Atomic mass, relative atomic mass(obsolete name - atomic weight) - the value of the mass of an atom, expressed in atomic mass units. Currently, the atomic mass unit is taken to be equal to 1/12 of the mass of a neutral atom of the most common isotope of carbon 12C, so the atomic mass of this isotope by definition is exactly 12. For any other isotope, the atomic mass is not an integer, although it is close to the mass number of this isotope (i.e., the total number of nucleons - protons and neutrons - in its nucleus). The difference between the atomic mass of an isotope and its mass number is called excess mass (usually expressed in MeVah). It can be either positive or negative; the reason for its occurrence is the nonlinear dependence of the binding energy of nuclei on the number of protons and neutrons, as well as the difference in the masses of the proton and neutron.
The dependence of the atomic mass on the mass number is as follows: the excess mass is positive for hydrogen-1, with increasing mass number it decreases and becomes negative until a minimum is reached for iron-56, then it begins to grow and increases to positive values for heavy nuclides. This corresponds to the fact that the fission of nuclei heavier than iron releases energy, while the fission of light nuclei requires energy. On the contrary, the fusion of nuclei lighter than iron releases energy, while the fusion of elements heavier than iron requires additional energy.
The atomic mass of a chemical element (also “average atomic mass”, “standard atomic mass”) is the weighted average atomic mass of all stable isotopes of a given chemical element, taking into account their natural abundance in the earth’s crust and atmosphere. It is this atomic mass that is presented in the periodic table and is used in stoichiometric calculations. The atomic mass of an element with a disturbed isotopic ratio (for example, enriched in some isotope) differs from the standard one.
The molecular mass of a mochemical compound is the sum of the atomic masses of the elements that compose it, multiplied by the stoichiometric coefficients of the elements according to the chemical formula of the compound. Strictly speaking, the mass of a molecule is less than the mass of its constituent atoms by an amount equal to the binding energy of the molecule. However, this mass defect is 9-10 orders of magnitude less than the mass of the molecule, and can be neglected.
The definition of a mole (and Avogadro's number) is chosen so that the mass of one mole of a substance (molar mass), expressed in grams, is numerically equal to the atomic (or molecular) mass of that substance. For example, the atomic mass of iron is 55.847. Therefore, one mole of iron atoms (i.e., their number is equal to Avogadro’s number, 6.022 1023) contains 55.847 grams.
Direct comparison and measurement of the masses of atoms and molecules is performed using mass spectrometric methods.
Story
Until the 1960s, atomic mass was defined so that the isotope oxygen-16 had an atomic mass of 16 (oxygen scale). However, the ratio of oxygen-17 and oxygen-18 in natural oxygen, which was also used in atomic mass calculations, resulted in two different tables of atomic masses. Chemists used a scale based on the fact that the natural mixture of oxygen isotopes would have an atomic mass of 16, while physicists assigned the same number of 16 to the atomic mass of the most common isotope of oxygen (which has eight protons and eight neutrons).
Wikipedia
The masses of atoms and molecules are very small, so it is convenient to choose the mass of one of the atoms as a unit of measurement and express the masses of the remaining atoms relative to it. This is exactly what the founder of atomic theory, Dalton, did, who compiled a table of atomic masses, taking the mass of the hydrogen atom as one.
Until 1961, in physics, 1/16 of the mass of the 16O oxygen atom was taken as an atomic mass unit (amu), and in chemistry – 1/16 of the average atomic mass of natural oxygen, which is a mixture of three isotopes. The chemical unit of mass was 0.03% larger than the physical one.
Currently, a unified measurement system has been adopted in physics and chemistry. 1/12 of the mass of a 12C carbon atom was chosen as the standard unit of atomic mass.
1 amu = 1/12 m(12С) = 1.66057×10-27 kg = 1.66057×10-24 g.
When calculating relative atomic mass, the abundance of isotopes of elements in the earth's crust is taken into account. For example, chlorine has two isotopes 35Сl (75.5%) and 37Сl (24.5%). The relative atomic mass of chlorine is:
Ar(Cl) = (0.755×m(35Сl) + 0.245×m(37Сl)) / (1/12×m(12С) = 35.5.
