The hydrogen atom has the electronic formula of the outer (and only) electron level 1 s 1 . On the one hand, in terms of the presence of one electron on the outer electronic level, the hydrogen atom is similar to alkali metal atoms. However, just like halogens, it only needs one electron to fill the outer electronic level, since the first electronic level can contain no more than 2 electrons. It turns out that hydrogen can be placed simultaneously in both the first and the penultimate (seventh) group of the periodic table, which is sometimes done in various versions periodic table:
From the point of view of the properties of hydrogen as a simple substance, it still has more in common with halogens. Hydrogen, like halogens, is a non-metal and forms diatomic molecules (H 2) like them.
IN normal conditions Hydrogen is a gaseous, low-active substance. The low activity of hydrogen is explained by the high strength of the bonds between the hydrogen atoms in the molecule, the breaking of which requires either strong heating, or the use of catalysts, or both.
Interaction of hydrogen with simple substances
with metals
Of the metals, hydrogen reacts only with alkali and alkaline earth metals! Alkali metals include the main metals subgroup I groups (Li, Na, K, Rb, Cs, Fr), and alkaline earth metals - metals of the main subgroup I Group I, except beryllium and magnesium (Ca, Sr, Ba, Ra)
When interacting with active metals, hydrogen exhibits oxidizing properties, i.e. lowers its oxidation state. In this case, hydrides of alkali and alkaline earth metals are formed, which have an ionic structure. The reaction occurs when heated:
It should be noted that interaction with active metals is the only case when molecular hydrogen H2 is an oxidizing agent.
with non-metals
Of the non-metals, hydrogen reacts only with carbon, nitrogen, oxygen, sulfur, selenium and halogens!
Carbon should be understood as graphite or amorphous carbon, since diamond is an extremely inert allotropic modification of carbon.
When interacting with non-metals, hydrogen can only perform the function of a reducing agent, that is, only increase its oxidation state:
Interaction of hydrogen with complex substances
with metal oxides
Hydrogen does not react with metal oxides that are in the activity series of metals up to aluminum (inclusive), however, it is capable of reducing many metal oxides to the right of aluminum when heated:
with non-metal oxides
Of the non-metal oxides, hydrogen reacts when heated with the oxides of nitrogen, halogens and carbon. Of all the interactions of hydrogen with non-metal oxides, especially noteworthy is its reaction with carbon monoxide CO.
The mixture of CO and H2 even has its own name - “synthesis gas”, since, depending on the conditions, such popular industrial products as methanol, formaldehyde and even synthetic hydrocarbons can be obtained from it:
with acids
WITH inorganic acids Hydrogen does not react!
Of organic acids, hydrogen reacts only with unsaturated acids, as well as with acids containing functional groups capable of reduction with hydrogen, in particular aldehyde, keto or nitro groups.
with salts
In the case of aqueous solutions of salts, their interaction with hydrogen does not occur. However, when hydrogen is passed over solid salts of some metals of medium and low activity, their partial or complete reduction is possible, for example:
Chemical properties of halogens
Halogens are the chemical elements of group VIIA (F, Cl, Br, I, At), as well as the simple substances they form. Here and further in the text, unless otherwise stated, halogens will be understood as simple substances.
All halogens have a molecular structure, which determines the low melting and boiling points of these substances. Halogen molecules are diatomic, i.e. their formula can be written as general view like Hal 2.
It should be noted such a specific physical property of iodine as its ability to sublimation or, in other words, sublimation. Sublimation, is a phenomenon in which a substance in a solid state does not melt when heated, but, bypassing the liquid phase, immediately passes into the gaseous state.
The electronic structure of the external energy level of an atom of any halogen has the form ns 2 np 5, where n is the number of the periodic table period in which the halogen is located. As you can see, the halogen atoms only need one electron to reach the eight-electron outer shell. From this it is logical to assume the predominantly oxidizing properties of free halogens, which is confirmed in practice. As is known, the electronegativity of nonmetals decreases when moving down a subgroup, and therefore the activity of halogens decreases in the series:
F 2 > Cl 2 > Br 2 > I 2
Interaction of halogens with simple substances
All halogens are highly active substances and react with the majority simple substances. However, it should be noted that fluorine, due to its extremely high reactivity, can react even with those simple substances with which other halogens cannot react. Such simple substances include oxygen, carbon (diamond), nitrogen, platinum, gold and some noble gases (xenon and krypton). Those. actually, fluorine does not react only with some noble gases.
