Studying the phenomenon of radioactivity, scientists in the first decade of the 20th century. discovered a large number of radioactive substances - about 40. There were significantly more of them than there were free places in the periodic table of elements between bismuth and uranium. The nature of these substances has been controversial. Some researchers considered them to be independent chemical elements, but in this case the question of their placement in the periodic table turned out to be insoluble. Others generally denied them the right to be called elements in the classical sense. In 1902, the English physicist D. Martin called such substances radioelements. As they were studied, it became clear that some radioelements have exactly the same chemical properties, but differ in atomic masses. This circumstance contradicted the basic provisions of the periodic law. The English scientist F. Soddy resolved the contradiction. In 1913, he called chemically similar radioelements isotopes (from Greek words meaning “same” and “place”), that is, they occupy the same place in the periodic table. The radioelements turned out to be isotopes of natural radioactive elements. All of them are combined into three radioactive families, the ancestors of which are isotopes of thorium and uranium.
Isotopes of oxygen. Isobars of potassium and argon (isobars are atoms of different elements with the same mass number).
Number of stable isotopes for even and odd elements.
It soon became clear that other stable chemical elements also have isotopes. The main credit for their discovery belongs to the English physicist F. Aston. He discovered stable isotopes of many elements.
From a modern point of view, isotopes are varieties of atoms of a chemical element: they have different atomic masses, but the same nuclear charge.
Their nuclei thus contain the same number of protons, but different numbers of neutrons. For example, natural isotopes of oxygen with Z = 8 contain 8, 9 and 10 neutrons in their nuclei, respectively. The sum of the numbers of protons and neutrons in the nucleus of an isotope is called the mass number A. Consequently, the mass numbers of the indicated oxygen isotopes are 16, 17 and 18. Nowadays, the following designation for isotopes is accepted: the value Z is given below to the left of the element symbol, the value A is given to the upper left. For example: 16 8 O, 17 8 O, 18 8 O.
Since the discovery of the phenomenon of artificial radioactivity, approximately 1,800 artificial radioactive isotopes have been produced using nuclear reactions for elements with Z from 1 to 110. The vast majority of artificial radioisotopes have very short half-lives, measured in seconds and fractions of seconds; only a few have a relatively long life expectancy (for example, 10 Be - 2.7 10 6 years, 26 Al - 8 10 5 years, etc.).
Stable elements are represented in nature by approximately 280 isotopes. However, some of them turned out to be weakly radioactive, with huge half-lives (for example, 40 K, 87 Rb, 138 La, l47 Sm, 176 Lu, 187 Re). The lifespan of these isotopes is so long that they can be considered stable.
There are still many challenges in the world of stable isotopes. Thus, it is unclear why their number varies so greatly among different elements. About 25% of stable elements (Be, F, Na, Al, P, Sc, Mn, Co, As, Y, Nb, Rh, I, Cs, Pt, Tb, Ho, Tu, Ta, Au) are present in nature only one type of atom. These are the so-called single elements. It is interesting that all of them (except Be) have odd Z values. In general, for odd elements the number of stable isotopes does not exceed two. In contrast, some even-Z elements consist of a large number of isotopes (for example, Xe has 9, Sn has 10 stable isotopes).
The set of stable isotopes of a given element is called a galaxy. Their content in the galaxy often fluctuates greatly. It is interesting to note that the highest content is of isotopes with mass numbers that are multiples of four (12 C, 16 O, 20 Ca, etc.), although there are exceptions to this rule.
The discovery of stable isotopes made it possible to solve the long-standing mystery of atomic masses - their deviation from whole numbers, explained by the different percentages of stable isotopes of elements in the galaxy.
In nuclear physics the concept of “isobars” is known. Isobars are isotopes of different elements (that is, with different Z values) that have the same mass numbers. The study of isobars contributed to the establishment of many important patterns in the behavior and properties of atomic nuclei. One of these patterns is expressed by the rule formulated by the Soviet chemist S. A. Shchukarev and the German physicist I. Mattauch. It says: if two isobars differ in Z values by 1, then one of them will definitely be radioactive. A classic example of a pair of isobars is 40 18 Ar - 40 19 K. In it, the potassium isotope is radioactive. The Shchukarev-Mattauch rule made it possible to explain why there are no stable isotopes in the elements technetium (Z = 43) and promethium (Z = 61). Since they have odd Z values, more than two stable isotopes could not be expected for them. But it turned out that the isobars of technetium and promethium, respectively the isotopes of molybdenum (Z = 42) and ruthenium (Z = 44), neodymium (Z = 60) and samarium (Z = 62), are represented in nature by stable varieties of atoms in a wide range of mass numbers . Thus, physical laws prohibit the existence of stable isotopes of technetium and promethium. This is why these elements do not actually exist in nature and had to be synthesized artificially.