From the definition of relative atomic mass it follows that the average absolute mass of an atom is equal to the relative atomic mass multiplied by amu:
m(Cl) = 35.5 × 1.66057 × 10-24 = 5.89 × 10-23 g.
Examples of problem solving
Relative atomic and molecular masses
This calculator is designed to calculate the atomic mass of elements.
Atomic mass(also called relative atomic mass) Is the value of the mass of one atom of a substance. Relative atomic mass is expressed in atomic mass units. Relative atomic mass distinctive(True) weight atom. At the same time, the actual mass of an atom is too small and therefore unsuitable for practical use.
The atomic mass of a substance affects the amount protons And neutrons in the nucleus of an atom.
The electron mass is ignored since it is very small.
To determine the atomic mass of a substance, you must enter the following information:
- Number of protons- how many protons are in the nucleus of the substance;
- Number of neutrons— how many neutrons are in the nucleus of a substance.
Based on this data, the calculator will calculate the atomic mass of the substance, expressed in atomic mass units.
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Table of chemical elements and their atomic mass
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Relative atomic mass of an elementTask status:Determine the mass of an oxygen molecule. Task no. 4.1.2 from the “Collection of problems in preparing upcoming exams in physics at USPTU” information:Solution:Consider a molecular oxygen molecule \(\nu\) (arbitrary number). Let us remember that the oxygen formula is O2. To find the mass (\m) of a given amount of oxygen, the molecular mass of oxygen\(M\) is multiplied by the number of moles\(\nu\). Using the periodic table, it is easy to establish that the molar mass of oxygen is \(M\) 32 g/mol or 0.032 kg/mol. In one mol, the number of avogadro molecules \(N_A\) and v\(\nu\) mol - v\(\nu\) is sometimes greater, i.e. To find the mass of one molecule \(m_0\), the total mass \(m\) must be divided by the number of molecules \(N\). \ [(m_0) = \frac (m) (N)\] \ [(m_0) = \frac ((\nu \cdot M)) ((\nu \cdot (N_A)))\] \ ((M_0) = \frac (M) (((N_A))) \] Avogadro's number (N_A1) is a tabular value equal to 6.022 1023 mol-1. We perform calculations: \[(M_0) = \frac ((0.032)) ((6.022\cdot ((10) * (23)))) = 5.3\cdot (10^(-26))\; = 5.3 kg\cdot(10^(-23))\; r\] Answer: 5.3 · 10-23 g.If you don't understand the solution and if you have any questions or found a bug, you can leave a comment below. Atoms are very small and very small. If we express the mass of an atom of a chemical element in grams, then it will be a number for which the decimal point is more than twenty zeros. Therefore, measuring the mass of atoms in grams is inappropriate. However, if we take a very small mass per unit, all other small masses can be expressed as a ratio between that unit. The unit of measurement for atomic mass is 1/12 of the mass of a carbon atom. It is called 1/12 of the mass of a carbon atom atomic mass(Ae. Atomic mass formulaRelative atomic mass the value is equal to the ratio of the actual mass of an atom of a particular chemical element to 1/12 of the actual mass of a carbon atom. This is an infinite value, since the two masses are separated. Ar = mathematics. / (1/12) mug. Nevertheless, absolute atomic mass equal to a relative value and has a measurement unit amu. This means that relative atomic mass shows how many times the mass of a given atom is greater than 1/12 of a carbon atom. If an Ar atom = 12, then its mass is 12 times greater than 1/12 the mass of a carbon atom or, in other words, 12 atomic mass units. This can only be for carbon (C). On the hydrogen atom (H) Ar = 1. This means that its mass is equal to the mass of 1/12 parts of the mass of the carbon atom. For oxygen (O), the relative atomic mass is 16 amu. This means that an oxygen atom is 16 times larger than a carbon atom, it has 16 atomic mass units. The lightest element is hydrogen. Its mass is about 1 amu. On the heaviest atoms the mass approaches 300 amu. Typically, for each chemical element, its value is the absolute mass of the atoms, expressed as a. For example. The meaning of atomic mass units is written in the periodic table. Concept used for molecules relative molecular weight (g). Relative molecular weight indicates how many times the mass of a molecule is greater than 1/12 the mass of a carbon atom. However, since the mass of a molecule is equal to the sum of the masses of its atomic atoms, the relative molecular mass can be found simply by adding the relative masses of those atoms. For example, a water molecule (H2O) contains two hydrogen