In the periodic table, hydrogen is located in two groups of elements that are completely opposite in their properties. This feature make it completely unique. Hydrogen is not just an element or substance, but is also integral part many complex compounds, organogenic and biogenic elements. Therefore, let's look at its properties and characteristics in more detail.
The release of flammable gas during the interaction of metals and acids was observed back in the 16th century, that is, during the formation of chemistry as a science. Famous English scientist Henry Cavendish studied the substance starting in 1766 and gave it the name “combustible air.” When burned, this gas produced water. Unfortunately, the scientist’s adherence to the theory of phlogiston (hypothetical “ultrafine matter”) prevented him from coming to the right conclusions.
The French chemist and naturalist A. Lavoisier, together with the engineer J. Meunier and with the help of special gasometers, synthesized water in 1783, and then analyzed it through the decomposition of water vapor with hot iron. Thus, scientists were able to come to the right conclusions. They found that “combustible air” is not only part of water, but can also be obtained from it.
In 1787, Lavoisier suggested that the gas under study was a simple substance and, accordingly, belonged to the number of primary chemical elements. He called it hydrogene (from the Greek words hydor - water + gennao - I give birth), i.e. “giving birth to water.”
The Russian name “hydrogen” was proposed in 1824 by the chemist M. Soloviev. The determination of the composition of water marked the end of the “phlogiston theory.” At the turn of the 18th and 19th centuries, it was established that the hydrogen atom is very light (compared to the atoms of other elements) and its mass was taken as the basic unit for comparing atomic masses, receiving a value equal to 1.
Physical properties
Hydrogen is the lightest substance known to science (it is 14.4 times lighter than air), its density is 0.0899 g/l (1 atm, 0 °C). This material melts (solidifies) and boils (liquefies), respectively, at -259.1 ° C and -252.8 ° C (only helium has lower boiling and melting temperatures).
The critical temperature of hydrogen is extremely low (-240 °C). For this reason, its liquefaction is a rather complex and costly process. Critical pressure substance - 12.8 kgf/cm², and the critical density is 0.0312 g/cm³. Among all gases, hydrogen has the highest thermal conductivity: at 1 atm and 0 °C it is equal to 0.174 W/(mxK).
The specific heat capacity of the substance under the same conditions is 14.208 kJ/(kgxK) or 3.394 cal/(rx°C). This element is slightly soluble in water (about 0.0182 ml/g at 1 atm and 20 °C), but well soluble in most metals (Ni, Pt, Pa and others), especially in palladium (about 850 volumes per volume of Pd ).
The latter property is associated with its ability to diffuse, and diffusion through a carbon alloy (for example, steel) can be accompanied by the destruction of the alloy due to the interaction of hydrogen with carbon (this process is called decarbonization). In the liquid state, the substance is very light (density - 0.0708 g/cm³ at t° = -253 °C) and fluid (viscosity - 13.8 spoise under the same conditions).
In many compounds, this element exhibits a +1 valency (oxidation state), like sodium and other alkali metals. It is usually considered as an analogue of these metals. Accordingly, he heads group I of the periodic system. In metal hydrides, the hydrogen ion exhibits a negative charge (the oxidation state is -1), that is, Na+H- has a structure similar to Na+Cl- chloride. In accordance with this and some other facts (proximity physical properties element “H” and halogens, the ability to replace it with halogens in organic compounds) Hydrogene belongs to group VII of the periodic system.
Under normal conditions, molecular hydrogen has low activity, directly combining only with the most active of non-metals (with fluorine and chlorine, with the latter in the light). In turn, when heated, it interacts with many chemical elements.
Atomic hydrogen has increased chemical activity (compared to molecular hydrogen). With oxygen it forms water according to the formula:
Н₂ + ½О₂ = Н₂О,
releasing 285.937 kJ/mol of heat or 68.3174 kcal/mol (25 °C, 1 atm). Under normal temperature conditions, the reaction proceeds rather slowly, and at t° >= 550 °C it is uncontrollable. The explosive limits of a hydrogen + oxygen mixture by volume are 4–94% H₂, and a hydrogen + air mixture is 4–74% H₂ (a mixture of two volumes of H₂ and one volume of O₂ is called detonating gas).
This element is used to reduce most metals, as it removes oxygen from oxides:
Fe₃O₄ + 4H₂ = 3Fe + 4H₂O,
CuO + H₂ = Cu + H₂O, etc.
Hydrogen forms hydrogen halides with different halogens, for example:
H₂ + Cl₂ = 2HCl.
However, when reacting with fluorine, hydrogen explodes (this also happens in the dark, at -252 ° C), with bromine and chlorine it reacts only when heated or illuminated, and with iodine - only when heated. When interacting with nitrogen, ammonia is formed, but only on a catalyst, at elevated pressures and temperatures:
ЗН₂ + N₂ = 2NN₃.
When heated, hydrogen reacts actively with sulfur:
H₂ + S = H₂S (hydrogen sulfide),
and much more difficult with tellurium or selenium. Hydrogen reacts with pure carbon without a catalyst, but at high temperatures:
2H₂ + C (amorphous) = CH₄ (methane).
This substance reacts directly with some of the metals (alkali, alkaline earth and others), forming hydrides, for example:
H₂ + 2Li = 2LiH.
Important practical significance have interactions between hydrogen and carbon(II) monoxide. In this case, depending on the pressure, temperature and catalyst, different organic compounds are formed: HCHO, CH₃OH, etc. Unsaturated hydrocarbons during the reaction become saturated, for example:
С n Н₂ n + Н₂ = С n Н₂ n ₊₂.
Hydrogen and its compounds play an exceptional role in chemistry. It determines the acidic properties of the so-called. protic acids, tends to form with different elements hydrogen bonding, which has a significant impact on the properties of many inorganic and organic compounds.
