Properties of oxides
Oxides are complex chemical substances that are chemical compounds simple elements with oxygen. They are salt-forming And non-salt forming. In this case, there are 3 types of salt-forming agents: main(from the word "foundation"), acidic And amphoteric.
An example of oxides that do not form salts are: NO (nitric oxide) - is a colorless, odorless gas. It is formed during a thunderstorm in the atmosphere. CO (carbon monoxide) is an odorless gas produced by the combustion of coal. It is usually called carbon monoxide. There are other oxides that do not form salts. Now let's take a closer look at each type of salt-forming oxides.
Basic oxides
Basic oxides- these are complex chemical substances related to oxides that form salts when chemical reaction with acids or acidic oxides and do not react with bases or basic oxides. For example, the main ones include the following:
K 2 O (potassium oxide), CaO (calcium oxide), FeO (ferrous oxide).
Let's consider Chemical properties oxides with examples
1. Interaction with water:
- interaction with water to form a base (or alkali)
CaO+H 2 O → Ca(OH) 2 (known lime slaking reaction, which releases large quantities warmth!)
2. Interaction with acids:
- interaction with acid to form salt and water (salt solution in water)
CaO+H 2 SO 4 → CaSO 4 + H 2 O (Crystals of this substance CaSO 4 are known to everyone under the name “gypsum”).
3. Interaction with acid oxides: salt formation
CaO+CO 2 → CaCO 3 (Everyone knows this substance - ordinary chalk!)
Acidic oxides
Acidic oxides- these are complex chemical substances related to oxides that form salts upon chemical interaction with bases or basic oxides and do not interact with acidic oxides.
Examples of acidic oxides can be:
CO 2 (everyone knows carbon dioxide), P 2 O 5 - phosphorus oxide (formed by the combustion of white phosphorus in air), SO 3 - sulfur trioxide - this substance is used to produce sulfuric acid.
Chemical reaction with water
CO 2 +H 2 O→ H 2 CO 3 - this substance - carbonic acid- one of the weak acids, it is added to carbonated water to create gas “bubbles”. With increasing temperature, the solubility of gas in water decreases, and its excess comes out in the form of bubbles.
Reaction with alkalis (bases):
CO 2 +2NaOH→ Na 2 CO 3 +H 2 O- the resulting substance (salt) is widely used in the household. Its name - soda ash or washing soda - is an excellent detergent for burnt pots, grease, and burnt marks. With bare hands I don't recommend working!
Reaction with basic oxides:
CO 2 +MgO→ MgCO 3 - the resulting salt is magnesium carbonate - also called “bitter salt”.
Amphoteric oxides
Amphoteric oxides- these are complex chemical substances, also related to oxides, which form salts during chemical interaction with acids (or acid oxides) and grounds (or basic oxides). Most frequent use the word "amphoteric" in our case refers to metal oxides.
Example amphoteric oxides can be:
ZnO - zinc oxide ( White powder, often used in medicine for the manufacture of masks and creams), Al 2 O 3 - aluminum oxide (also called “alumina”).
The chemical properties of amphoteric oxides are unique in that they can enter into chemical reactions with both bases and acids. For example:
Reaction with acid oxide:
ZnO+H 2 CO 3 → ZnCO 3 + H 2 O - The resulting substance is a solution of the salt “zinc carbonate” in water.
Reaction with bases:
ZnO+2NaOH→ Na 2 ZnO 2 +H 2 O - the resulting substance - double salt sodium and zinc.
Obtaining oxides
Obtaining oxides produce different ways. This can happen through physical and chemical means. The most in a simple way is chemical reaction simple elements with oxygen. For example, the result of the combustion process or one of the products of this chemical reaction are oxides. For example, if a hot iron rod, and not only iron (you can take zinc Zn, tin Sn, lead Pb, copper Cu - basically whatever is at hand) is placed in a flask with oxygen, then a chemical reaction of iron oxidation will occur, which accompanied by a bright flash and sparks. The reaction product will be black iron oxide powder FeO:
2Fe+O 2 → 2FeO
Chemical reactions with other metals and non-metals are completely similar. Zinc burns in oxygen to form zinc oxide
2Zn+O 2 → 2ZnO
Coal combustion is accompanied by the formation of two oxides at once: carbon monoxide and carbon dioxide.
