but even chemists most commonly refer to this substance as carbon monoxide.

So, an oxide is a combination of oxygen with a chemical element. The main science that studies their formation and interactions is chemistry. Oxides, their classification and properties are several important topics in the science of chemistry, without understanding which it is impossible to understand everything else. Oxides are gases, minerals, and powders. Some oxides should be known in detail not only by scientists, but also by ordinary people, because they can even be dangerous for life on this earth. Oxides are a very interesting and fairly easy topic. Oxide compounds are very common in everyday life.

Modern chemical science is a wide variety of branches, and each of them, in addition to the theoretical base, is of great applied and practical importance. Whatever you touch, everything around is the products of chemical production. The main sections are inorganic and organic chemistry. Consider what main classes of substances are classified as inorganic and what properties they have.

Main categories of inorganic compounds

These include the following:

1. Oxides.
2. Salt.
3. Foundations.
4. Acids.

Each of the classes is represented by a wide variety of inorganic compounds and is important in almost any structure of human economic and industrial activity. All the main properties characteristic of these compounds, being in nature and obtaining are studied in the school chemistry course without fail, in grades 8-11.

There is a general table of oxides, salts, bases, acids, which presents examples of each of the substances and their state of aggregation, being in nature. It also shows interactions that describe chemical properties. However, we will consider each of the classes separately and in more detail.

Group of compounds - oxides

4. Reactions, as a result of which elements change CO

Me + n O + C = Me 0 + CO

1. Reagent water: acid formation (SiO 2 exception)

KO + water = acid

2. Reactions with bases:

CO 2 + 2CsOH \u003d Cs 2 CO 3 + H 2 O

3. Reactions with basic oxides: salt formation

P 2 O 5 + 3MnO \u003d Mn 3 (PO 3) 2

4. OVR reactions:

CO 2 + 2Ca \u003d C + 2CaO,

They show dual properties, interact according to the principle of the acid-base method (with acids, alkalis, basic oxides, acid oxides). They do not interact with water.

1. With acids: formation of salts and water

AO + acid \u003d salt + H 2 O

2. With bases (alkalis): formation of hydroxo complexes

Al 2 O 3 + LiOH + water \u003d Li

3. Reactions with acid oxides: preparation of salts

FeO + SO 2 \u003d FeSO 3

4. Reactions with RO: formation of salts, fusion

MnO + Rb 2 O = double salt Rb 2 MnO 2

5. Fusion reactions with alkalis and alkali metal carbonates: formation of salts

Al 2 O 3 + 2LiOH \u003d 2LiAlO 2 + H 2 O

They do not form acids or alkalis. They exhibit highly specific properties.

Each higher oxide, formed both by a metal and a non-metal, when dissolved in water, gives a strong acid or alkali.

Acids organic and inorganic

In classical terms (based on the positions of ED - electrolytic dissociation - acids are compounds that dissociate into H + cations and An - acid residue anions in an aqueous medium. However, today acids have been carefully studied under anhydrous conditions, so there are many different theories for hydroxides.

Empirical formulas of oxides, bases, acids, salts are made up only of symbols, elements and indices indicating their amount in a substance. For example, inorganic acids are expressed by the formula H + acid residue n-. Organic substances have a different theoretical mapping. In addition to the empirical one, it is possible to write down a full and abbreviated structural formula for them, which will reflect not only the composition and amount of the molecule, but also the arrangement of atoms, their relationship to each other and the main functional group for carboxylic acids -COOH.

In the inorganic, all acids are divided into two groups:

• anoxic - HBr, HCN, HCL and others;
• oxygen-containing (oxo acids) - HClO 3 and everything where there is oxygen.

Also, inorganic acids are classified according to stability (stable or stable - everything except carbonic and sulphurous, unstable or unstable - carbonic and sulphurous). By strength, acids can be strong: sulfuric, hydrochloric, nitric, perchloric and others, as well as weak: hydrogen sulfide, hypochlorous and others.

Organic chemistry does not offer such diversity at all. Acids that are organic in nature are carboxylic acids. Their common feature is the presence of a functional group -COOH. For example, HCOOH (antic), CH 3 COOH (acetic), C 17 H 35 COOH (stearic) and others.