Hydrogen production
The main types of raw materials for industrial production This element includes oil refining gases, natural combustible and coke oven gases. It is also obtained from water through electrolysis (in places where electricity is available). One of the most important methods for producing material from natural gas The catalytic interaction of hydrocarbons, mainly methane, with water vapor (the so-called conversion) is considered. For example:
CH₄ + H₂O = CO + ZN₂.
Incomplete oxidation of hydrocarbons with oxygen:
CH₄ + ½O₂ = CO + 2H₂.
The synthesized carbon monoxide (II) undergoes conversion:
CO + H₂O = CO₂ + H₂.
Hydrogen produced from natural gas is the cheapest.
For the electrolysis of water, direct current is used, which is passed through a solution of NaOH or KOH (acids are not used to avoid corrosion of the equipment). In laboratory conditions, the material is obtained by electrolysis of water or as a result of a reaction between hydrochloric acid and zinc. However, ready-made factory material in cylinders is more often used.
This element is isolated from oil refining gases and coke oven gas by removing all other components of the gas mixture, since they liquefy more easily during deep cooling.
This material began to be produced industrially at the end of the 18th century. Then it was used for filling balloons. At the moment, hydrogen is widely used in industry, mainly in the chemical industry, for the production of ammonia.
Mass consumers of the substance are producers of methyl and other alcohols, synthetic gasoline and many other products. They are obtained by synthesis from carbon monoxide (II) and hydrogen. Hydrogene is used for hydrogenation of heavy and solid liquid fuel, fats, etc., for the synthesis of HCl, hydrotreating of petroleum products, as well as in cutting/welding of metals. The most important elements For nuclear energy are its isotopes - tritium and deuterium.
Biological role of hydrogen
About 10% of the mass of living organisms (on average) comes from this element. It is part of water and the most important groups of natural compounds, including proteins, nucleic acids, lipids, and carbohydrates. What is it used for?
This material plays a decisive role: in maintaining the spatial structure of proteins (quaternary), in implementing the principle of complementarity nucleic acids(i.e. in the implementation and storage of genetic information), in general in “recognition” at the molecular level.
The hydrogen ion H+ takes part in important dynamic reactions/processes in the body. Including: in biological oxidation, which provides living cells with energy, in biosynthetic reactions, in photosynthesis in plants, in bacterial photosynthesis and nitrogen fixation, in maintaining acid-base balance and homeostasis, in membrane transport processes. Along with carbon and oxygen, it forms the functional and structural basis of life phenomena.
In our Everyday life There are things that are so common that almost every person knows about them. For example, everyone knows that water is a liquid, it is easily accessible and does not burn, therefore, it can extinguish fire. But have you ever wondered why this is so?
Image source: pixabay.com
Water consists of hydrogen and oxygen atoms. Both of these elements support combustion. So, based on general logic (not scientific), it follows that water should also burn, right? However, this does not happen.
When does combustion occur?
Combustion is a chemical process in which molecules and atoms combine to release energy in the form of heat and light. To burn something you need two things - a fuel as a combustion source (for example, a sheet of paper, a piece of wood, etc.) and an oxidizer (oxygen contained in the earth's atmosphere is the main oxidizer). We also need the heat necessary to reach the ignition temperature of the substance in order for the combustion process to begin.
Image source auclip.ru
For example, consider the process of burning paper using matches. The paper in this case will be the fuel, the gaseous oxygen contained in the air will act as an oxidizing agent, and the ignition temperature will be achieved due to the burning match.
Structure of the chemical composition of water
Image source: water-service.com.ua
Water consists of two hydrogen atoms and one oxygen atom. Her chemical formula H2O. Now, it is interesting to note that the two constituents of water are indeed flammable substances.
Why is hydrogen a flammable substance?
Hydrogen atoms have only one electron and therefore easily combine with other elements. As a rule, hydrogen occurs in nature in the form of a gas whose molecules consist of two atoms. This gas is highly reactive and oxidizes quickly in the presence of an oxidizing agent, making it flammable.
Image source: myshared.ru
When hydrogen is burned, a large amount of energy is released, so it is often used in liquefied form to launch spacecraft into space.
Oxygen supports combustion
As mentioned earlier, any combustion requires an oxidizer. There are many chemical oxidizing agents, including oxygen, ozone, hydrogen peroxide, fluorine, etc. Oxygen is the main oxidizing agent found in abundance in the Earth's atmosphere. It is typically the primary oxidizing agent in most fires. That is why a constant supply of oxygen is necessary to maintain a fire.
Water puts out fire
Water can extinguish fire for a number of reasons, one of which is that it is a non-flammable liquid, despite being composed of two elements that could separately create a fiery inferno.
Water is the most common means of extinguishing fires. Image source: pixabay.com
As we said earlier, hydrogen is highly flammable, all it needs is an oxidizing agent and ignition temperature to start the reaction. Since oxygen is the most common oxidizing agent on Earth, it quickly combines with hydrogen atoms, releasing large amounts of light and heat, and water molecules are formed. Here's how it happens:
Please note that a mixture of hydrogen with a small amount of oxygen or air is explosive and is called detonating gas, it burns extremely quickly with a loud bang, which is perceived as an explosion. The Hindenburg airship disaster in New Jersey in 1937 claimed dozens of lives as a result of the ignition of hydrogen that filled the airship's shell. The easy flammability of hydrogen and its explosiveness in combination with oxygen is main reason the fact that we do not obtain water chemically in laboratories.
Hydrogen H is the most common element in the Universe (about 75% by mass), and on Earth it is the ninth most abundant. The most important natural hydrogen compound is water.
Hydrogen ranks first in the periodic table (Z = 1). It has the simplest atomic structure: the nucleus of the atom is 1 proton, surrounded by an electron cloud consisting of 1 electron.