2C+O 2 → 2CO - formation of carbon monoxide.
C+O 2 → CO 2 - formation of carbon dioxide. This gas is formed if there is more than enough oxygen, that is, in any case, the reaction first occurs with the formation of carbon monoxide, and then the carbon monoxide is oxidized, turning into carbon dioxide.
Obtaining oxides can be done in another way - through a chemical decomposition reaction. For example, to obtain iron oxide or aluminum oxide, it is necessary to calcinate the corresponding bases of these metals over a fire:
Fe(OH) 2 → FeO+H 2 O
Solid aluminum oxide - mineral corundum 
Iron(III) oxide. The surface of the planet Mars is reddish-orange in color due to the presence of iron (III) oxide in the soil. Solid aluminum oxide - corundum 
2Al(OH) 3 → Al 2 O 3 +3H 2 O,
as well as during the decomposition of individual acids:
H 2 CO 3 → H 2 O+CO 2 - decomposition of carbonic acid
H 2 SO 3 → H 2 O+SO 2 - decomposition of sulfurous acid
Obtaining oxides can be made from metal salts at high heat:
CaCO 3 → CaO+CO 2 - calcination of chalk produces calcium oxide (or quicklime) and carbon dioxide.
2Cu(NO 3) 2 → 2CuO + 4NO 2 + O 2 - in this decomposition reaction two oxides are obtained at once: copper CuO (black) and nitrogen NO 2 (it is also called brown gas because of its really brown color).
Another way in which oxides can be produced is through redox reactions.
Cu + 4HNO 3 (conc.) → Cu(NO 3) 2 + 2NO 2 + 2H 2 O
S + 2H 2 SO 4 (conc.) → 3SO 2 + 2H 2 O
Chlorine oxides
ClO2 molecule 
Cl 2 O 7 molecule 
Nitrous oxide N2O 
Nitrogenous anhydride N 2 O 3 
Nitric anhydride N 2 O 5 
Brown gas NO 2 The following are known chlorine oxides: Cl 2 O, ClO 2, Cl 2 O 6, Cl 2 O 7. All of them, with the exception of Cl 2 O 7, are yellow or orange in color and are not stable, especially ClO 2, Cl 2 O 6. All chlorine oxides are explosive and are very strong oxidizing agents.
Reacting with water, they form the corresponding oxygen-containing and chlorine-containing acids:
So, Cl 2 O - acid chlorine oxide hypochlorous acid.
Cl 2 O + H 2 O → 2HClO - Hypochlorous acid
ClO2 - acid chlorine oxide hypochlorous and hypochlorous acid, since during a chemical reaction with water it forms two of these acids at once:
ClO 2 + H 2 O→ HClO 2 + HClO 3
Cl 2 O 6 - too acid chlorine oxide perchloric and perchloric acids:
Cl 2 O 6 + H 2 O → HClO 3 + HClO 4
And finally, Cl 2 O 7 - a colorless liquid - acid chlorine oxide perchloric acid:
Cl 2 O 7 + H 2 O → 2HClO 4
Nitrogen oxides
Nitrogen is a gas that forms 5 different compounds with oxygen - 5 nitrogen oxides. Namely:
N2O- nitric oxide. Its other name is known in medicine as laughing gas or nitrous oxide- It is colorless, sweetish and pleasant to the taste of gas.
- NO - nitrogen monoxide- a colorless, odorless, tasteless gas.