The remaining halogens, i.e. chlorine, bromine and iodine are also active substances, but less active than fluorine. They react with almost all simple substances except oxygen, nitrogen, carbon in the form of diamond, platinum, gold and noble gases.
Interaction of halogens with non-metals
hydrogen
When all halogens interact with hydrogen, they form hydrogen halides With general formula HHal. In this case, the reaction of fluorine with hydrogen begins spontaneously even in the dark and proceeds with an explosion in accordance with the equation:
The reaction of chlorine with hydrogen can be initiated by intense ultraviolet irradiation or by heating. Also proceeds with explosion:
Bromine and iodine react with hydrogen only when heated, and at the same time, the reaction with iodine is reversible:
phosphorus
The interaction of fluorine with phosphorus leads to the oxidation of phosphorus to highest degree oxidation (+5). In this case, phosphorus pentafluoride is formed:
When chlorine and bromine interact with phosphorus, it is possible to obtain phosphorus halides both in the oxidation state + 3 and in the oxidation state +5, which depends on the proportions of the reacting substances:
Moreover, in the case of white phosphorus in an atmosphere of fluorine, chlorine or liquid bromine, the reaction begins spontaneously.
The interaction of phosphorus with iodine can lead to the formation of only phosphorus triodide due to its significantly lower oxidizing ability than that of other halogens:
gray
Fluorine oxidizes sulfur to the highest oxidation state +6, forming sulfur hexafluoride:
Chlorine and bromine react with sulfur, forming compounds containing sulfur in the oxidation states +1 and +2, which are extremely unusual for it. These interactions are very specific, and for passing the Unified State Exam in chemistry, the ability to write equations for these interactions is not necessary. Therefore, the following three equations are given rather for reference:
Interaction of halogens with metals
As mentioned above, fluorine is capable of reacting with all metals, even such inactive ones as platinum and gold:
The remaining halogens react with all metals except platinum and gold:
Reactions of halogens with complex substances
Substitution reactions with halogens
More active halogens, i.e. the chemical elements of which are located higher in the periodic table are capable of displacing less active halogens from the hydrohalic acids and metal halides they form:
Similarly, bromine and iodine displace sulfur from solutions of sulfides and or hydrogen sulfide:
Chlorine is a stronger oxidizing agent and oxidizes hydrogen sulfide in its aqueous solution not to sulfur, but to sulfuric acid:
Reaction of halogens with water
Water burns in fluorine with a blue flame in accordance with the reaction equation:
Bromine and chlorine react differently with water than fluorine. If fluorine acted as an oxidizing agent, then chlorine and bromine are disproportionate in water, forming a mixture of acids. In this case, the reactions are reversible:
The interaction of iodine with water occurs to such an insignificant degree that it can be neglected and it can be assumed that the reaction does not occur at all.
Interaction of halogens with alkali solutions
Fluorine, when interacting with an aqueous alkali solution, again acts as an oxidizing agent:
Ability to take notes given equation not required to pass the Unified State Exam. It is enough to know the fact about the possibility of such an interaction and the oxidative role of fluorine in this reaction.
Unlike fluorine, other halogens in alkali solutions are disproportionate, that is, they simultaneously increase and decrease their oxidation state. In this case, in the case of chlorine and bromine, depending on the temperature, it is possible to flow through two different directions. In particular, in the cold the reactions proceed as follows:
and when heated:
Iodine reacts with alkalis exclusively according to the second option, i.e. with the formation of iodate, because hypoiodite is not stable not only when heated, but also at ordinary temperatures and even in the cold.
Hydrogen is a special element that occupies two cells at once in Mendeleev’s periodic table. It is located in two groups of elements that have opposite properties, and this feature makes it unique. Hydrogen is a simple substance and integral part many complex compounds, it is an organogenic and biogenic element. It is worth familiarizing yourself in detail with its main features and properties.