Under some conditions, hydrogen exhibits metallic properties (donates an electron), while in others it exhibits nonmetallic properties (accepts an electron).
Hydrogen isotopes found in nature are: 1H - protium (the nucleus consists of one proton), 2H - deuterium (D - the nucleus consists of one proton and one neutron), 3H - tritium (T - the nucleus consists of one proton and two neutrons).
Simple substance hydrogen
A hydrogen molecule consists of two atoms connected by a covalent nonpolar bond.
Physical properties. Hydrogen is a colorless, odorless, tasteless, non-toxic gas. The hydrogen molecule is not polar. Therefore, the forces of intermolecular interaction in hydrogen gas are small. This manifests itself in low temperatures boiling (-252.6 0С) and melting (-259.2 0С).
Hydrogen is lighter than air, D (by air) = 0.069; slightly soluble in water (2 volumes of H2 dissolve in 100 volumes of H2O). Therefore, hydrogen, when produced in the laboratory, can be collected by air or water displacement methods.
Hydrogen production
In the laboratory:
1. Effect of dilute acids on metals:
Zn +2HCl → ZnCl 2 +H 2
2. Interaction between alkaline and metals with water:
Ca +2H 2 O → Ca(OH) 2 +H 2
3. Hydrolysis of hydrides: metal hydrides are easily decomposed by water to form the corresponding alkali and hydrogen:
NaH +H 2 O → NaOH +H 2
CaH 2 + 2H 2 O = Ca(OH) 2 + 2H 2
4.The effect of alkalis on zinc or aluminum or silicon:
2Al +2NaOH +6H 2 O → 2Na +3H 2
Zn +2KOH +2H 2 O → K 2 +H 2
Si + 2NaOH + H 2 O → Na 2 SiO 3 + 2H 2
5. Electrolysis of water. To increase the electrical conductivity of water, an electrolyte is added to it, for example NaOH, H 2 SO 4 or Na 2 SO 4. 2 volumes of hydrogen are formed at the cathode, and 1 volume of oxygen at the anode.
2H 2 O → 2H 2 +O 2
Industrial production of hydrogen
1. Methane conversion with steam, Ni 800 °C (cheapest):
CH 4 + H 2 O → CO + 3 H 2
CO + H 2 O → CO 2 + H 2
In total:
CH 4 + 2 H 2 O → 4 H 2 + CO 2
2. Water vapor through hot coke at 1000 o C:
C + H 2 O → CO + H 2
CO +H 2 O → CO 2 + H 2
The resulting carbon monoxide (IV) is absorbed by water, and 50% of industrial hydrogen is produced in this way.
3. By heating methane to 350°C in the presence of an iron or nickel catalyst:
CH 4 → C + 2H 2
4. Electrolysis of aqueous solutions of KCl or NaCl as a by-product:
2H 2 O + 2NaCl → Cl 2 + H 2 + 2NaOH
Chemical properties of hydrogen
- In compounds, hydrogen is always monovalent. It is characterized by an oxidation state of +1, but in metal hydrides it is equal to -1.
- The hydrogen molecule consists of two atoms. The emergence of a connection between them is explained by the formation of a generalized pair of electrons H:H or H 2
- Thanks to this generalization of electrons, the H 2 molecule is more energetically stable than its individual atoms. To break 1 mole of hydrogen molecules into atoms, it is necessary to expend 436 kJ of energy: H 2 = 2H, ∆H° = 436 kJ/mol
- This explains the relatively low activity of molecular hydrogen at ordinary temperatures.
- With many non-metals, hydrogen forms gaseous compounds such as RH 4, RH 3, RH 2, RH.
1) Forms hydrogen halides with halogens:
H 2 + Cl 2 → 2HCl.
At the same time, it explodes with fluorine, reacts with chlorine and bromine only when illuminated or heated, and with iodine only when heated.
2) With oxygen:
2H 2 + O 2 → 2H 2 O
with heat release. At normal temperatures the reaction proceeds slowly, above 550°C it explodes. A mixture of 2 volumes of H 2 and 1 volume of O 2 is called detonating gas.
3) When heated, it reacts vigorously with sulfur (much more difficult with selenium and tellurium):
H 2 + S → H 2 S (hydrogen sulfide),
4) With nitrogen with the formation of ammonia only on a catalyst and at elevated temperatures and pressures:
ZN 2 + N 2 → 2NH 3
5) With carbon at high temperatures:
2H 2 + C → CH 4 (methane)
6) Forms hydrides with alkali and alkaline earth metals (hydrogen is an oxidizing agent):
H 2 + 2Li → 2LiH
in metal hydrides, the hydrogen ion is negatively charged (oxidation state -1), that is, Na + H hydride - built similar to Na + Cl chloride -
7) With metal oxides (used to reduce metals):
CuO + H 2 → Cu + H 2 O
Fe 3 O 4 + 4H 2 → 3Fe + 4H 2 O
8) with carbon monoxide (II):
CO + 2H 2 → CH 3 OH
Synthesis - gas (a mixture of hydrogen and carbon monoxide) is of important practical importance, because depending on temperature, pressure and catalyst, various organic compounds are formed, for example HCHO, CH 3 OH and others.
9) Unsaturated hydrocarbons react with hydrogen, becoming saturated:
C n H 2n + H 2 → C n H 2n+2.
§3. Reaction equation and how to write it
Interaction hydrogen With oxygen, as Sir Henry Cavendish established, leads to the formation of water. Let's get on with it simple example let's learn how to compose chemical reaction equations.