- N 2 O 3 - nitrous anhydride- colorless crystalline substance
- NO 2 - nitrogen dioxide. Its other name is brown gas- the gas really has a brownish-brown color
- N 2 O 5 - nitric anhydride- blue liquid, boiling at a temperature of 3.5 0 C
Of all these listed nitrogen compounds, NO - nitrogen monoxide and NO 2 - nitrogen dioxide are of greatest interest in industry. Nitrogen monoxide(NO) and nitrous oxide N 2 O does not react with water or alkalis. (N 2 O 3) when reacting with water forms a weak and unstable nitrous acid HNO 2, which in air gradually turns into a more stable Chemical substance nitric acid Let's look at some chemical properties of nitrogen oxides:
Reaction with water:
2NO 2 + H 2 O → HNO 3 + HNO 2 - 2 acids are formed at once: Nitric acid HNO 3 and nitrous acid.
Reaction with alkali:
2NO 2 + 2NaOH → NaNO 3 + NaNO 2 + H 2 O - two salts are formed: sodium nitrate NaNO 3 (or sodium nitrate) and sodium nitrite (a salt of nitrous acid).
Reaction with salts:
2NO 2 + Na 2 CO 3 → NaNO 3 + NaNO 2 + CO 2 - two salts are formed: sodium nitrate and sodium nitrite, and carbon dioxide is released.
Nitrogen dioxide (NO 2) is obtained from nitrogen monoxide (NO) using a chemical reaction of combining with oxygen:
2NO + O 2 → 2NO 2
Iron oxides
Iron forms two oxide:FeO- iron oxide(2-valent) - black powder, which is obtained by reduction iron oxide(3-valent) carbon monoxide by the following chemical reaction:
Fe 2 O 3 +CO→ 2FeO+CO 2
This is a basic oxide that reacts easily with acids. It has reducing properties and quickly oxidizes into iron oxide(3-valent).
4FeO +O 2 → 2Fe 2 O 3
Iron oxide(3-valent) - red-brown powder (hematite), which has amphoteric properties (can interact with both acids and alkalis). But acid properties of this oxide are so weakly expressed that it is most often used as basic oxide.
There are also so-called mixed iron oxide Fe 3 O 4 . It is formed when iron burns and conducts well electricity and has magnetic properties(it is called magnetic iron ore or magnetite). If iron burns, then as a result of the combustion reaction, scale is formed, consisting of two oxides: iron oxide(III) and (II) valence.
Sulfur oxide
Sulphur dioxide SO 2 Sulfur oxide SO 2 - or sulphur dioxide refers to acid oxides, but does not form acid, although it dissolves perfectly in water - 40 liters of sulfur oxide in 1 liter of water (for ease of preparation chemical equations This solution is called sulfurous acid).
Under normal circumstances, it is a colorless gas with a pungent and suffocating odor of burnt sulfur. At a temperature of only -10 0 C it can be converted into a liquid state.
In the presence of a catalyst - vanadium oxide (V 2 O 5) sulfur oxide attaches oxygen and turns into sulfur trioxide
2SO 2 +O 2 → 2SO 3
Dissolved in water sulphur dioxide- sulfur oxide SO2 - oxidizes very slowly, as a result of which the solution itself turns into sulfuric acid
If sulphur dioxide pass an alkali, for example, sodium hydroxide, through a solution, then sodium sulfite is formed (or hydrosulfite - depending on how much alkali and sulfur dioxide you take)
NaOH + SO 2 → NaHSO 3 - sulphur dioxide taken in excess
2NaOH + SO 2 → Na 2 SO 3 + H 2 O
If sulfur dioxide does not react with water, then why is it water solution gives a sour reaction?! Yes, it does not react, but it itself oxidizes in water, adding oxygen to itself. And it turns out that free hydrogen atoms accumulate in water, which give an acidic reaction (you can check with some indicator!)
Oxides are substances in which the molecules consist of an oxygen atom with oxidation state - 2 and atoms of some second element.