Hydrogen in the periodic table of Mendeleev
The main features of hydrogen indicated in:
- the serial number of the element is 1 (there are the same number of protons and electrons);
- atomic mass is 1.00795;
- hydrogen has three isotopes, each of which has special properties;
- due to the content of only one electron, hydrogen is capable of exhibiting reducing and oxidizing properties, and after donating an electron, hydrogen has a free orbital that takes part in the formation of chemical bonds according to the donor-acceptor mechanism;
- hydrogen is a light element with low density;
- hydrogen is a strong reducing agent, it opens the group of alkali metals in the first group to the main subgroup;
- when hydrogen reacts with metals and other strong reducing agents, it accepts their electron and becomes an oxidizing agent. Such compounds are called hydrides. According to this characteristic, hydrogen conventionally belongs to the group of halogens (in the table it is given above fluorine in parentheses), with which it is similar.
Hydrogen as a simple substance
Hydrogen is a gas whose molecule consists of two. This substance was discovered in 1766 by the British scientist Henry Cavendish. He proved that hydrogen is a gas that explodes when it reacts with oxygen. After studying hydrogen, chemists found that this substance is the lightest of all known to man.
Another scientist, Lavoisier, gave the element the name “hydrogenium,” which translated from Latin means “giving birth to water.” In 1781, Henry Cavendish proved that water is a combination of oxygen and hydrogen. In other words, water is the product of the reaction of hydrogen with oxygen. The flammable properties of hydrogen were known to ancient scientists: the corresponding records were left by Paracelsus, who lived in the 16th century.
Molecular hydrogen is a product formed naturally a gaseous compound common in nature, which consists of two atoms and when brought to the surface of a burning splinter. A hydrogen molecule can disintegrate into atoms that turn into helium nuclei, since they are capable of participating in nuclear reactions. Such processes regularly occur in space and on the Sun.
Hydrogen and its physical properties
Hydrogen has the following physical parameters:
- boils at -252.76 °C;
- melts at -259.14 °C; *within the specified temperature limits, hydrogen is an odorless, colorless liquid;
- Hydrogen is slightly soluble in water;
- hydrogen can theoretically transform into a metallic state if special conditions are provided (low temperatures and high pressure);
- pure hydrogen is an explosive and flammable substance;
- hydrogen is able to diffuse through the thickness of metals, therefore it dissolves well in them;
- hydrogen is 14.5 times lighter than air;
- at high blood pressure snow-like crystals of solid hydrogen can be obtained.
Chemical properties of hydrogen
Laboratory methods:
- interaction of dilute acids with active metals and metals of intermediate activity;
- hydrolysis of metal hydrides;
- reaction of alkali and alkaline earth metals with water.
Hydrogen compounds:
Hydrogen halides; volatile hydrogen compounds of non-metals; hydrides; hydroxides; hydrogen hydroxide (water); hydrogen peroxide; organic compounds (proteins, fats, hydrocarbons, vitamins, lipids, essential oils, hormones). Click to see safe experiments to study the properties of proteins, fats and carbohydrates.
To collect the hydrogen produced, you need to hold the test tube upside down. Hydrogen cannot be collected as carbon dioxide, because it is much lighter than air. Hydrogen quickly evaporates, and when mixed with air (or in large accumulations) it explodes. Therefore, it is necessary to invert the test tube. Immediately after filling, the tube is closed with a rubber stopper.
To test the purity of hydrogen, you need to hold a lit match to the neck of the test tube. If a dull and quiet bang occurs, the gas is clean and air impurities are minimal. If the cotton is loud and whistling, the gas in the test tube is dirty and contains a large proportion of foreign components.
Attention! Do not try to repeat these experiments yourself!
Industrial methods for producing simple substances depend on the form in which the corresponding element is found in nature, that is, what can be the raw material for its production. Thus, oxygen available in a free state is obtained physically- release from liquid air. Almost all of hydrogen is in the form of compounds, so to obtain it they use chemical methods. In particular, decomposition reactions can be used. One way to produce hydrogen is through the decomposition of water by electric current.