atoms with Ar = 1 and one oxygen atom with Ar = 16. Therefore, gentleman (H2O) = 18. Many substances have a non-molecular structure, such as metals. In this case, their relative molecular mass is equal to their relative atomic mass. Chemistry is called a significant amount mass fraction of a chemical element in a molecule or substance. It shows the relative molecular weight of that element. For example, in water, hydrogen has 2 parts (as both atoms) and oxygen 16. This means that when hydrogen is mixed with 1 kg and 8 kg of oxygen, they react without a residue. The mass fraction of hydrogen is 2/18 = 1/9, and the oxygen content is 16/18 = 8/9. Microbalance otherwise support, atomic equilibrium(English microbial or English nanotubes) is a term referring to:
descriptionOne of the first references to the microglob is in 1910, when William Ramsay was informed of the extent to which it had evolved, allowing the weight range of 0.1 mm3 of body to be determined to be 10-9 g (1 ng). The term "microbial" is now more commonly used to refer to devices that can measure and detect mass changes in the microgram range (10-6 grams). Microbiologists have become common practice in modern research and industrial laboratories and are available in different versions with varying sensitivities and associated costs. At the same time, measurement techniques are being developed in the nanogram field. chemistry. how to find relative atomic mass?When we talk about measuring mass at the nanogram level, which is important for measuring the mass of atoms, molecules or clusters, we first consider mass spectrometry. In this case, it should be borne in mind that measuring mass using this method implies the need to convert the weighed objects into ions, which is sometimes very undesirable. This is not necessary when using another practically important and widely used instrument for the accurate measurement of mass quartz microbes, the mechanism of action of which is described in the corresponding article. links
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Atomic mass is the sum of the masses of all protons, neutrons and electrons that make up an atom or molecule. Compared to protons and neutrons, the mass of electrons is very small, so it is not taken into account in calculations. Although this is not formally correct, the term is often used to refer to the average atomic mass of all isotopes of an element. This is actually relative atomic mass, also called atomic weight element. Atomic weight is the average of the atomic masses of all isotopes of an element found in nature. Chemists must differentiate between these two types of atomic mass when doing their work—an incorrect atomic mass value can, for example, result in an incorrect result for the yield of a reaction.
Steps
Finding atomic mass from the periodic table of elements
- The atomic mass unit characterizes the mass one mole of a given element in grams. This value is very useful in practical calculations, since it can be used to easily convert the mass of a given number of atoms or molecules of a given substance into moles, and vice versa.
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Find the atomic mass in the periodic table. Most standard periodic tables contain the atomic masses (atomic weights) of each element. Typically, they are listed as a number at the bottom of the element cell, below the letters representing the chemical element. Usually this is not a whole number, but a decimal fraction.
Remember that the periodic table gives the average atomic masses of elements. As noted earlier, the relative atomic masses given for each element in the periodic table are the average of the masses of all isotopes of the atom. This average value is valuable for many practical purposes: for example, it is used in calculating the molar mass of molecules consisting of several atoms. However, when you are dealing with individual atoms, this value is usually not enough.
- Since the average atomic mass is an average of several isotopes, the value shown in the periodic table is not accurate the value of the atomic mass of any single atom.
- The atomic masses of individual atoms must be calculated taking into account the exact number of protons and neutrons in a single atom.
Learn how atomic mass is written. Atomic mass, that is, the mass of a given atom or molecule, can be expressed in standard SI units - grams, kilograms, and so on. However, because atomic masses expressed in these units are extremely small, they are often written in unified atomic mass units, or amu for short. – atomic mass units. One atomic mass unit is equal to 1/12 the mass of the standard isotope carbon-12.