Scientists have long been trying to develop a periodic system of isotopes. Of course, it is based on different principles than the basis of the periodic table of elements. But these attempts have not yet led to satisfactory results. True, physicists have proven that the sequence of filling proton and neutron shells in atomic nuclei is, in principle, similar to the construction of electron shells and subshells in atoms (see Atom).
The electron shells of isotopes of a given element are constructed in exactly the same way. Therefore, their chemical and physical properties are almost identical. Only hydrogen isotopes (protium and deuterium) and their compounds exhibit noticeable differences in properties. For example, heavy water (D 2 O) freezes at +3.8, boils at 101.4 ° C, has a density of 1.1059 g/cm 3, and does not support the life of animals and plant organisms. During the electrolysis of water into hydrogen and oxygen, predominantly H 2 0 molecules are decomposed, while heavy water molecules remain in the electrolyzer.
Separating isotopes of other elements is an extremely difficult task. However, in many cases, isotopes of individual elements with significantly altered abundances compared to natural abundance are required. For example, when solving the problem of atomic energy, it became necessary to separate the isotopes 235 U and 238 U. For this purpose, the mass spectrometry method was first used, with the help of which the first kilograms of uranium-235 were obtained in the USA in 1944. However, this method proved to be too expensive and was replaced by the gas diffusion method, which used UF 6. There are now several methods for separating isotopes, but they are all quite complex and expensive. And yet the problem of “dividing the inseparable” is being successfully solved.
A new scientific discipline has emerged - isotope chemistry. She studies the behavior of various isotopes of chemical elements in chemical reactions and isotope exchange processes. As a result of these processes, the isotopes of a given element are redistributed between the reacting substances. Here is the simplest example: H 2 0 + HD = HD0 + H 2 (a water molecule exchanges a protium atom for a deuterium atom). The geochemistry of isotopes is also developing. She studies variations in the isotopic composition of different elements in the earth's crust.
The most widely used are so-called labeled atoms - artificial radioactive isotopes of stable elements or stable isotopes. With the help of isotopic indicators - labeled atoms - they study the paths of movement of elements in inanimate and living nature, the nature of the distribution of substances and elements in various objects. Isotopes are used in nuclear technology: as materials for the construction of nuclear reactors; as nuclear fuel (isotopes of thorium, uranium, plutonium); in thermonuclear fusion (deuterium, 6 Li, 3 He). Radioactive isotopes are also widely used as radiation sources.
Repeat the main points of the topic “Basic concepts of chemistry” and solve the proposed problems. Use Nos. 6-17.
Basic provisions
1. Substance(simple and complex) is any collection of atoms and molecules located in a certain state of aggregation.
Transformations of substances accompanied by changes in their composition and (or) structure are called chemical reactions .
2. Structural units substances:
· Atom- the smallest electrically neutral particle of a chemical element or simple substance, possessing all its chemical properties and then physically and chemically indivisible.
· Molecule- the smallest electrically neutral particle of a substance, possessing all its chemical properties, physically indivisible, but chemically divisible.
3. Chemical element - This is a type of atom with a certain nuclear charge.
4. Compound atom :
|
Particle |
How to determine? |
Charge |
Weight |
||
|
Cl |
conventional units |
a.e.m. |
|||
|
Electron |
By ordinal Number (N) |
1.6 ∙ 10 -19 |
9.10 ∙ 10 -28 |
0.00055 |
|
|
Proton |
By ordinal number (N) |
1.6 ∙ 10 -19 |
1.67 ∙ 10 -24 |
1.00728 |
|
|
Neutron |
Ar–N |
1.67 ∙ 10 -24 |
1.00866 |
||
5. Compound atomic nucleus :
The nucleus contains elementary particles ( nucleons) –
protons(1 1 p ) and neutrons(1 0 n ).
· Because Almost all the mass of an atom is concentrated in the nucleus and m p ≈ m n≈ 1 amu, That rounded valueA rof a chemical element is equal to the total number of nucleons in the nucleus.