Basic industrial method producing hydrogen - the reaction of methane, which is part of natural gas, with water. It is carried out at high temperature(it’s easy to see that when passing methane even through boiling water, no reaction occurs):
CH 4 + 2H 2 0 = CO 2 + 4H 2 - 165 kJ
In the laboratory, to obtain simple substances, they do not necessarily use natural raw materials, but choose those starting materials from which it is easier to isolate the required substance. For example, in the laboratory, oxygen is not obtained from the air. The same applies to the production of hydrogen. One of the laboratory methods for producing hydrogen, which is sometimes used in industry, is the decomposition of water by electric current.
Typically, hydrogen is produced in the laboratory by reacting zinc with hydrochloric acid.
In industry
1.Electrolysis of aqueous salt solutions:
2NaCl + 2H 2 O → H 2 + 2NaOH + Cl 2
2.Passing water vapor over hot coke at temperatures around 1000°C:
H 2 O + C ⇄ H 2 + CO
3.From natural gas.
Steam conversion: CH 4 + H 2 O ⇄ CO + 3H 2 (1000 °C) Catalytic oxidation with oxygen: 2CH 4 + O 2 ⇄ 2CO + 4H 2
4. Cracking and reforming of hydrocarbons during oil refining.
In the laboratory
1.The effect of dilute acids on metals. To carry out this reaction, zinc and hydrochloric acid are most often used:
Zn + 2HCl → ZnCl 2 + H 2
2.Interaction of calcium with water:
Ca + 2H 2 O → Ca(OH) 2 + H 2
3.Hydrolysis of hydrides:
NaH + H 2 O → NaOH + H 2
4.Effect of alkalis on zinc or aluminum:
2Al + 2NaOH + 6H 2 O → 2Na + 3H 2 Zn + 2KOH + 2H 2 O → K 2 + H 2
5.Using electrolysis. During the electrolysis of aqueous solutions of alkalis or acids, hydrogen is released at the cathode, for example:
2H 3 O + + 2e - → H 2 + 2H 2 O
- Bioreactor for hydrogen production
Physical properties
Hydrogen gas can exist in two forms (modifications) - in the form of ortho - and para-hydrogen.
In a molecule of orthohydrogen (mp. −259.10 °C, bp −252.56 °C) the nuclear spins are directed identically (parallel), and in parahydrogen (mp. −259.32 °C, bp. boiling point -252.89 °C) - opposite to each other (antiparallel).
Allotropic forms of hydrogen can be separated by adsorption on active carbon at liquid nitrogen temperature. At very low temperatures the equilibrium between orthohydrogen and parahydrogen is almost completely shifted towards the latter. At 80 K the ratio of forms is approximately 1:1. When heated, desorbed parahydrogen transforms into orthohydrogen until equilibrium is formed at room temperature mixtures (ortho-para: 75:25). Without a catalyst, the transformation occurs slowly, which makes it possible to study the properties of individual allotropic forms. The hydrogen molecule is diatomic - H₂. Under normal conditions, it is a colorless, odorless, and tasteless gas. Hydrogen is the lightest gas, its density is many times less than the density of air. Obviously, the smaller the mass of the molecules, the higher their speed at the same temperature. As the lightest molecules, hydrogen molecules move faster than the molecules of any other gas and thus can transfer heat from one body to another faster. It follows that hydrogen has the highest thermal conductivity among gaseous substances. Its thermal conductivity is approximately seven times higher than the thermal conductivity of air.
Chemical properties
Hydrogen molecules H₂ are quite strong, and in order for hydrogen to react, a lot of energy must be expended: H 2 = 2H - 432 kJ Therefore, at ordinary temperatures, hydrogen reacts only with very active metals, for example calcium, forming calcium hydride: Ca + H 2 = CaH 2 and with the only non-metal - fluorine, forming hydrogen fluoride: F 2 + H 2 = 2HF With most metals and non-metals, hydrogen reacts at elevated temperature or under other influences, such as lighting. It can “take away” oxygen from some oxides, for example: CuO + H 2 = Cu + H 2 0 The written equation reflects the reduction reaction. Reduction reactions are processes in which oxygen is removed from a compound; Substances that take away oxygen are called reducing agents (they themselves oxidize). Further, another definition of the concepts “oxidation” and “reduction” will be given. A this definition, historically the first, remains important today, especially in organic chemistry. The reduction reaction is the opposite of the oxidation reaction. Both of these reactions always occur simultaneously as one process: when one substance is oxidized (reduced), the reduction (oxidation) of another necessarily occurs simultaneously.
