Big encyclopedia of oil and gas. How to determine the nature of the oxide

Chemical compounds consisting of oxygen and any other element of the periodic system are called oxides. Depending on their properties, they are classified into basic, amphoteric and acidic. The nature of oxides can be determined theoretically and practically.

You will need

• - periodic system;
• - glassware;
• - chemical reagents.

Instruction

You need to have a good idea of ​​how the properties of chemical elements change depending on their location in the D.I. table. Mendeleev. Therefore, repeat the periodic law, the electronic structure of atoms (the degree of oxidation of elements depends on it), and so on.

Without resorting to practical steps, you can establish the nature of the oxide using only the periodic table. After all, it is known that in periods, in the direction from left to right, the alkaline properties of oxides change to amphoteric, and then to acidic. For example, in period III, sodium oxide (Na2O) exhibits basic properties, the compound of aluminum with oxygen (Al2O3) is amphoteric, and chlorine oxide (ClO2) is acidic.

Keep in mind that in the main subgroups, the alkaline properties of oxides increase from top to bottom, while acidity, on the contrary, weakens. So, in group I, cesium oxide (CsO) has a stronger basicity than lithium oxide (LiO). In group V, nitric oxide (III) is acidic, and bismuth oxide (Bi2O5) is already basic.

Another way to determine the nature of oxides. Let's suppose that the task is given to experimentally prove the basic, amphoteric and acidic properties of calcium oxide (CaO), pentavalent phosphorus oxide (P2O5(V)) and zinc oxide (ZnO).

First, take two clean test tubes. From the bottles, using a chemical spatula, pour some CaO into one and P2O5 into the other. Then pour 5-10 ml of distilled water into both reagents. Stir with a glass rod until the powder is completely dissolved. Dip pieces of litmus paper into both test tubes. Where calcium oxide is located, the indicator will become of blue color, which is proof of the basic character of the compound under study. In a test tube with phosphorus (V) oxide, the paper will turn red, therefore, P2O5 is an acidic oxide.

Since zinc oxide is insoluble in water, test it with acid and hydroxide to prove it is amphoteric. In either case, ZnO crystals will enter into a chemical reaction. For example:
ZnO + 2KOH = K2ZnO2 + H2O
3ZnO + 2H3PO4 Zn3(PO4)2 + 3H2O

note

Remember, the nature of the properties of the oxide directly depends on the valence of the element included in its composition.

Helpful advice

Do not forget that there are still so-called indifferent (non-salt-forming) oxides that do not react under normal conditions with either hydroxides or acids. These include oxides of non-metals with valences I and II, for example: SiO, CO, NO, N2O, etc., but there are also “metallic” ones: MnO2 and some others.

All interesting

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Instruction

You need to have a good idea of ​​how the properties of chemical elements change depending on their location in the D.I. table. Mendeleev. Therefore, repeat, the electronic structure of atoms (the degree of oxidation of elements depends on it), and so on.

Without resorting to practical steps, you can establish the nature of the oxide using only the periodic table. After all, it is known that in periods, in the direction from left to right, the alkaline properties of oxides change to amphoteric, and then to acidic. For example, in the III period, sodium oxide (Na2O) has the main properties, the compound of aluminum with oxygen (Al2O3) has a character, and chlorine oxide (ClO2) -.

Keep in mind that in the main subgroups, the alkaline properties of oxides increase from top to bottom, while acidity, on the contrary, weakens. So, in group I, cesium oxide (CsO) has a stronger basicity than lithium oxide (LiO). In group V, nitric oxide (III) is acidic, and oxide (Bi2O5) is already basic.

First, take two clean test tubes. From the bottles, using a chemical spatula, pour some CaO into one and P2O5 into the other. Then pour 5-10 ml of distilled water into both reagents. Stir with a glass rod until the powder is completely dissolved. Dip pieces of litmus paper into both test tubes. There, - the indicator will turn blue, which is proof of the basic nature of the compound under study. In a test tube with phosphorus (V) oxide, the paper will turn red, therefore, P2O5 -.