There are a number of acids, which are especially carefully emphasized when considering this topic in a school chemistry course.

1. Salt.
2. Nitrogen.
3. Orthophosphoric.
4. Hydrobromic.
5. Coal.
6. Iodine.
7. Sulfuric.
8. Acetic, or ethane.
9. Butane or oil.
10. Benzoic.

These 10 acids in chemistry are the fundamental substances of the corresponding class both in the school course and in general in industry and synthesis.

Properties of inorganic acids

The main physical properties should be attributed primarily to a different state of aggregation. After all, there are a number of acids that have the form of crystals or powders (boric, orthophosphoric) under normal conditions. The vast majority of known inorganic acids are different liquids. Boiling and melting points also vary.

Acids can cause severe burns, as they have the power to destroy organic tissues and skin. Indicators are used to detect acids:

• methyl orange (in normal environment - orange, in acids - red),
• litmus (in neutral - violet, in acids - red) or some others.

The most important chemical properties include the ability to interact with both simple and complex substances.

Chemical properties of inorganic acids

What do they interact with?	Reaction example
1. With simple substances-metals. Mandatory condition: the metal must stand in the ECHRNM before hydrogen, since the metals standing after hydrogen are not able to displace it from the composition of acids. As a result of the reaction, hydrogen is always formed in the form of a gas and a salt.
2. With bases. The result of the reaction is salt and water. Such reactions of strong acids with alkalis are called neutralization reactions.	Any acid (strong) + soluble base = salt and water
3. With amphoteric hydroxides. Bottom line: salt and water.	2HNO 2 + beryllium hydroxide \u003d Be (NO 2) 2 (medium salt) + 2H 2 O
4. With basic oxides. Outcome: water, salt.	2HCL + FeO = iron (II) chloride + H 2 O
5. With amphoteric oxides. Final effect: salt and water.	2HI + ZnO = ZnI 2 + H 2 O
6. With salts formed by weaker acids. Final effect: salt and weak acid.	2HBr + MgCO 3 = magnesium bromide + H 2 O + CO 2

When interacting with metals, not all acids react in the same way. Chemistry (grade 9) at school involves a very shallow study of such reactions, however, even at this level, the specific properties of concentrated nitric and sulfuric acid are considered when interacting with metals.

Hydroxides: alkalis, amphoteric and insoluble bases

Oxides, salts, bases, acids - all these classes of substances have a common chemical nature, which is explained by the structure of the crystal lattice, as well as the mutual influence of atoms in the composition of molecules. However, if for oxides it was possible to give a very specific definition, then for acids and bases it is more difficult to do so.

Just like acids, according to the ED theory, bases are substances that can decompose in an aqueous solution into metal cations Me n + and anions of hydroxo groups OH -.

• Soluble or alkali (strong bases that change Formed by metals of groups I, II. Example: KOH, NaOH, LiOH (that is, elements of only the main subgroups are taken into account);
• Slightly soluble or insoluble (medium strength, do not change the color of the indicators). Example: magnesium hydroxide, iron (II), (III) and others.
• Molecular (weak bases, in an aqueous medium they reversibly dissociate into ions-molecules). Example: N 2 H 4, amines, ammonia.
• Amphoteric hydroxides (show dual basic-acid properties). Example: beryllium, zinc and so on.

Each group represented is studied in the school chemistry course in the "Foundations" section. Chemistry grades 8-9 involves a detailed study of alkalis and sparingly soluble compounds.

The main characteristic properties of the bases

All alkalis and sparingly soluble compounds are found in nature in a solid crystalline state. At the same time, their melting points are, as a rule, low, and poorly soluble hydroxides decompose when heated. The base color is different. If the alkalis are white, then the crystals of sparingly soluble and molecular bases can be of very different colors. The solubility of most compounds of this class can be viewed in the table, which presents the formulas of oxides, bases, acids, salts, shows their solubility.

Alkalis are able to change the color of indicators as follows: phenolphthalein - raspberry, methyl orange - yellow. This is ensured by the free presence of hydroxo groups in solution. That is why sparingly soluble bases do not give such a reaction.

The chemical properties of each group of bases are different.