N 2 + 3H 2 → 2 NH 3
Forms with halogens hydrogen halides:
F 2 + H 2 → 2 HF, the reaction occurs explosively in the dark and at any temperature, Cl 2 + H 2 → 2 HCl, the reaction occurs explosively, only in the light.
It interacts with soot under high heat:
C + 2H 2 → CH 4
Interaction with alkali and alkaline earth metals
Hydrogen forms with active metals hydrides:
Na + H 2 → 2 NaH Ca + H 2 → CaH 2 Mg + H 2 → MgH 2
Hydrides- salt-like, solid substances, easily hydrolyzed:
CaH 2 + 2H 2 O → Ca(OH) 2 + 2H 2
Interaction with metal oxides (usually d-elements)
Oxides are reduced to metals:
CuO + H 2 → Cu + H 2 O Fe 2 O 3 + 3H 2 → 2 Fe + 3H 2 O WO 3 + 3H 2 → W + 3H 2 O
Hydrogenation of organic compounds
When hydrogen acts on unsaturated hydrocarbons in the presence of a nickel catalyst and at elevated temperatures, a reaction occurs hydrogenation:
CH 2 =CH 2 + H 2 → CH 3 -CH 3
Hydrogen reduces aldehydes to alcohols:
CH 3 CHO + H 2 → C 2 H 5 OH.
Geochemistry of hydrogen
Hydrogen - basic construction material universe. It is the most common element, and all elements are formed from it as a result of thermonuclear and nuclear reactions.
Free hydrogen H2 is relatively rare in terrestrial gases, but in the form of water it takes an extremely important part in geochemical processes.
Hydrogen can be present in minerals in the form of ammonium ion, hydroxyl ion and crystalline water.
In the atmosphere, hydrogen is continuously produced as a result of the decomposition of water solar radiation. It migrates to the upper atmosphere and escapes into space.
Application
- Hydrogen energy
Atomic hydrogen is used for atomic hydrogen welding.
IN Food Industry hydrogen is registered as food additives E949, like packaging gas.
Features of treatment
Hydrogen, when mixed with air, forms an explosive mixture - the so-called detonating gas. This gas is most explosive when the volume ratio of hydrogen and oxygen is 2:1, or hydrogen and air is approximately 2:5, since air contains approximately 21% oxygen. Hydrogen is also a fire hazard. Liquid hydrogen can cause severe frostbite if it comes into contact with the skin.
Explosive concentrations of hydrogen and oxygen occur from 4% to 96% by volume. When mixed with air from 4% to 75(74)% by volume.
Hydrogen use
IN chemical industry Hydrogen is used in the production of ammonia, soap and plastics. In the food industry using hydrogen from liquids vegetable oils make margarine. Hydrogen is very light and always rises in the air. Once upon a time, airships and Balloons filled with hydrogen. But in the 30s. XX century several happened terrible disasters when airships exploded and burned. Nowadays, airships are filled with helium gas. Hydrogen is also used as rocket fuel. Someday, hydrogen may be widely used as a fuel for cars and trucks. Hydrogen engines do not pollute environment and only release water vapor (although the production of hydrogen itself leads to some environmental pollution). Our Sun is mostly made of hydrogen. Solar heat and light is the result of the release of nuclear energy from the fusion of hydrogen nuclei.