What comes out of hydrogen And oxygen, we already know:
H 2 + O 2 → H 2 O
Now let us take into account that atoms of chemical elements in chemical reactions do not disappear and do not appear from nothing, do not transform into each other, but combine in new combinations, forming new molecules. So in the equation chemical reaction there must be the same number of atoms of each type before reactions ( left from the equal sign) and after the end of the reaction ( on right from the equal sign), like this:
2H 2 + O 2 = 2H 2 O
That's what it is reaction equation - conditional recording of an ongoing chemical reaction using formulas of substances and coefficients.
This means that in the given reaction two moles hydrogen must react with one mole oxygen, and the result will be two moles water.
Interaction hydrogen With oxygen- not a simple process at all. It leads to a change in the oxidation states of these elements. To select coefficients in such equations, they usually use the " electronic balance".
When water is formed from hydrogen and oxygen, it means that hydrogen changed its oxidation state from 0 before +I, A oxygen- from 0 before −II. In this case, several passed from hydrogen atoms to oxygen atoms. (n) electrons:
Hydrogen donating electrons serves here reducing agent, and oxygen accepting electrons is oxidizing agent.
Oxidizing agents and reducing agents
Let's now see what the processes of giving and receiving electrons look like separately. Hydrogen, having met with the “robber” oxygen, loses all its assets - two electrons, and its oxidation state becomes equal +I:
N 2 0 − 2 e− = 2Н +I
Happened oxidation half-reaction equation hydrogen.
And the bandit- oxygen O 2, having taken the last electrons from the unfortunate hydrogen, is very pleased with his new oxidation state -II:
O2+4 e− = 2O −II
This reduction half-reaction equation oxygen.
It remains to add that both the “bandit” and his “victim” have lost their chemical individuality and are made from simple substances - gases with diatomic molecules H 2 And O 2 turned into components of a new chemical substance - water H 2 O.
Further we will reason as follows: how many electrons the reducing agent gave to the oxidizing bandit, that’s how many electrons he received. The number of electrons donated by the reducing agent must be equal to the number of electrons accepted by the oxidizing agent.
So it's necessary equalize the number of electrons in the first and second half-reactions. In chemistry, the following conventional form of writing half-reaction equations is accepted:
2 N 2 0 − 2 e− = 2Н +I |
|
1 O 2 0 + 4 e− = 2O −II |
Here, the numbers 2 and 1 to the left of the curly brace are factors that will help ensure that the number of electrons given and received is equal. Let's take into account that in the half-reaction equations 2 electrons are given, and 4 are accepted. To equalize the number of accepted and given electrons, find the least common multiple and additional factors. In our case, the least common multiple is 4. The additional factors for hydrogen will be 2 (4: 2 = 2), and for oxygen - 1 (4: 4 = 1)
The resulting multipliers will serve as the coefficients of the future reaction equation:
2H 2 0 + O 2 0 = 2H 2 +I O −II
Hydrogen oxidizes not only when meeting with oxygen. They act on hydrogen in approximately the same way. fluorine F 2, a halogen and a known "robber", and seemingly harmless nitrogen N 2:
H 2 0 + F 2 0 = 2H +I F −I |
3H 2 0 + N 2 0 = 2N −III H 3 +I |
In this case it turns out hydrogen fluoride HF or ammonia NH 3.
In both compounds the oxidation state is hydrogen becomes equal +I, because he gets molecule partners who are “greedy” for other people’s electronic goods, with high electronegativity - fluorine F And nitrogen N. U nitrogen the value of electronegativity is considered equal to three conventional units, and fluoride In general, the highest electronegativity among all chemical elements is four units. So it’s no wonder they left the poor hydrogen atom without any electronic environment.
But hydrogen maybe restore- accept electrons. This happens if alkali metals or calcium, which have a lower electronegativity than hydrogen, participate in the reaction with it.
10.1.Hydrogen
The name "hydrogen" refers to both a chemical element and a simple substance. Element hydrogen consists of hydrogen atoms. Simple substance hydrogen consists of hydrogen molecules.
A) Chemical element hydrogen
In the natural series of elements, the serial number of hydrogen is 1. In the system of elements, hydrogen is in the first period in group IA or VIIA.
Hydrogen is one of the most common elements on Earth. The mole fraction of hydrogen atoms in the atmosphere, hydrosphere and lithosphere of the Earth (collectively called the earth's crust) is 0.17. It is found in water, many minerals, oil, natural gas, plants and animals. The average human body contains about 7 kilograms of hydrogen.
There are three isotopes of hydrogen:
a) light hydrogen – protium,
b) heavy hydrogen – deuterium(D),
c) superheavy hydrogen – tritium(T).
Tritium is an unstable (radioactive) isotope, so it is practically never found in nature. Deuterium is stable, but there is very little of it: w D = 0.015% (of the mass of all terrestrial hydrogen). Therefore, the atomic mass of hydrogen differs very little from 1 Dn (1.00794 Dn).
b) Hydrogen atom
From previous sections of the chemistry course, you already know the following characteristics of the hydrogen atom:
The valence capabilities of a hydrogen atom are determined by the presence of one electron in a single valence orbital. A high ionization energy makes a hydrogen atom not inclined to give up an electron, and a not too high electron affinity energy leads to a slight tendency to accept one. Consequently, in chemical systems the formation of the H cation is impossible, and compounds with the H anion are not very stable. Thus, the hydrogen atom is most likely to form a covalent bond with other atoms due to its one unpaired electron. Both in the case of the formation of an anion and in the case of the formation of a covalent bond, the hydrogen atom is monovalent.
In a simple substance, the oxidation state of hydrogen atoms is zero; in most compounds, hydrogen exhibits an oxidation state of +I, and only in the hydrides of the least electronegative elements does hydrogen have an oxidation state of –I.
Information about the valence capabilities of the hydrogen atom is given in Table 28. The valence state of a hydrogen atom bound by one covalent bond to any atom is indicated in the table by the symbol “H-”.