7. Isotopes- a variety of atoms of the same chemical element, differing from each other only in their mass.
· Isotopic notation: to the left of the element symbol indicate the mass number (top) and atomic number of the element (bottom)
· Why do isotopes have different masses?
Assignment: Determine the atomic composition of chlorine isotopes: 35 17Cland 37 17Cl?
· Isotopes have different masses due to different numbers of neutrons in their nuclei.
8. In nature, chemical elements exist in the form of mixtures of isotopes.
The isotopic composition of the same chemical element is expressed in atomic fractions(ω at.), which indicate what part the number of atoms of a given isotope makes up of the total number of atoms of all isotopes of a given element, taken as one or 100%.
For example:
ω at (35 17 Cl) = 0.754
ω at (37 17 Cl) = 0.246
9. The periodic table shows the average values of the relative atomic masses of chemical elements, taking into account their isotopic composition. Therefore, Ar indicated in the table are fractional.
A rWed= ω at.(1) ∙ Ar (1) + … + ω at.(n ) ∙ Ar ( n )
For example:
A rWed(Cl) = 0.754 ∙ 35 + 0.246 ∙ 37 = 35.453
10. Problem to solve:
No. 1. Determine the relative atomic mass of boron if it is known that the molar fraction of the 10 B isotope is 19.6%, and the 11 B isotope is 80.4%.
11. The masses of atoms and molecules are very small. Currently, a unified measurement system has been adopted in physics and chemistry.
1 amu =m(a.u.m.) = 1/12 m(12 C) = 1.66057 ∙ 10 -27 kg = 1.66057 ∙ 10 -24 g.
Absolute masses of some atoms:
m( C) =1.99268 ∙ 10 -23 g
m( H) =1.67375 ∙ 10 -24 g
m( O) =2.656812 ∙ 10 -23 g
A r– shows how many times a given atom is heavier than 1/12 of a 12 C atom. Mr∙ 1.66 ∙ 10 -27 kg
13. The number of atoms and molecules in ordinary samples of substances is very large, therefore, when characterizing the amount of a substance, the unit of measurement is used -mole .
· Mole (ν)– a unit of quantity of a substance that contains the same number of particles (molecules, atoms, ions, electrons) as there are atoms in 12 g of isotope 12 C
· Mass of 1 atom 12 C is equal to 12 amu, so the number of atoms in 12 g of isotope 12 C equals:
N A= 12 g / 12 ∙ 1.66057 ∙ 10 -24 g = 6.0221 ∙ 10 23
· Physical quantity N A called Avogadro's constant (Avogadro's number) and has the dimension [N A] = mol -1.
14. Basic formulas:
M = Mr = ρ ∙ Vm(ρ – density; V m – volume at zero level)
Problems to solve independently
No. 1. Calculate the number of nitrogen atoms in 100 g of ammonium carbonate containing 10% non-nitrogen impurities.
No. 2. Under normal conditions, 12 liters of a gas mixture consisting of ammonia and carbon dioxide have a mass of 18 g. How many liters of each gas does the mixture contain?
No. 3. When exposed to excess hydrochloric acid, 8.24 g of a mixture of manganese oxide (IV) with the unknown oxide MO 2, which does not react with hydrochloric acid, 1.344 liters of gas were obtained at ambient conditions. In another experiment, it was established that the molar ratio of manganese oxide (IV) to the unknown oxide is 3:1. Determine the formula of the unknown oxide and calculate its mass fraction in the mixture.
There is probably not a person on earth who has not heard of isotopes. But not everyone knows what it is. The phrase “radioactive isotopes” sounds especially frightening. These strange chemical elements terrify humanity, but in reality they are not as scary as they might seem at first glance.
Definition
To understand the concept of radioactive elements, it is necessary to first say that isotopes are samples of the same chemical element, but with different masses. What does it mean? The questions will disappear if we first remember the structure of the atom. It consists of electrons, protons and neutrons. The number of the first two elementary particles in the nucleus of an atom is always constant, while neutrons, which have their own mass, can occur in the same substance in different quantities. This circumstance gives rise to a variety of chemical elements with different physical properties.
Now we can give a scientific definition to the concept under study. So, isotopes are a collective set of chemical elements that are similar in properties, but have different masses and physical properties. According to more modern terminology, they are called a galaxy of nucleotides of a chemical element.