Since zinc oxide is insoluble in water, test it with acid and hydroxide to prove it is amphoteric. In either case, ZnO crystals will enter into a chemical reaction. For example:
ZnO + 2KOH = K2ZnO2 + H2O
3ZnO + 2H3PO4→ Zn3(PO4)2↓ + 3H2O

note

Remember, the nature of the properties of the oxide directly depends on the valence of the element included in its composition.

Helpful advice

Do not forget that there are still so-called indifferent (non-salt-forming) oxides that do not react under normal conditions with either hydroxides or acids. These include oxides of non-metals with valences I and II, for example: SiO, CO, NO, N2O, etc., but there are also “metallic” ones: MnO2 and some others.

Sources:

• basic character of oxides

Oxide calcium- This is ordinary quicklime. But, despite such a simple nature, this substance is very widely used in economic activity. From construction, as a base for lime cement, to cooking, as a food additive E-529 oxide calcium finds application. Oxide can be obtained both in industrial and at home conditions calcium from carbonate calcium thermal decomposition reaction.

You will need

• Calcium carbonate in the form of limestone or chalk. Ceramic crucible for annealing. Propane or acetylene torch.

Instruction

Prepare the crucible for carbonate annealing. Mount it firmly on fireproof supports or special fixtures. The crucible must be firmly installed and, if possible, secured.

Grind the carbonate calcium. Grinding must be done for better heat transfer inside. It is not necessary to grind limestone or chalk into dust. It is enough to produce a rough inhomogeneous grinding.

Fill the annealing crucible with crushed carbonate calcium. Do not fill the crucible completely, because when carbon dioxide is released, part of the substance may be thrown out. Fill the crucible to about a third or less.

Start heating the crucible. Install well and secure it. Carry out a smooth heating of the crucible from different sides in order to avoid its destruction due to uneven thermal expansion. Continue heating the crucible on the gas burner. After a while, the thermal decomposition of carbonate will begin calcium.

Wait complete passage thermal decay. During the reaction, the upper layers of the substance in the crucible can be poorly heated. They can be mixed several times with a steel spatula.

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note

Be careful when working with a gas burner and a heated crucible. During the reaction, the crucible will be heated to a temperature above 1200 degrees Celsius.

Helpful advice

Instead of trying to produce large quantities of calcium oxide on your own (for example, for the subsequent production of lime cement), it is better to buy a finished product from specialized trading floors.

Sources:

• Write down the reaction equations that you can use to

According to generally accepted views, acids are complex substances consisting of one or more hydrogen atoms that can be replaced by metal atoms and acid residues. They are divided into anoxic and oxygen-containing, monobasic and polybasic, strong, weak, etc. How to determine whether a substance has acidic properties?

You will need

• - indicator paper or litmus solution;
• - hydrochloric acid (preferably diluted);
• - sodium carbonate powder (soda ash);
• - a little silver nitrate in solution;
• - flat-bottomed flasks or beakers.

Instruction

The first and easiest test is a test using indicator litmus paper or litmus solution. If paper strip or the solution has a pink tint, which means that there are hydrogen ions in the test substance, and this is a sure sign of an acid. You can easily understand that the more intense the color (up to red-burgundy), the acid.

There are many other ways to check. For example, you are tasked with determining whether a clear liquid is hydrochloric acid. How to do it? You know the reaction to the chloride ion. It is detected by adding even the smallest amounts of lapis solution - AgNO3.

Pour a little of the investigated liquid into a separate container and drip a little bit of the lapis solution. In this case, a "curdled" white precipitate of insoluble silver chloride will instantly fall out. That is, there is definitely a chloride ion in the composition of a substance molecule. But maybe it's still not, but a solution of some kind of chlorine-containing salt? Like sodium chloride?

Remember another property of acids. Strong acids (and hydrochloric acid, of course, is one of them) can displace weak acids from them. Place a little soda powder - Na2CO3 into a flask or beaker and slowly add the test liquid. If a hiss is immediately heard and the powder literally “boils” - there will be no doubt left - this is hydrochloric acid.