Table 28.Valence possibilities of the hydrogen atom
Valence state |
Examples of chemicals |
|||
I |
HCl, H 2 O, H 2 S, NH 3, CH 4, C 2 H 6, NH 4 Cl, H 2 SO 4, NaHCO 3, KOH |
|||
NaH, KH, CaH 2, BaH 2 |
c) Hydrogen molecule
The diatomic hydrogen molecule H2 is formed when hydrogen atoms are bonded with the only covalent bond possible for them. The connection is formed by an exchange mechanism. According to the way electron clouds overlap, this is an s-bond (Fig. 10.1 A). Since the atoms are the same, the bond is non-polar.
Interatomic distance (more precisely, equilibrium interatomic distance, because atoms vibrate) in a hydrogen molecule r(H–H) = 0.74 A (Fig. 10.1 V), which is significantly less than the sum of the orbital radii (1.06 A). Consequently, the electron clouds of bonded atoms overlap deeply (Fig. 10.1 b), and the bond in the hydrogen molecule is strong. This is pretty much the same thing great importance binding energy (454 kJ/mol).
If we characterize the shape of the molecule by the boundary surface (similar to the boundary surface of the electron cloud), then we can say that the hydrogen molecule has the shape of a slightly deformed (elongated) ball (Fig. 10.1 G).
d) Hydrogen (substance)
Under normal conditions, hydrogen is a colorless and odorless gas. In small quantities it is non-toxic. Solid hydrogen melts at 14 K (–259 °C), and liquid hydrogen boils at 20 K (–253 °C). Low melting and boiling points, a very small temperature range for the existence of liquid hydrogen (only 6 °C), as well as small values of the molar heats of fusion (0.117 kJ/mol) and vaporization (0.903 kJ/mol) indicate that intermolecular bonds in hydrogen very weak.
Hydrogen density r(H 2) = (2 g/mol): (22.4 l/mol) = 0.0893 g/l. For comparison: the average air density is 1.29 g/l. That is, hydrogen is 14.5 times “lighter” than air. It is practically insoluble in water.
At room temperature, hydrogen is inactive, but when heated it reacts with many substances. In these reactions, hydrogen atoms can either increase or decrease their oxidation state: H 2 + 2 e– = 2Н –I, Н 2 – 2 e– = 2Н +I.
In the first case, hydrogen is an oxidizing agent, for example, in reactions with sodium or calcium: 2Na + H 2 = 2NaH, ( t) Ca + H 2 = CaH 2 . ( t)
But the reducing properties of hydrogen are more characteristic: O 2 + 2H 2 = 2H 2 O, ( t)
CuO + H 2 = Cu + H 2 O. ( t)
When heated, hydrogen is oxidized not only by oxygen, but also by some other non-metals, for example, fluorine, chlorine, sulfur and even nitrogen.
In the laboratory, hydrogen is produced as a result of the reaction
Zn + H 2 SO 4 = ZnSO 4 + H 2.
Instead of zinc, you can use iron, aluminum and some other metals, and instead of sulfuric acid, you can use some other dilute acids. The resulting hydrogen is collected in a test tube by displacing water (see Fig. 10.2 b) or simply into an inverted flask (Fig. 10.2 A).
In industry, hydrogen is produced in large quantities from natural gas (mainly methane) by reacting it with water vapor at 800 °C in the presence of a nickel catalyst:
CH 4 + 2H 2 O = 4H 2 +CO 2 ( t, Ni)
or treat coal at high temperature with water vapor:
2H 2 O + C = 2H 2 + CO 2. ( t)
Pure hydrogen is obtained from water by decomposing it with electric current (subjecting to electrolysis):
2H 2 O = 2H 2 + O 2 (electrolysis).
e) Hydrogen compounds
Hydrides (binary compounds containing hydrogen) are divided into two main types:
a) volatile
(molecular) hydrides,
b) salt-like (ionic) hydrides.
Elements of groups IVA – VIIA and boron form molecular hydrides. Of these, only the hydrides of elements forming nonmetals are stable:
B 2 H 6 ; CH 4 ; NH3; H2O; HF
SiH 4 ;PH 3 ; H2S; HCl
AsH3; H2Se; HBr
H2Te; HI
With the exception of water, all these compounds are gaseous substances at room temperature, hence their name - “volatile hydrides”.
Some of the elements that form nonmetals are also found in more complex hydrides. For example, carbon forms compounds with the general formulas C n H 2 n+2 , C n H 2 n, C n H 2 n–2 and others, where n can be very large (these compounds are studied in organic chemistry).
Ionic hydrides include hydrides of alkali, alkaline earth elements and magnesium. The crystals of these hydrides consist of H anions and metal cations in the highest oxidation state Me or Me 2 (depending on the group of the element system).
| LiH | |
| NaH | MgH 2 |
| KH | CaH2 |
| RbH | SrH 2 |
| CsH | BaH 2 |
Both ionic and almost all molecular hydrides (except H 2 O and HF) are reducing agents, but ionic hydrides exhibit reducing properties much stronger than molecular ones.
In addition to hydrides, hydrogen is part of hydroxides and some salts. You will become familiar with the properties of these more complex hydrogen compounds in the following chapters.
The main consumers of hydrogen produced in industry are plants for the production of ammonia and nitrogen fertilizers, where ammonia is obtained directly from nitrogen and hydrogen:
N 2 +3H 2 2NH 3 ( R, t, Pt – catalyst).
Hydrogen is used in large quantities to produce methyl alcohol (methanol) by the reaction 2H 2 + CO = CH 3 OH ( t, ZnO – catalyst), as well as in the production of hydrogen chloride, which is obtained directly from chlorine and hydrogen:
H 2 + Cl 2 = 2HCl.