Oxides are formed directly by the interaction of oxygen with another substance or indirectly by the decomposition of bases, salts, and acids. This type of compound is very common in nature and can exist in the form of gas, liquid or B earth's crust there are also oxides. So, sand, rust, and even ordinary water - that’s all
There are both salt-forming and non-salt-forming oxides. Salt-formers produce salts as a result of a chemical reaction. These include oxides of non-metals and metals, which in reaction with water form an acid, and in reaction with a base - salts, normal and acidic. Salt-forming agents include, for example,
Accordingly, it is impossible to obtain salt from non-salt-forming substances. Examples include dinitrogen oxide and
Salt-forming oxides are divided, in turn, into basic, acidic and amphoteric. Let's talk in more detail about the main ones.
So, basic oxides are oxides of some metals, the corresponding hydroxides belong to the class of bases. That is, when interacting with acid, such substances form water and salt. For example, these are K2O, CaO, MgO, etc. Under normal conditions, basic oxides are solid crystalline formations. The degree of oxidation of metals in such compounds, as a rule, does not exceed +2 or rarely +3.
Chemical properties of basic oxides
1. Reaction with acid
It is in the reaction with an acid that the oxide exhibits its basic properties, so a similar experiment can prove the type of a particular oxide. If salt and water are formed, then it is a basic oxide. Acidic oxides in such a reaction form an acid. And amphoteric ones can exhibit either acidic or basic properties - it depends on the conditions. These are the main differences between non-salt-forming oxides.
2. Reaction with water
Those oxides that are formed by metals from the electrical voltage range that are in front of magnesium interact with water. When reacting with water they form soluble bases. This is a group of alkaline earth oxides (barium oxide, lithium oxide, etc.). Acidic oxides form acid in water, while amphoteric oxides do not react to water.
3. Reaction with amphoteric and acidic oxides
Chemically opposite substances react with each other, forming salts. For example, basic oxides can interact with acidic ones, but do not react with other representatives of their group. The most active are the oxides of alkali metals, alkaline earths and magnesium. Even under normal conditions, they fuse with solid amphoteric oxides and with solid and gaseous acidic oxides. When reacting with acidic oxides, they form the corresponding salts.
But the basic oxides of other metals are less active and practically do not react with gaseous (acidic) oxides. They can only undergo an addition reaction when fused with solid acid oxides.
4. Redox properties
Oxides of active alkali metals do not exhibit pronounced reducing or oxidizing properties. Conversely, oxides of less active metals can be reduced by coal, hydrogen, ammonia or carbon monoxide.
Preparation of basic oxides
1. Decomposition of hydroxides: When heated, insoluble bases decompose into water and a basic oxide.
2. Oxidation of metals: an alkali metal, when burned in oxygen, forms a peroxide, which then, upon reduction, forms a basic oxide.
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Chemical properties of basic oxides
You can read in detail about oxides, their classification and methods of preparation. .
1. Interaction with water. Only basic oxides, which correspond to soluble hydroxides (alkalis), can react with water. Alkalies form alkali metals (lithium, sodium, potassium, rubidium and cesium) and alkaline earth metals (calcium, strontium, barium). Oxides of other metals do not react chemically with water. Magnesium oxide reacts with water when boiled.
CaO + H 2 O → Ca(OH) 2
CuO + H 2 O ≠
2. Interaction with acid oxides and acids. When basic oxides interact with acids, a salt of this acid and water are formed. When a basic oxide interacts with an acidic one, a salt is formed:
basic oxide + acid = salt + water
basic oxide + acidic oxide = salt
When basic oxides interact with acids and their oxides, the following rule applies:
At least one of the reagents must correspond to a strong hydroxide (alkali or strong acid).
In other words, basic oxides, which correspond to alkalis, react with all acidic oxides and their acids. Basic oxides, which correspond to insoluble hydroxides, react only with strong acids and their oxides (N 2 O 5, NO 2, SO 3, etc.).