Why? Because such a reaction: 2HCl + Na2CO3 = 2NaCl + H2CO3. Carbonic acid was formed, which is so weak that it instantly decomposes into water and carbon dioxide. It was his bubbles that caused this "seething and hissing."

Related videos

note

Hydrochloric acid, even diluted, is corrosive! Remember safety precautions.

Helpful advice

In no case should you resort to taste tests (if the tongue is sour, then there is acid). At the very least, it can be very dangerous! After all, many acids are extremely caustic.

Sources:

• how acid properties change in 2019

Phosphorus is a chemical element that has the 15th serial number in the periodic table. It is located in her V group. A classic non-metal discovered by the alchemist Brand in 1669. There are three main modifications of phosphorus: red (which is part of the mixture for lighting matches), white and black. At very high pressures (of the order of 8.3 * 10^10Pa), black phosphorus passes into another allotropic state (“metallic phosphorus”) and begins to conduct current. phosphorus in various substances?

Instruction

Remember degree. This is the value corresponding to the charge of the ion in the molecule, provided that the electron pairs that carry out the bond are shifted towards the more electronegative element (located to the right and above in the Periodic Table).

It is also necessary to know the main condition: the sum of the electric charges of all the ions that make up the molecule, taking into account the coefficients, must always be equal to zero.

The oxidation state does not always quantitatively coincide with the valency. best example- carbon, which in organic always has equal to 4, and the oxidation state can be equal to -4, and 0, and +2, and +4.

What is the oxidation state in a phosphine PH3 molecule, for example? With all that said, this question is very easy to answer. Since hydrogen is the very first element in the Periodic Table, it, by definition, cannot be located there "more to the right and higher" than. Therefore, it is phosphorus that will attract hydrogen electrons to itself.

Each hydrogen atom, having lost an electron, will turn into a positively charged oxidation ion +1. Therefore, the total positive charge is +3. Hence, taking into account the rule that the total charge of the molecule is zero, the oxidation state of phosphorus in the phosphine molecule is -3.

Well, what is the oxidation state of phosphorus in P2O5 oxide? Take the periodic table. Oxygen is located in group VI, to the right of phosphorus, and also higher, therefore, it is definitely more electronegative. That is, the oxidation state of oxygen in this compound will be with a minus sign, and phosphorus with a plus sign. What are these degrees so that the molecule as a whole is neutral? It can be easily seen that the least common multiple of the numbers 2 and 5 is 10. Therefore, the oxidation state of oxygen is -2, and that of phosphorus is +5.