Sometimes hydrogen is used in metallurgy as a reducing agent in the production of pure metals, for example: Fe 2 O 3 + 3H 2 = 2Fe + 3H 2 O.
1. What particles do the nuclei of a) protium, b) deuterium, c) tritium consist of?
2.Compare the ionization energy of the hydrogen atom with the ionization energy of atoms of other elements. Which element is hydrogen closest to in terms of this characteristic?
3.Do the same for electron affinity energy
4. Compare the direction of polarization of the covalent bond and the degree of oxidation of hydrogen in the compounds: a) BeH 2, CH 4, NH 3, H 2 O, HF; b) CH 4, SiH 4, GeH 4.
5.Write down the simplest, molecular, structural and spatial formula of hydrogen. Which one is most often used?
6. They often say: “Hydrogen is lighter than air.” What does this mean? In what cases can this expression be taken literally, and in what cases can it not?
7.Make up the structural formulas of potassium and calcium hydrides, as well as ammonia, hydrogen sulfide and hydrogen bromide.
8.Knowing the molar heats of melting and vaporization of hydrogen, determine the values of the corresponding specific quantities.
9.For each of the four reactions illustrating the main Chemical properties hydrogen, create an electron balance. Label the oxidizing and reducing agents.
10. Determine the mass of zinc required to produce 4.48 liters of hydrogen using a laboratory method.
11. Determine the mass and volume of hydrogen that can be obtained from 30 m 3 of a mixture of methane and water vapor, taken in a volume ratio of 1:2, with a yield of 80%.
12. Make up equations for the reactions that occur during the interaction of hydrogen a) with fluorine, b) with sulfur.
13. The reaction schemes below illustrate the basic chemical properties of ionic hydrides:
a) MH + O 2 MOH ( t); b) MH + Cl 2 MCl + HCl ( t);
c) MH + H 2 O MOH + H 2 ; d) MH + HCl(p) MCl + H 2
Here M is lithium, sodium, potassium, rubidium or cesium. Write down the equations for the corresponding reactions if M is sodium. Illustrate the chemical properties of calcium hydride using reaction equations.
14.Using the electron balance method, create equations for the following reactions illustrating the reducing properties of some molecular hydrides:
a) HI + Cl 2 HCl + I 2 ( t); b) NH 3 + O 2 H 2 O + N 2 ( t); c) CH 4 + O 2 H 2 O + CO 2 ( t).
10.2 Oxygen
As with hydrogen, the word "oxygen" is the name of both a chemical element and simple substance. Apart from simple matter" oxygen"(dioxygen) chemical element oxygen forms another simple substance called " ozone"(trioxygen). This allotropic modifications oxygen. The substance oxygen consists of oxygen molecules O 2 , and the substance ozone consists of ozone molecules O 3 .
a) Chemical element oxygen
In the natural series of elements, the serial number of oxygen is 8. In the system of elements, oxygen is in the second period in the VIA group.
Oxygen is the most abundant element on Earth. In the earth's crust, every second atom is an oxygen atom, that is, the molar fraction of oxygen in the atmosphere, hydrosphere and lithosphere of the Earth is about 50%. Oxygen (substance) – component air. The volume fraction of oxygen in the air is 21%. Oxygen (an element) is found in water, many minerals, and plants and animals. The human body contains on average 43 kg of oxygen.
Natural oxygen consists of three isotopes (16 O, 17 O and 18 O), of which the lightest isotope 16 O is the most common. Therefore, the atomic mass of oxygen is close to 16 Dn (15.9994 Dn).
b) Oxygen atom
You know the following characteristics of the oxygen atom.
Table 29.Valence possibilities of the oxygen atom
Valence state |
Examples of chemicals |
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Al 2 O 3 , Fe 2 O 3 , Cr 2 O 3 * |
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–II |
H 2 O, SO 2, SO 3, CO 2, SiO 2, H 2 SO 4, HNO 2, HClO 4, COCl 2, H 2 O 2 |
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NaOH, KOH, Ca(OH) 2, Ba(OH) 2 |
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Li 2 O, Na 2 O, MgO, CaO, BaO, FeO, La 2 O 3 |
* These oxides can also be considered as ionic compounds.
** The oxygen atoms in the molecule are not in this valence state; this is just an example of a substance with an oxidation state of oxygen atoms equal to zero
The high ionization energy (like that of hydrogen) prevents the formation of a simple cation from the oxygen atom. The electron affinity energy is quite high (almost twice that of hydrogen), which provides a greater propensity for the oxygen atom to gain electrons and the ability to form O 2A anions. But the electron affinity energy of the oxygen atom is still lower than that of halogen atoms and even other elements of the VIA group. Therefore, oxygen anions ( oxide ions) exist only in compounds of oxygen with elements whose atoms give up electrons very easily.
By sharing two unpaired electrons, an oxygen atom can form two covalent bonds. Two lone pairs of electrons, due to the impossibility of excitation, can only enter into donor-acceptor interaction. Thus, without taking into account the bond multiplicity and hybridization, the oxygen atom can be in one of five valence states (Table 29).
The most typical valence state for the oxygen atom is W k = 2, that is, the formation of two covalent bonds due to two unpaired electrons.
The very high electronegativity of the oxygen atom (higher only for fluorine) leads to the fact that in most of its compounds oxygen has an oxidation state of –II. There are substances in which oxygen exhibits other oxidation states, some of them are given in Table 29 as examples, and the comparative stability is shown in Fig. 10.3.
c) Oxygen molecule
It has been experimentally established that the diatomic oxygen molecule O 2 contains two unpaired electrons. Using the valence bond method, this electronic structure of this molecule cannot be explained. However, the bond in the oxygen molecule is close in properties to a covalent one. The oxygen molecule is non-polar. Interatomic distance ( r o–o = 1.21 A = 121 nm) is less than the distance between atoms connected by a single bond. The molar binding energy is quite high and amounts to 498 kJ/mol.
d) Oxygen (substance)
Under normal conditions, oxygen is a colorless and odorless gas. Solid oxygen melts at 55 K (–218 °C), and liquid oxygen boils at 90 K (–183 °C).