Calculation of the atomic mass of an individual atom
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Find the atomic number of a given element or its isotope. Atomic number is the number of protons in the atoms of an element and never changes. For example, all hydrogen atoms, and only they have one proton. The atomic number of sodium is 11 because it has eleven protons in its nucleus, while the atomic number of oxygen is eight because it has eight protons in its nucleus. You can find the atomic number of any element in the periodic table - in almost all its standard versions, this number is indicated above the letter designation of the chemical element. The atomic number is always a positive integer.
- Suppose we are interested in the carbon atom. Carbon atoms always have six protons, so we know that its atomic number is 6. In addition, we see that in the periodic table, at the top of the cell with carbon (C) is the number "6", indicating that the atomic carbon number is six.
- Note that the atomic number of an element is not uniquely related to its relative atomic mass in the periodic table. Although, especially for the elements at the top of the table, it may appear that an element's atomic mass is twice its atomic number, it is never calculated by multiplying the atomic number by two.
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Find the number of neutrons in the nucleus. The number of neutrons can be different for different atoms of the same element. When two atoms of the same element with the same number of protons have different numbers of neutrons, they are different isotopes of that element. Unlike the number of protons, which never changes, the number of neutrons in the atoms of a given element can often change, so the average atomic mass of an element is written as a decimal fraction with a value lying between two adjacent whole numbers.
Add up the number of protons and neutrons. This will be the atomic mass of this atom. Ignore the number of electrons that surround the nucleus - their total mass is extremely small, so they have virtually no effect on your calculations.
Calculating the relative atomic mass (atomic weight) of an element
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Determine which isotopes are contained in the sample. Chemists often determine the isotope ratios of a particular sample using a special instrument called a mass spectrometer. However, in training, this data will be provided to you in assignments, tests, and so on in the form of values taken from the scientific literature.
- In our case, let's say that we are dealing with two isotopes: carbon-12 and carbon-13.
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Determine the relative abundance of each isotope in the sample. For each element, different isotopes occur in different ratios. These ratios are almost always expressed as percentages. Some isotopes are very common, while others are very rare—sometimes so rare that they are difficult to detect. These values can be determined using mass spectrometry or found in a reference book.
- Let's assume that the concentration of carbon-12 is 99% and carbon-13 is 1%. Other isotopes of carbon really exist, but in quantities so small that in this case they can be neglected.
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Multiply the atomic mass of each isotope by its concentration in the sample. Multiply the atomic mass of each isotope by its percentage abundance (expressed as a decimal). To convert percentages to a decimal, simply divide them by 100. The resulting concentrations should always add up to 1.
- Our sample contains carbon-12 and carbon-13. If carbon-12 makes up 99% of the sample and carbon-13 makes up 1%, then multiply 12 (the atomic mass of carbon-12) by 0.99 and 13 (the atomic mass of carbon-13) by 0.01.
- The reference books give percentages based on the known quantities of all isotopes of a particular element. Most chemistry textbooks contain this information in a table at the end of the book. For the sample being studied, the relative concentrations of isotopes can also be determined using a mass spectrometer.
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Add up the results. Sum up the multiplication results you got in the previous step. As a result of this operation, you will find the relative atomic mass of your element - the average value of the atomic masses of the isotopes of the element in question. When considering an element as a whole, rather than a specific isotope of a given element, this is the value used.
- In our example, 12 x 0.99 = 11.88 for carbon-12, and 13 x 0.01 = 0.13 for carbon-13. The relative atomic mass in our case is 11.88 + 0.13 = 12,01 .
- Some isotopes are less stable than others: they break down into atoms of elements with fewer protons and neutrons in the nucleus, releasing particles that make up the atomic nucleus. Such isotopes are called radioactive.
One of the fundamental properties of atoms is their mass. Absolute (true) mass of an atom– the value is extremely small. It is impossible to weigh atoms on a balance because such precise scales do not exist. Their masses were determined using calculations.
For example, the mass of one hydrogen atom is 0.000 000 000 000 000 000 000 001 663 grams! The mass of a uranium atom, one of the heaviest atoms, is approximately 0.000 000 000 000 000 000 000 4 grams.