Related videos

Oxides are called complex substances consisting of two elements, one of which is oxygen (K - O - K; Ca "O; 0" Sb0, etc.). All oxides are divided into non-salt and salt-forming. A few non-salt-forming oxides do not interact with either acids or bases. These include nitric oxide (I) N20, nitric oxide (I) N0, etc. Salt-forming oxides are divided into basic, acidic and amphoteric. Basic oxides are called oxides, which form salts when interacting with acids or acid oxides. So, for example: CuO + H2S04 - CuS04 + H20, MgO + CO2 = MgC03. Only metal oxides can be basic. However, not all metal oxides are basic - many of them are amphoteric or acidic (for example, Cr203 is amphoteric, and Cr03 is acidic oxide). Part of the basic oxides dissolves in water, forming the corresponding bases: Na20 + H20 - 2NaOH. Acidic oxides are oxides that form salts when interacting with bases or basic oxides. So, for example: S02 + 2K0H - K2S03 + H20, P4O10 + bCaO \u003d 2Ca3 (P04) 2. Acidic oxides are typical non-metal oxides, as well as oxides of a number of metals in higher oxidation states (B203; N205; Mn207). Many acidic oxides (also called anhydrides) combine with water to form acids: N203 + H20 - 2HN02. Amphoteric are oxides that form salts when interacting with both acids and bases. Amphoteric oxides include: ZnO; A1203; Cr203; Mn02; Fe203, etc. For example, the amphoteric nature of zinc oxide manifests itself when it interacts with both hydrochloric acid and potassium hydroxide: ZnO + 2HC1 = ZnCl2 + H20, ZnO + 2 KOH = K2Zn02 + H20, ZnO + 2KOH + H20 - K2. The amphoteric nature of oxides, insoluble in acid solutions, and hydroxides is proved using more complex reactions. Thus, calcined oxides of aluminum and chromium (III) are practically insoluble in acid solutions and in alkalis. In the reaction of their fusion with potassium disulfate, the main properties of oxides are manifested: Al203 + 3K2S207 - 3K2S04 + Al2(S04)3. When fused with hydroxides, the acidic properties of oxides are revealed: A1203 + 2KOH - 2KA102 4- H20. Thus, amphoteric oxides have the properties of both basic and acidic oxides. Note that for various amphoteric oxides, the duality of properties can be expressed in terms of varying degrees. For example, zinc oxide is equally easily soluble in both acids and alkalis, i.e., in this oxide, the basic and acidic functions are approximately equally expressed. Iron oxide (III) - Fe203 - has predominantly basic properties; exhibits acidic properties only by interacting with alkalis at high temperatures: Fe203 + 2NaOH - 2NaFe02 + H20. Methods for obtaining oxides [T] Obtaining from simple substances: 2Ca + 02 = 2CaO. \2\ Decomposition complex substances: a) decomposition of oxides 4CrO3 = 2Cr2O3 + 302!; b) decomposition of hydroxides Ca(OH)2 = CaO + H20; c) decomposition of acids H2CO3 = H2O + CO2T; d) decomposition of salts Interaction of acids - oxidizing agents with metals and nonmetals: high temperature: Na2COn + Si02 = Na2Si03 + С02 f. fusion Questions and tasks for independent solution L Specify which inorganic substances are called oxides. , acidic and amphoteric 2. Determine what type the following oxides belong to: CaO, SiO, BaO, Si02, S03, P4O10, FeO, CO, ZnO, Cr203, NO 3. Specify which bases correspond to the following oxides: Na20, CaO, A1203, CuO, FeO, Fe203 4. Indicate which acid anhydrides are the following oxides: С02, S02, S03, N203, N205, Cr03, P4O10 5. Indicate which of the following oxides are soluble in water: CaO, CuO, Cr203, Si02, FeO, K20, CO, N02, Cr03, ZnO, A1203 6. Specify those with which of the following substances carbon monoxide (IV) will react: S02, KOH, H20, Ca (OH) 2, CaO. 7. Write reaction equations reflecting the properties of the following basic oxides: FeO, Cs20, HgO, Bi203. Write reaction equations that prove the acidic nature of the following oxides: S03, Mn207, P4O10, Cr03, Si02. 9. Show how the amphoteric nature of the following oxides can be proved: ZnO, A1203, Cr203. 10. Using the example of reactions for producing sulfur oxide (IV), indicate the main methods for producing oxides. 11. Complete the equations of the following chemical reactions, reflecting the methods for obtaining oxides: 1) Li + 02 -> 2) Si2H6 + 02 - 3) PbS + 02 4) Ca3P2 + 02 5) A1 (OH) s - 6) Pb (N03) 2 U 7) HgCl2 + Ba(OH)2 8) MgC03 + HN03 - 9) Ca3(PO4)2 + SiO2 - 10) CO2 + C £ 11) Cu + HNO3(30o/o) £ 12) C + H2S04 ( conc) 12. Determine the formula of the oxide formed by an element with an oxidation state of +2, if it is known that 3.73 g of hydrochloric acid was required to dissolve 4.05 g of it. Answer: SIO. 13. When carbon monoxide (IV) reacted with caustic soda, 21 g of sodium bicarbonate was formed. Determine the volume of carbon monoxide (IV) and the mass of sodium hydroxide spent to obtain salt. Answer: 5.6 liters of CO2; 10 g NaOH. 14. During the electrolysis of 40 mol of water, 620 g of oxygen were released. Determine the oxygen output. Answer: 96.9%. Determine the mass of acid and medium salt, which can be obtained by reacting 5.6 liters of SO2 with potassium hydroxide. What is the mass of alkali in each individual case? Answer: 30g KHS03; 39.5 g K2SO3; 14 g KOH; 28 g CON. 16. Determine the simplest formula compound containing 68.4% chromium and 31.6% oxygen. Answer: SG203. 17. Determine the oxidation state of manganese in the oxide, if it is known that 1.02 g of oxygen falls on 1 g of manganese. Answer: +7. 18. In the oxide of a monovalent element, the mass fraction of oxygen is 53.3%. Name the element. Answer: lithium. 19. Determine the mass of water needed to dissolve 188 g of potassium oxide, if you get a solution with mass fraction KOH 5.6%. Answer: 3812. 20. When 32 g of iron oxide (III) was reduced with carbon, 20.81 g of iron was formed. Determine the yield of iron. Answer: 90%.