A little history
At the beginning of the last century, scientists discovered that the same chemical compound under different conditions can have different masses of electron nuclei. From a purely theoretical point of view, such elements could be considered new and they could begin to fill empty cells in D. Mendeleev’s periodic table. But there are only nine free cells in it, and scientists discovered dozens of new elements. In addition, mathematical calculations showed that the discovered compounds cannot be considered previously unknown, because their chemical properties fully corresponded to the characteristics of existing ones.
After lengthy discussions, it was decided to call these elements isotopes and place them in the same box as those whose nuclei contain the same number of electrons. Scientists have been able to determine that isotopes are just some variations of chemical elements. However, the causes of their occurrence and life expectancy have been studied for almost a century. Even at the beginning of the 21st century, it is impossible to say that humanity knows absolutely everything about isotopes.
Persistent and unstable variations
Each chemical element has several isotopes. Due to the fact that there are free neutrons in their nuclei, they do not always enter into stable bonds with the rest of the atom. After some time, free particles leave the nucleus, which changes its mass and physical properties. In this way, other isotopes are formed, which ultimately leads to the formation of a substance with an equal number of protons, neutrons and electrons.
Those substances that decay very quickly are called radioactive isotopes. They release a large number of neutrons into space, forming powerful ionizing gamma radiation, known for its strong penetrating power, which negatively affects living organisms.
More stable isotopes are not radioactive, since the number of free neutrons released by them is not capable of generating radiation and significantly affecting other atoms.
Quite a long time ago, scientists established one important pattern: each chemical element has its own isotopes, persistent or radioactive. Interestingly, many of them were obtained in laboratory conditions, and their presence in natural form is small and is not always detected by instruments.
Distribution in nature
Under natural conditions, substances are most often found whose isotope mass is directly determined by its ordinal number in D. Mendeleev’s table. For example, hydrogen, denoted by the symbol H, has an atomic number of 1, and its mass is equal to one. Its isotopes, 2H and 3H, are extremely rare in nature.
Even the human body has some radioactive isotopes. They enter through food in the form of carbon isotopes, which, in turn, are absorbed by plants from the soil or air and become part of organic matter during the process of photosynthesis. Therefore, humans, animals, and plants emit a certain background radiation. Only it is so low that it does not interfere with normal functioning and growth.
The sources that contribute to the formation of isotopes are the inner layers of the earth's core and radiation from space.
As you know, the temperature on a planet largely depends on its hot core. But only very recently it became clear that the source of this heat is a complex thermonuclear reaction in which radioactive isotopes participate.
Isotopic Decay
Since isotopes are unstable formations, it can be assumed that over time they always decay into more permanent nuclei of chemical elements. This statement is true because scientists have not been able to detect huge amounts of radioactive isotopes in nature. And most of those that were extracted in laboratories lasted from a couple of minutes to several days, and then turned back into ordinary chemical elements.
But there are also isotopes in nature that turn out to be very resistant to decay. They can exist for billions of years. Such elements were formed in those distant times, when the earth was still being formed, and there was not even a solid crust on its surface.
Radioactive isotopes decay and form again very quickly. Therefore, in order to facilitate the assessment of the stability of the isotope, scientists decided to consider the category of its half-life.
Half life
It may not be immediately clear to all readers what is meant by this concept. Let's define it. The half-life of an isotope is the time during which a conventional half of the substance taken will cease to exist.
This does not mean that the rest of the connection will be destroyed in the same amount of time. In relation to this half, it is necessary to consider another category - the period of time during which its second part, that is, a quarter of the original amount of substance, will disappear. And this consideration continues ad infinitum. It can be assumed that it is simply impossible to calculate the time for complete disintegration of the initial amount of a substance, since this process is practically endless.
However, scientists, knowing the half-life, can determine how much of the substance existed at the beginning. These data are successfully used in related sciences.
In the modern scientific world, the concept of complete decay is practically not used. For each isotope, it is customary to indicate its half-life, which varies from a few seconds to many billions of years. The lower the half-life, the more radiation comes from the substance and the higher its radioactivity.
Fossil beneficiation
In some branches of science and technology, the use of relatively large quantities of radioactive substances is considered mandatory. However, under natural conditions there are very few such compounds.
It is known that isotopes are uncommon variants of chemical elements. Their number is measured in several percent of the most resistant variety. This is why scientists need to artificially enrich fossil materials.
Over the years of research, we have learned that the decay of an isotope is accompanied by a chain reaction. The released neutrons of one substance begin to influence another. As a result of this, heavy nuclei disintegrate into lighter ones and new chemical elements are obtained.