Intermolecular bonds in solid and liquid oxygen are somewhat stronger than in hydrogen, as evidenced by the larger temperature range of existence of liquid oxygen (36 °C) and larger molar heats of fusion (0.446 kJ/mol) and vaporization (6. 83 kJ/mol).
Oxygen is slightly soluble in water: at 0 °C, only 5 volumes of oxygen (gas!) dissolve in 100 volumes of water (liquid!).
The high propensity of oxygen atoms to gain electrons and high electronegativity lead to the fact that oxygen exhibits only oxidizing properties. These properties are especially pronounced at high temperatures.
Oxygen reacts with many metals: 2Ca + O 2 = 2CaO, 3Fe + 2O 2 = Fe 3 O 4 ( t);
non-metals: C + O 2 = CO 2, P 4 + 5O 2 = P 4 O 10,
and complex substances: CH 4 + 2O 2 = CO 2 + 2H 2 O, 2H 2 S + 3O 2 = 2H 2 O + 2SO 2.
Most often, as a result of such reactions, various oxides are obtained (see Chapter II § 5), but active alkali metals, for example sodium, when burned, turn into peroxides:
2Na + O 2 = Na 2 O 2.
The structural formula of the resulting sodium peroxide is (Na) 2 (O-O).
A smoldering splinter placed in oxygen bursts into flames. This is a convenient and easy way to detect pure oxygen.
In industry, oxygen is obtained from air by rectification (complex distillation), and in the laboratory - by subjecting certain oxygen-containing compounds to thermal decomposition, for example:
2KMnO 4 = K 2 MnO 4 + MnO 2 + O 2 (200 °C);
2KClO 3 = 2KCl + 3O 2 (150 °C, MnO 2 – catalyst);
2KNO 3 = 2KNO 2 + 3O 2 (400 °C)
and, in addition, by the catalytic decomposition of hydrogen peroxide at room temperature: 2H 2 O 2 = 2H 2 O + O 2 (MnO 2 catalyst).
Pure oxygen is used in industry to intensify those processes in which oxidation occurs and to create a high-temperature flame. In rocket technology, liquid oxygen is used as an oxidizer.
Oxygen is of great importance for maintaining the life of plants, animals and humans. Under normal conditions, a person has enough oxygen in the air to breathe. But in conditions where there is not enough air, or there is no air at all (in airplanes, during diving work, in spaceships etc.), special gas mixtures containing oxygen are prepared for breathing. Oxygen is also used in medicine for diseases that cause difficulty breathing.
e) Ozone and its molecules
Ozone O 3 is the second allotropic modification of oxygen.
The triatomic ozone molecule has a corner structure intermediate between the two structures represented by the following formulas:
Ozone is a dark blue gas with a pungent odor. Due to its strong oxidizing activity, it is poisonous. Ozone is one and a half times “heavier” than oxygen and slightly more soluble in water than oxygen.
Ozone is formed in the atmosphere from oxygen during lightning electrical discharges:
3O 2 = 2O 3 ().
At normal temperatures, ozone slowly turns into oxygen, and when heated, this process occurs explosively.
Ozone is contained in the so-called "ozone layer" earth's atmosphere, protecting all life on Earth from harmful effects solar radiation.
In some cities, ozone is used instead of chlorine to disinfect (disinfect) drinking water.
Draw the structural formulas of the following substances: OF 2, H 2 O, H 2 O 2, H 3 PO 4, (H 3 O) 2 SO 4, BaO, BaO 2, Ba(OH) 2. Name these substances. Describe the valence states of oxygen atoms in these compounds.
Determine the valence and oxidation state of each oxygen atom.
2. Make up equations for the combustion reactions of lithium, magnesium, aluminum, silicon, red phosphorus and selenium in oxygen (selenium atoms are oxidized to the oxidation state +IV, atoms of other elements are oxidized to the highest oxidation state). What classes of oxides do the products of these reactions belong to?
3. How many liters of ozone can be obtained (under normal conditions) a) from 9 liters of oxygen, b) from 8 g of oxygen?
Water is the most abundant substance in the earth's crust. The mass of earth's water is estimated at 10 18 tons. Water is the basis of the hydrosphere of our planet; in addition, it is contained in the atmosphere, in the form of ice it forms the Earth’s polar caps and high-mountain glaciers, and is also part of various rocks. The mass fraction of water in the human body is about 70%.
Water is the only substance that has its own special names in all three states of aggregation.
Electronic structure of a water molecule (Fig. 10.4 A) we studied in detail earlier (see § 7.10).
Due to the polarity of the O–H bonds and the angular shape, the water molecule is electric dipole.
To characterize the polarity of an electric dipole, a physical quantity called " electric moment of an electric dipole" or simply " dipole moment".
In chemistry, the dipole moment is measured in debyes: 1 D = 3.34. 10 –30 Class. m
In a water molecule there are two polar covalent bonds, that is, two electric dipoles, each of which has its own dipole moment (u). The total dipole moment of a molecule is equal to the vector sum of these two moments (Fig. 10.5):
(H 2 O) = 
,
Where q 1 and q 2 – partial charges (+) on hydrogen atoms, and and – interatomic O – H distances in the molecule. Because q 1 = q 2 = q, and , then
The experimentally determined dipole moments of the water molecule and some other molecules are given in the table.