3. Interaction with amphoteric oxides and hydroxides.
When basic oxides interact with amphoteric ones, salts are formed:
basic oxide + amphoteric oxide = salt
They interact with amphoteric oxides during fusion only basic oxides, which correspond to alkalis . This creates salt. The metal in the salt comes from the more basic oxide, the acid residue from the more acidic one. IN in this case the amphoteric oxide forms an acid residue.
K 2 O + Al 2 O 3 → 2KAlO 2
CuO + Al 2 O 3 ≠ (the reaction does not occur, because Cu(OH) 2 is an insoluble hydroxide)
(to determine the acidic residue, we add a water molecule to the formula of an amphoteric or acidic oxide: Al 2 O 3 + H 2 O = H 2 Al 2 O 4 and divide the resulting indices in half if the oxidation state of the element is odd: HAlO 2. The result is an aluminate ion AlO 2 - The charge of the ion can be easily determined by the number of attached hydrogen atoms - if there is 1 hydrogen atom, then the charge of the anion will be -1, if there are 2 hydrogens, then -2, etc.).
Amphoteric hydroxides decompose when heated, so they actually cannot react with basic oxides.
4. Interaction of basic oxides with reducing agents.
Thus, some metal ions are oxidizing agents (the more to the right in the voltage series, the stronger). When interacting with reducing agents, metals go into oxidation state 0.
4.1. Reduction with coal or carbon monoxide.
Carbon (coal) reduces from oxides only metals located in the activity series after aluminum. The reaction occurs only when heated.
FeO + C → Fe + CO
Carbon monoxide also reduces from oxides only metals located after aluminum in the electrochemical series:
Fe 2 O 3 + CO → Al 2 O 3 + CO 2
CuO + CO → Cu + CO 2
4.2. Reduction with hydrogen .
Hydrogen reduces from oxides only metals located in the activity series to the right of aluminum. The reaction with hydrogen occurs only under harsh conditions - under pressure and heating.
CuO + H 2 → Cu + H 2 O
4.3. Reduction with more active metals (in melt or solution, depending on the metal)
In this case, more active metals displace less active ones. That is, the metal added to the oxide must be located to the left in the activity series than the metal from the oxide. Reactions usually occur when heated.
For example , Zinc oxide reacts with aluminum:
3ZnO + 2Al → Al 2 O 3 + 3Zn
but does not interact with copper:
ZnO + Cu ≠
Reducing metals from oxides using other metals is a very common process. Aluminum and magnesium are often used to restore metals. But alkali metals are not very suitable for this - they are too chemically active, which creates difficulties when working with them.
For example, cesium explodes in air.
Aluminothermy– is the reduction of metals from oxides with aluminum.
For example : aluminum reduces copper(II) oxide from the oxide:
3CuO + 2Al → Al 2 O 3 + 3Cu
Magniethermy– is the reduction of metals from oxides with magnesium.
CuO + H 2 → Cu + H 2 O
4.4. Reduction with ammonia.
Only oxides of inactive metals can be reduced with ammonia. The reaction only occurs at high temperatures.
For example , ammonia reduces copper(II) oxide:
3CuO + 2NH 3 → 3Cu + 3H 2 O + N 2
5. Interaction of basic oxides with oxidizing agents.
Under the influence of oxidizing agents, some basic oxides (in which metals can increase the oxidation state, for example Fe 2+, Cr 2+, Mn 2+, etc.) can act as reducing agents.
For example ,Iron(II) oxide can be oxidized with oxygen to iron(III) oxide:
4FeO + O 2 → 2Fe 2 O 3
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 reacts 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, since 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 a medium or basic salt (as a rule, a less soluble substance is obtained): 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 carbonic and sulfurous acids 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(sediment, gas, water).
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 weak acids (carbonates, sulfides, sulfites, nitrites), and a gaseous product is released.
Reactions between concentrated acids and crystalline salts are possible if a more volatile acid is obtained: 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 most 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
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 oxide (III), 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 black, calcium oxide CaO 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 acid 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 until 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 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).