This phenomenon is called a chain reaction, as a result of which more stable but less common isotopes can be obtained, which are subsequently used in the national economy.
Application of decay energy
Scientists also found that during the decay of a radioactive isotope, a huge amount of free energy is released. Its amount is usually measured by the Curie unit, equal to the fission time of 1 g of radon-222 in 1 second. The higher this indicator, the more energy is released.
This became the reason for developing ways to use free energy. This is how atomic reactors appeared, into which a radioactive isotope is placed. Most of the energy released by it is collected and converted into electricity. Based on these reactors, nuclear power plants are created that provide the cheapest electricity. Smaller versions of such reactors are installed on self-propelled mechanisms. Given the danger of accidents, submarines are most often used as such vehicles. In the event of a reactor failure, the number of casualties on the submarine will be easier to minimize.
Another very scary option for using half-life energy is atomic bombs. During World War II, they were tested on humans in the Japanese cities of Hiroshima and Nagasaki. The consequences were very sad. Therefore, there is an agreement in the world on the non-use of these dangerous weapons. At the same time, large states with a focus on militarization continue research in this area today. In addition, many of them, secretly from the world community, are producing atomic bombs, which are thousands of times more dangerous than those used in Japan.
Isotopes in medicine
For peaceful purposes, they have learned to use the decay of radioactive isotopes in medicine. By directing radiation to the affected area of the body, it is possible to stop the course of the disease or help the patient to recover completely.
But more often radioactive isotopes are used for diagnostics. The thing is that their movement and the nature of the cluster are most easily determined by the radiation they produce. Thus, a certain non-hazardous amount of a radioactive substance is introduced into the human body, and doctors use instruments to observe how and where it gets into.
In this way, they diagnose the functioning of the brain, the nature of cancerous tumors, and the peculiarities of the functioning of the endocrine and exocrine glands.
Application in archeology
It is known that living organisms always contain radioactive carbon-14, the half-life of which isotope is 5570 years. In addition, scientists know how much of this element is contained in the body until the moment of death. This means that all cut trees emit the same amount of radiation. Over time, the radiation intensity decreases.
This helps archaeologists determine how long ago the wood from which the galley or any other ship was built died, and therefore the time of construction itself. This research method is called radioactive carbon analysis. Thanks to it, it is easier for scientists to establish the chronology of historical events.
A certain element that has the same but different. They have nuclei with the same number and diversity. number, have the same structure of electron shells and occupy the same place in the periodicity. chemical system elements. The term "isotopes" was proposed in 1910 by F. Soddy to designate chemically indistinguishable varieties that differ in their physical properties. (primarily radioactive) Saints. Stable isotopes were first discovered in 1913 by J. Thomson using the so-called he developed. the method of parabolas - the prototype of the modern one. . He found that Ne has at least 2 varieties with a wt. parts 20 and 22. The names and symbols of isotopes are usually the names and symbols of the corresponding chemicals. elements; point to the top left of the symbol. For example, to indicate natural isotopes use the notation 35 Cl and 37 Cl; sometimes the element is also indicated at the bottom left, i.e. write 35 17 Cl and 37 17 Cl. Only isotopes of the lightest element, hydrogen, with wt. parts 1, 2 and 3 have special. names and symbols: (1 1 H), (D, or 2 1 H) and (T, or 3 1 H), respectively. Due to the large difference in masses, the behavior of these isotopes differs significantly (see,). Stable isotopes occur in all even and most odd elements with[ 83. The number of stable isotopes of elements with even numbers may be. equals 10 (e.g. y); Odd-numbered elements have no more than two stable isotopes. Known approx. 280 stable and more than 2000 radioactive isotopes of 116 natural and artificially obtained elements. For each element, the content of individual isotopes in nature. the mixture undergoes small fluctuations, which can often be neglected. More means. fluctuations in the isotopic composition are observed for meteorites and other celestial bodies. The constancy of the isotopic composition leads to the constancy of the elements found on Earth, which is the average value of the mass of a given element, found taking into account the abundance of isotopes in nature. Fluctuations in the isotopic composition of light elements are associated, as a rule, with changes in the isotopic composition during decomposition. processes occurring in nature (, etc.). For the heavy element Pb, variations in the isotopic composition of different samples are explained by different factors. content in, and other sources and - the ancestors of nature. . Differences in the properties of isotopes of a given element are called. . Important practical The task is to obtain from nature. mixtures of individual isotopes -