Using hydrogen as a fuel (cost-effective)
The most important characteristic of substances used as fuel is their heat of combustion. From the course general chemistry It is known that the reaction between hydrogen and oxygen occurs with the release of heat. If we take 1 mol H 2 (2 g) and 0.5 mol O 2 (16 g) under standard conditions and excite the reaction, then according to the equation
H 2 + 0.5 O 2 = H 2 O
after completion of the reaction, 1 mol of H 2 O (18 g) is formed with the release of energy 285.8 kJ/mol (for comparison: the heat of combustion of acetylene is 1300 kJ/mol, propane - 2200 kJ/mol). 1 m³ of hydrogen weighs 89.8 g (44.9 mol). Therefore, to produce 1 m³ of hydrogen, 12832.4 kJ of energy will be expended. Taking into account the fact that 1 kWh = 3600 kJ, we get 3.56 kWh of electricity. Knowing the tariff for 1 kWh of electricity and the cost of 1 m³ of gas, we can conclude that it is advisable to switch to hydrogen fuel.
For example, the 3rd generation Honda FCX experimental model with a 156 liter hydrogen tank (contains 3.12 kg of hydrogen under a pressure of 25 MPa) travels 355 km. Accordingly, from 3.12 kg H2, 123.8 kWh is obtained. Per 100 km, energy consumption will be 36.97 kWh. Knowing the cost of electricity, the cost of gas or gasoline, their consumption for a car per 100 km can easily be calculated as negative economic effect transition of cars to hydrogen fuel. Let's say (Russia 2008), 10 cents per kWh of electricity leads to the fact that 1 m³ of hydrogen leads to a price of 35.6 cents, and taking into account the efficiency of water decomposition of 40-45 cents, the same amount of kWh from burning gasoline costs 12832.4 kJ/42000 kJ/0.7 kg/l*80 cents/l=34 cents at retail prices, whereas for hydrogen we calculated perfect option, without taking into account transportation, depreciation of equipment, etc. For methane with a combustion energy of about 39 MJ per m³, the result will be two to four times lower due to the difference in price (1 m³ for Ukraine costs $179, and for Europe $350) . That is, an equivalent amount of methane will cost 10-20 cents.
However, we should not forget that when we burn hydrogen, we get clean water from which it was extracted. That is, we have a renewable hoarder energy without harm to the environment, unlike gas or gasoline, which are primary sources of energy.
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Oxygen is the most abundant element on Earth. Together with nitrogen and a small amount of other gases, free oxygen forms the Earth's atmosphere. Its content in the air is 20.95% by volume or 23.15% by mass. IN earth's crust 58% of the atoms are bound oxygen atoms (47% by mass). Oxygen is part of water (the reserves of bound oxygen in the hydrosphere are extremely large), rocks, many minerals and salts, are found in the fats, proteins and carbohydrates that make up living organisms. Almost all of the Earth's free oxygen is created and preserved as a result of the process of photosynthesis.
Physical properties.
Oxygen is a colorless, tasteless, and odorless gas, slightly heavier than air. It is slightly soluble in water (31 ml of oxygen dissolves in 1 liter of water at 20 degrees), but it is still better than other atmospheric gases, so water is enriched with oxygen. Oxygen density under normal conditions is 1.429 g/l. At a temperature of -183 0 C and a pressure of 101.325 kPa, oxygen turns into a liquid state. Liquid oxygen has a bluish color, is drawn into a magnetic field, and at -218.7 ° C, forms blue crystals.
Natural oxygen has three isotopes O 16, O 17, O 18.
Allotropy- the ability of a chemical element to exist in the form of two or more simple substances that differ only in the number of atoms in the molecule or in structure.
Ozone O 3 – exists in the upper layers of the atmosphere at an altitude of 20-25 km from the Earth’s surface and forms the so-called “ozone layer”, which protects the Earth from the harmful ultraviolet radiation of the Sun; pale purple, poisonous large quantities gas with a specific, pungent, but pleasant odor. The melting point is -192.7 0 C, the boiling point is 111.9 0 C. We dissolve oxygen better in water.
Ozone - strong oxidizing agent. Its oxidative activity is based on the ability of the molecule to decompose with the release of atomic oxygen:
It oxidizes many simple and complex substances. With some metals it forms ozonides, for example potassium ozonide:
K + O 3 = KO 3
Ozone is produced in special devices - ozonizers. In them, under the influence of an electric discharge, molecular oxygen is converted into ozone:
A similar reaction occurs under the influence of lightning discharges.
The use of ozone is due to its strong oxidizing properties: it is used for bleaching fabrics, disinfecting drinking water, in medicine as a disinfectant.