Table 30.Dipole moments of some polar molecules
Molecule |
Molecule |
Molecule |
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Given the dipole nature of the water molecule, it is often schematically represented as follows:
Pure water is a colorless liquid without taste or smell. Some basic physical characteristics of water are given in the table.
Table 31.Some physical characteristics of water
The large values of the molar heats of melting and vaporization (an order of magnitude greater than those of hydrogen and oxygen) indicate that water molecules, both in solid and liquid matter, are quite tightly bound together. These connections are called " hydrogen bonds".
ELECTRIC DIPOLE, DIPOLE MOMENT, BOND POLARITY, MOLECULE POLARITY.
How many valence electrons of an oxygen atom take part in the formation of bonds in a water molecule?
2. When what orbitals overlap, bonds are formed between hydrogen and oxygen in a water molecule?
3.Make a diagram of the formation of bonds in a molecule of hydrogen peroxide H 2 O 2. What can you say about the spatial structure of this molecule?
4. Interatomic distances in HF, HCl and HBr molecules are equal to 0.92, respectively; 1.28 and 1.41. Using the table of dipole moments, calculate and compare the partial charges on the hydrogen atoms in these molecules.
5. The interatomic distances S – H in the hydrogen sulfide molecule are 1.34, and the angle between the bonds is 92°. Determine the values of the partial charges on the sulfur and hydrogen atoms. What can you say about the hybridization of the valence orbitals of the sulfur atom?
10.4. Hydrogen bond
As you already know, due to the significant difference in electronegativity of hydrogen and oxygen (2.10 and 3.50), the hydrogen atom in the water molecule acquires a large positive partial charge ( q h = 0.33 e), and the oxygen atom has an even greater negative partial charge ( q h = –0.66 e). Recall also that the oxygen atom has two lone pairs of electrons per sp 3-hybrid AO. The hydrogen atom of one water molecule is attracted to the oxygen atom of another molecule, and, in addition, the half-empty 1s-AO of the hydrogen atom partially accepts a pair of electrons of the oxygen atom. As a result of these interactions between molecules, a special kind intermolecular bonds - hydrogen bond.
In the case of water, hydrogen bond formation can be represented schematically as follows:
In the last structural formula, three dots (dotted line, not electrons!) indicate a hydrogen bond.
Hydrogen bonds exist not only between water molecules. It is formed if two conditions are met:
1) the molecule has a highly polar H–E bond (E is the symbol of an atom of a fairly electronegative element),
2) the molecule contains an E atom with a large negative partial charge and a lone pair of electrons.
The element E can be fluorine, oxygen and nitrogen. Hydrogen bonds are significantly weaker if E is chlorine or sulfur.
Examples of substances with hydrogen bonds between molecules: hydrogen fluoride, solid or liquid ammonia, ethyl alcohol and many others.
In liquid hydrogen fluoride, its molecules are linked by hydrogen bonds into fairly long chains, and in liquid and solid ammonia three-dimensional networks are formed.
The strength of the hydrogen bond is intermediate between chemical bond and other types of intermolecular bonds. The molar energy of a hydrogen bond usually ranges from 5 to 50 kJ/mol.
In solid water (i.e., ice crystals), all hydrogen atoms are hydrogen bonded to oxygen atoms, with each oxygen atom forming two hydrogen bonds (using both lone pairs of electrons). This structure makes ice more “loose” compared to liquid water, where some of the hydrogen bonds are broken, and the molecules are able to “pack” a little more tightly. This feature of the structure of ice explains why, unlike most other substances, water in the solid state has a lower density than in the liquid state. Water reaches its maximum density at 4 °C - at this temperature quite a lot of hydrogen bonds are broken, and thermal expansion does not yet have a very strong effect on the density.
Hydrogen bonds are very important in our lives. Let's imagine for a moment that hydrogen bonds have stopped forming. Here are some consequences:
- water at room temperature would become gaseous as its boiling point would drop to about -80 °C;
- all bodies of water would begin to freeze from the bottom, since the density of ice would be greater than the density of liquid water;
- The double helix of DNA and much more would cease to exist.
The examples given are enough to understand that in this case nature on our planet would become completely different.
HYDROGEN BOND, CONDITIONS OF ITS FORMATION.
The formula of ethyl alcohol is CH 3 – CH 2 – O – H. Between which atoms of different molecules of this substance are hydrogen bonds formed? Write structural formulas illustrating their formation.
2. Hydrogen bonds exist not only in individual substances, but also in solutions. Show with structural formulas How are hydrogen bonds formed in aqueous solution a) ammonia, b) hydrogen fluoride, c) ethanol (ethyl alcohol). = 2H 2 O.
Both of these reactions occur in water constantly and at the same speed, therefore, there is an equilibrium in water: 2H 2 O AN 3 O + OH.
This equilibrium is called equilibrium of autoprotolysis water.
The direct reaction of this reversible process is endothermic, therefore, when heated, autoprotolysis increases, but at room temperature the equilibrium is shifted to the left, that is, the concentration of H 3 O and OH ions is negligible. What are they equal to?
According to the law of mass action
But due to the fact that the number of reacted water molecules is insignificant compared to the total number of water molecules, we can assume that the concentration of water during autoprotolysis practically does not change, and 2 = const Such a low concentration of oppositely charged ions in clean water explains why this liquid, although poorly, still conducts electric current.
AUTOPROTOLYSIS OF WATER, AUTOPROTOLYSIS CONSTANT (IONIC PRODUCT) OF WATER.
The ionic product of liquid ammonia (boiling point –33 °C) is 2·10 –28. Write an equation for the autoprotolysis of ammonia. Determine the concentration of ammonium ions in pure liquid ammonia. Which substance has greater electrical conductivity, water or liquid ammonia?
1. Production of hydrogen and its combustion (reducing properties).
2. Obtaining oxygen and burning substances in it (oxidizing properties).