The exact mass of the uranium atom is 3.952 ∙ 10−22 g, and the hydrogen atom, the lightest among all atoms, is 1.673 ∙ 10−24 g.
It is inconvenient to perform calculations with small numbers. Therefore, instead of the absolute masses of atoms, their relative masses are used.
Relative atomic mass
The mass of any atom can be judged by comparing it with the mass of another atom (find the ratio of their masses). Since the determination of the relative atomic masses of elements, various atoms have been used as comparisons. At one time, hydrogen and oxygen atoms were unique standards for comparison.
A unified scale of relative atomic masses and a new unit of atomic mass, adopted International Congress of Physicists (1960) and unified by the International Congress of Chemists (1961).
To this day, the standard for comparison is 1/12 of the mass of a carbon atom. This value is called the atomic mass unit, abbreviated a.u.m.
Atomic mass unit (amu) – mass of 1/12 of a carbon atom
Let’s compare how many times the absolute mass of a hydrogen and uranium atom differs from 1 amu, to do this we divide these numbers by one another:
The values obtained in the calculations are the relative atomic masses of the elements - relative 1/12 the mass of a carbon atom.
Thus, the relative atomic mass of hydrogen is approximately 1, and that of uranium is 238. Please note that relative atomic mass does not have units of measurement, since the units of absolute mass (grams) are canceled out when dividing.
The relative atomic masses of all elements are indicated in the Periodic Table of Chemical Elements by D.I. Mendeleev. The symbol used to indicate relative atomic mass is Аr (the letter r is an abbreviation for the word relative, which means relative).
The relative atomic masses of elements are used in many calculations. As a rule, values given in the Periodic Table are rounded to whole numbers. Note that the elements in the Periodic Table are arranged in order of increasing relative atomic masses.
For example, using the Periodic Table we determine the relative atomic masses of a number of elements:
Ar(O) = 16; Ar(Na) = 23; Ar(P) = 31.
The relative atomic mass of chlorine is usually written as 35.5!
Ar(Cl) = 35.5
- Relative atomic masses are proportional to the absolute masses of atoms
- The standard for determining relative atomic mass is 1/12 of the mass of a carbon atom
- 1 amu = 1.662 ∙ 10−24 g
- Relative atomic mass is denoted by Ar
- For calculations, the values of relative atomic masses are rounded to whole numbers, with the exception of chlorine, for which Ar = 35.5
- Relative atomic mass has no units of measurement
Contents of the article
ATOMIC MASS. The concept of this quantity has undergone long-term changes in accordance with changes in the concept of atoms. According to Dalton's theory (1803), all atoms of the same chemical element are identical and its atomic mass is a number equal to the ratio of their mass to the mass of an atom of a certain standard element. However, by about 1920 it became clear that elements found in nature were of two types: some actually represented by identical atoms, while others had atoms with the same nuclear charge but different masses; These types of atoms were called isotopes. Dalton's definition is thus valid only for elements of the first type. The atomic mass of an element represented by several isotopes is the average of the mass numbers of all its isotopes, taken as a percentage corresponding to their abundance in nature.
In the 19th century Chemists used hydrogen or oxygen as a standard when determining atomic masses. In 1904, 1/16 of the average mass of an atom of natural oxygen (oxygen unit) was adopted as the standard, and the corresponding scale was called chemical. Mass spectrographic determination of atomic masses was carried out on the basis of 1/16 of the mass of the 16 O isotope, and the corresponding scale was called physical. In the 1920s, it was discovered that natural oxygen consists of a mixture of three isotopes: 16 O, 17 O and 18 O. This raised two problems. First, it turns out that the relative abundance of natural oxygen isotopes varies slightly, which means that the chemical scale is based on a value that is not an absolute constant. Secondly, physicists and chemists obtained different values for such derivative constants as molar volumes, Avogadro’s number, etc. The solution to the problem was found in 1961, when 1/12 of the mass was taken as the atomic mass unit (amu) carbon isotope 12 C (carbon unit). (1 amu, or 1D (dalton), in SI mass units is 1.66057Х10 –27 kg.) Natural carbon also consists of two isotopes: 12 C – 99% and 13 C – 1%, but new values of atomic masses of elements are associated only with the first of them. As a result, a universal table of relative atomic masses was obtained. The 12 C isotope also turned out to be convenient for physical measurements.