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 (rarely +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 (rarely +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.

Acid. Oxides whose hydrates are oxygen-containing acids. Oxides of non-metals. Examples: P 2 O 3 - phosphorus oxide (III), CO 2 - carbon monoxide (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 an increase in the degree of oxidation of the metal

Physical Properties

Oxides are solid, liquid and gaseous, of various 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 normal 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 and others.

Gaseous:

CO 2 , SO 2 , N 2 O, NO, NO 2 and others.

Solubility in water

Soluble:

a) basic oxides of alkali and alkaline earth metals;

b) almost all acidic 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 scheme:

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

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

2. Redox properties

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

Examples of reactions in which oxides act as reducing agents:

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

2C +2 O + O 2 \u003d 2C +4 O 2

2S +4 O 2 + O 2 \u003d 2S +6 O 3

2N +2 O + O 2 \u003d 2N +4 O 2

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

C +2 O + FeO \u003d Fe + 2C +4 O 2

C +2 O + H 2 O \u003d H 2 + 2C +4 O 2

Examples of reactions in which oxides act as oxidizing agents:

Recovery of oxides with high o.d. elements to oxides with low s. O. or down to simple substances.

C +4 O 2 + C \u003d 2C +2 O

2S +6 O 3 + H 2 S \u003d 4S +4 O 2 + H 2 O

C +4 O 2 + Mg \u003d C 0 + 2MgO

Cr +3 2 O 3 + 2Al \u003d 2Cr 0 + 2Al 2 O 3

Cu +2 O + H 2 \u003d Cu 0 + H 2 O

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 \u003d NaNO 2 + NaNO 3 + H 2 O

How to get

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

4Li + O 2 = 2Li 2 O;

2Cu + O 2 \u003d 2CuO;

4P + 5O 2 \u003d 2P 2 O 5

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

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

2Al(OH) 3 \u003d Al 2 O 3 + 3H 2 O

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

H 2 SiO 3 \u003d SiO 2 + H 2 O

3. Decomposition of some salts:

2Cu(NO 3) 2 \u003d 2CuO + 4NO 2 + O 2

CaCO 3 \u003d CaO + CO 2

(CuOH) 2 CO 3 \u003d 2CuO + CO 2 + H 2 O

4. Oxidation of complex substances with oxygen:

CH 4 + 2O 2 \u003d CO 2 + H 2 O

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

4NH 3 + 5O 2 \u003d 4NO + 6H 2 O

5. Recovery of oxidizing acids by 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 (razb) + S \u003d H 2 SO 4 + 2NO

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

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.

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

In the very general view 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.

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, having a melting point of about 2000 o C, has to be mixed with Na 3 cryolite in order to lower the melting point to ~ 1000 o 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.

The acid-base character of oxides depends on the oxidation state of the metal to a greater extent than 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 less 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 more 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. Among d-metals, oxides are amphotericZnO, Cr2O 3 , Au 2 O 3 , PdO and almost all metal oxides in the +4 oxidation state except for the basic ZrO 2 and HfO 2 .