Inhaling ozone in large quantities is harmful: it irritates the mucous membranes of the eyes and respiratory organs.
Chemical properties.
IN chemical reactions with atoms of other elements (except fluorine), oxygen exhibits exclusively oxidizing properties
The most important chemical property is the ability to form oxides with almost all elements. At the same time, oxygen reacts directly with most substances, especially when heated.
As a result of these reactions, as a rule, oxides are formed, less often peroxides:
2Ca + O 2 = 2CaO
2Ba + O 2 = 2BaO
2Na + O 2 = Na 2 O 2
Oxygen does not interact directly with halogens, gold, and platinum; their oxides are obtained indirectly. When heated, sulfur, carbon, and phosphorus burn in oxygen.
The interaction of oxygen with nitrogen begins only at a temperature of 1200 0 C or at electrical discharge:
N 2 + O 2 = 2NO
With hydrogen, oxygen forms water:
2H 2 + O 2 = 2H 2 O
During this reaction, it is released significant amount warmth.
A mixture of two volumes of hydrogen with one volume of oxygen explodes when ignited; it is called detonating gas.
Many metals upon contact with atmospheric oxygen are subject to destruction - corrosion. Some metals under normal conditions are oxidized only from the surface (for example, aluminum, chromium). The resulting oxide film prevents further interaction.
4Al + 3O 2 = 2Al 2 O 3
Under certain conditions, complex substances also interact with oxygen. In this case, oxides are formed, and in some cases, oxides and simple substances.
CH 4 + 2O 2 = CO 2 + 2H 2 O
H 2 S + O 2 = 2SO 2 + 2H 2 O
4NН 3 +ЗО 2 =2N 2 +6Н 2 О
4CH 3 NH 2 + 9O 2 = 4CO 2 + 2N 2 + 10H 2 O
When interacting with complex substances, oxygen acts as an oxidizing agent. Its importance is based on the oxidative activity of oxygen. property - ability support combustion substances.
Oxygen also forms a compound with hydrogen - hydrogen peroxide H 2 O 2 - a colorless transparent liquid with a pungent astringent taste, highly soluble in water. IN chemically Hydrogen peroxide is a very interesting compound. Its low stability is characteristic: when standing, it slowly decomposes into water and oxygen:
H 2 O 2 = H 2 O + O 2
Light, heat, the presence of alkalis, and contact with oxidizing or reducing agents accelerate the decomposition process. The oxidation state of oxygen in hydrogen peroxide = - 1, i.e. has an intermediate value between the oxidation state of oxygen in water (-2) and in molecular oxygen (0), so hydrogen peroxide exhibits redox duality. Oxidative properties hydrogen peroxide are much more pronounced than reducing ones, and they manifest themselves in acidic, alkaline and neutral environments.
H 2 O 2 + 2KI + H 2 SO 4 = K 2 SO 4 + I 2 + 2H 2 O
Hydrogen occupies a special position in the Periodic Table chemical elements DI. Mendeleev. In terms of the number of valence electrons and the ability to form the hydration ion H + in solutions, it is similar to alkali metals, and should be placed in group I. Based on the number of electrons required to complete the outer electron shell, the value of the ionization energy, the ability to exhibit a negative oxidation state, and the small atomic radius, hydrogen should be placed in group VII of the periodic table. Thus, the placement of hydrogen in one group or another of the periodic table is largely arbitrary, but in most cases it is placed in group VII.
Electronic formula of hydrogen 1 s 1 . The only valence electron is directly in the range of action atomic nucleus. The simplicity of the electronic configuration of hydrogen does not mean that Chemical properties of this element are simple. In contrast, the chemistry of hydrogen differs in many ways from the chemistry of other elements. Hydrogen in its compounds is capable of exhibiting oxidation states of +1 and –1.
There are a large number of methods for producing hydrogen. In the laboratory it is obtained by reacting certain metals with acids, for example:
Hydrogen can be obtained by electrolysis of aqueous solutions of sulfuric acid or alkalis. In this case, the process of hydrogen evolution occurs at the cathode and oxygen at the anode.