Chemical properties
alkalis	sparingly soluble bases	Amphoteric hydroxides
I. Interact with KO (total - salt and water): 2LiOH + SO 3 \u003d Li 2 SO 4 + water II. Interact with acids (salt and water): conventional neutralization reactions (see acids) III. Interact with AO to form a hydroxocomplex of salt and water: 2NaOH + Me + n O \u003d Na 2 Me + n O 2 + H 2 O, or Na 2 IV. Interact with amphoteric hydroxides to form hydroxo complex salts: The same as with AO, only without water V. Interact with soluble salts to form insoluble hydroxides and salts: 3CsOH + iron (III) chloride = Fe(OH) 3 + 3CsCl VI. Interact with zinc and aluminum in an aqueous solution to form salts and hydrogen: 2RbOH + 2Al + water = complex with hydroxide ion 2Rb + 3H 2	I. When heated, they can decompose: insoluble hydroxide = oxide + water II. Reactions with acids (total: salt and water): Fe(OH) 2 + 2HBr = FeBr 2 + water III. Interact with KO: Me + n (OH) n + KO \u003d salt + H 2 O	I. React with acids to form salt and water: (II) + 2HBr = CuBr 2 + water II. React with alkalis: result - salt and water (condition: fusion) Zn(OH) 2 + 2CsOH \u003d salt + 2H 2 O III. They react with strong hydroxides: the result is salts, if the reaction takes place in an aqueous solution: Cr(OH) 3 + 3RbOH = Rb 3

These are the most chemical properties that bases exhibit. The chemistry of bases is quite simple and obeys the general laws of all inorganic compounds.

Class of inorganic salts. Classification, physical properties

Based on the provisions of the ED, salts can be called inorganic compounds that dissociate in an aqueous solution into metal cations Me + n and anions of acid residues An n- . So you can imagine salt. Chemistry gives more than one definition, but this is the most accurate.

At the same time, according to their chemical nature, all salts are divided into:

• Acidic (containing a hydrogen cation). Example: NaHSO4.
• Basic (having a hydroxo group). Example: MgOHNO 3 , FeOHCL 2.
• Medium (consist only of a metal cation and an acid residue). Example: NaCL, CaSO 4.
• Double (include two different metal cations). Example: NaAl(SO 4) 3.
• Complex (hydroxocomplexes, aquacomplexes and others). Example: K 2 .

The formulas of salts reflect their chemical nature, and also speak of the qualitative and quantitative composition of the molecule.

Oxides, salts, bases, acids have different solubility, which can be seen in the corresponding table.

If we talk about the state of aggregation of salts, then you need to notice their uniformity. They exist only in a solid, crystalline or powdered state. The color scheme is quite varied. Solutions of complex salts, as a rule, have bright saturated colors.

Chemical interactions for the class of medium salts

They have similar chemical properties of bases, acids, salts. Oxides, as we have already considered, differ somewhat from them in this factor.

In total, 4 main types of interactions can be distinguished for medium salts.

I. Interaction with acids (only strong in terms of ED) with the formation of another salt and a weak acid:

KCNS + HCL = KCL + HCNS

II. Reactions with soluble hydroxides with the appearance of salts and insoluble bases:

CuSO 4 + 2LiOH = 2LiSO 4 soluble salt + Cu(OH) 2 insoluble base

III. Interaction with another soluble salt to form an insoluble salt and a soluble one:

PbCL 2 + Na 2 S = PbS + 2NaCL

IV. Reactions with metals to the left of the one that forms the salt in the EHRNM. In this case, the metal entering into the reaction should not, under normal conditions, interact with water:

Mg + 2AgCL = MgCL 2 + 2Ag

These are the main types of interactions that are characteristic of medium salts. The formulas of complex, basic, double and acidic salts speak for themselves about the specificity of the manifested chemical properties.

The formulas of oxides, bases, acids, salts reflect the chemical essence of all representatives of these classes of inorganic compounds, and in addition, give an idea of ​​the name of the substance and its physical properties. Therefore, special attention should be paid to their writing. A huge variety of compounds offers us a generally amazing science - chemistry. Oxides, bases, acids, salts - this is only part of the vast variety.