In industry, hydrogen is produced mainly from natural and associated gases, fuel gasification products and coke oven gas.
Simple substance hydrogen (H 2) is a flammable gas, colorless and odorless. Boiling point –252.8 °C. Hydrogen is 14.5 times lighter than air and slightly soluble in water.
The hydrogen molecule is stable and has great strength. Due to the high dissociation energy (435 kJ/mol), the decomposition of H 2 molecules into atoms occurs to a noticeable extent only at temperatures above 2000 °C.
For hydrogen, positive and negative degree oxidation, therefore in chemical reactions hydrogen can exhibit both oxidizing and reducing properties. In cases where hydrogen acts as an oxidizing agent, it behaves like halogens, forming hydrides similar to halides ( hydrides name a group of chemical compounds of hydrogen with metals and elements less electronegative than it):
In terms of oxidative activity, hydrogen is significantly inferior to halogens. Therefore, only hydrides of alkali and alkaline earth metals exhibit ionic character. Ionic as well as complex hydrides, for example, are strong reducing agents. They are widely used in chemical syntheses.
In most reactions, hydrogen behaves as a reducing agent. Under normal conditions, hydrogen does not react with oxygen, but when ignited, the reaction occurs explosively:
A mixture of two volumes of hydrogen with one volume of oxygen is called detonating gas. During controlled combustion, release occurs large quantity heat, and the temperature of the hydrogen-oxygen flame reaches 3000 °C.
The reaction with halogens proceeds in different ways, depending on the nature of the halogen:
With fluorine, this reaction occurs explosively even at low temperatures. With chlorine in the light, the reaction also occurs explosively. With bromine the reaction is much slower, but with iodine it does not reach completion even at high temperatures. The mechanism of these reactions is radical.
At elevated temperatures, hydrogen interacts with elements of group VI - sulfur, selenium, tellurium, for example:
The reaction of hydrogen with nitrogen is very important. This reaction is reversible. To shift the equilibrium towards the formation of ammonia, use high blood pressure. In industry, this process is carried out at a temperature of 450–500 °C in the presence of various catalysts:
Hydrogen reduces many metals from oxides, for example:
This reaction is used to obtain some pure metals.
Hydrogenation reactions play a huge role organic compounds, which are widely used both in laboratory practice and in industrial organic synthesis.
Reduction natural sources hydrocarbon raw materials, environmental pollution by fuel combustion products are increasing interest in hydrogen as an environmentally friendly fuel. Hydrogen will probably play important role in the energy sector of the future.
Currently, hydrogen is widely used in industry for the synthesis of ammonia, methanol, hydrogenation of solid and liquid fuel, in organic synthesis, for welding and cutting metals, etc.
Water H 2 O, hydrogen oxide, is the most important chemical compound. Under normal conditions, water is a colorless liquid, odorless and tasteless. Water is the most abundant substance on the surface of the Earth. IN human body contains 63–68% water.
The physical properties of water are in many ways anomalous. Under normal conditions atmospheric pressure water boils at 100 °C. Freezing point clean water 0°C. Unlike other liquids, the density of water does not increase monotonically when cooled, but has a maximum at +4 °C. The heat capacity of water is very high and amounts to 418 kJ/mol·K. The heat capacity of ice at 0 °C is 2.038 kJ/mol·K. The heat of melting of ice is abnormally high. The electrical conductivity of water is very low. Abnormal physical properties waters explain its structure. The H–O–H bond angle is 104.5°. The water molecule is a distorted tetrahedron, at two vertices of which hydrogen atoms are located, and the other two are occupied by the orbitals of lone pairs of electrons of the oxygen atom, which are not involved in the formation of chemical bonds.
Water is a stable compound; its decomposition into oxygen and hydrogen occurs only under the influence of constant electric current or at a temperature of about 2000 °C:
Water directly interacts with metals in the range of standard electronic potentials up to hydrogen. Depending on the nature of the metal, the reaction products can be the corresponding hydroxides and oxides. The reaction rate, depending on the nature of the metal, also varies within wide limits. Thus, sodium reacts with water already at room temperature, the reaction is accompanied by the release of a large amount of heat; iron reacts with water at a temperature of 800 °C:
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