METHODS OF DETERMINATION
Atomic mass can be determined either by physical or chemical methods. Chemical methods differ in that at one stage they involve not the atoms themselves, but their combinations.
Chemical methods.
According to atomic theory, the numbers of atoms of elements in compounds are related to each other as small integers (the law of multiple ratios, which was discovered by Dalton). Therefore, for a compound of known composition, it is possible to determine the mass of one of the elements, knowing the masses of all the others. In some cases, the mass of a compound can be measured directly, but it is usually found by indirect methods. Let's look at both of these approaches.
The atomic mass of Al was recently determined as follows. Known quantities of Al were converted to nitrate, sulfate or hydroxide and then calcined to aluminum oxide (Al 2 O 3), the amount of which was precisely determined. From the relationship between two known masses and the atomic masses of aluminum and oxygen (15.9)
found the atomic mass of Al. However, by direct comparison with the atomic mass of oxygen, the atomic masses of only a few elements can be determined. For most elements, they were determined indirectly by analyzing chlorides and bromides. Firstly, these compounds for many elements can be obtained in pure form, and secondly, for their precise quantitative determination, chemists have at their disposal a sensitive analytical method based on comparing their masses with the mass of silver. To do this, accurately determine the mass of the analyzed compounds and the mass of silver necessary to interact with them. The atomic mass of the desired element is calculated based on the atomic mass of silver - the reference value in such definitions. The atomic mass of silver (107.870) in carbon units was determined by an indirect chemical method.
Physical methods.
In the middle of the 20th century. There was only one physical method for determining atomic masses; today four are the most widely used.
Gas density.
The very first physical method was based on determining the density of a gas and on the fact that, in accordance with Avogadro's law, equal volumes of gases at the same temperature and pressure contain the same number of molecules. Therefore, if a certain volume of pure CO 2 has a mass 1.3753 greater than the same volume of oxygen under the same conditions, then the CO 2 molecule must be 1.3753 times heavier than the oxygen molecule (molecular mass of O 2 = 31.998) , i.e. the mass of a CO 2 molecule on the chemical scale is 44.008. If we subtract the mass of two oxygen atoms, equal to 31.998, from this value, we get the atomic mass of carbon - 12.01. To obtain a more accurate value, it is necessary to introduce a number of corrections, which complicates this method. Nevertheless, with its help some very valuable data were obtained. Thus, after the discovery of noble gases (He, Ne, Ar, Kr, Xe), the method based on density measurements turned out to be the only suitable one for determining their atomic masses.
Mass spectroscopy.
Soon after the First World War, F. Aston created the first mass spectroscope to accurately determine the mass numbers of various isotopes and thereby opened a new era in the history of determining atomic masses. Today there are two main types of mass spectroscopes: mass spectrometers and mass spectrographs (the latter is, for example, the Aston instrument). A mass spectrograph is designed to study the behavior of a flow of electrically charged atoms or molecules in a strong magnetic field. The deflection of charged particles in this field is proportional to the ratio of their masses to the charge, and they are recorded in the form of lines on a photographic plate. By comparing the positions of the lines corresponding to certain particles with the position of the line for an element with a known atomic mass, it is possible to determine the atomic mass of the desired element with sufficient accuracy. A good illustration of the method is to compare the mass of a CH 4 (methane) molecule with the mass number of the lightest isotope of oxygen, 16 O. Equally charged methane and 16 O ions are simultaneously admitted into the mass spectrograph chamber and their position is recorded on a photographic plate. The difference in the position of their lines corresponds to a mass difference of 0.036406 (on the physical scale). This is significantly higher accuracy than any chemical method can provide.
If the element under study does not have isotopes, then determining its atomic mass is not difficult. Otherwise, it is necessary to determine not only the mass of each isotope, but also their relative abundance in the mixture. This value cannot be determined with sufficient accuracy, which limits the use of the mass spectrographic method for finding the atomic masses of isotopic elements, especially heavy ones. Recently, using mass spectrometry, it was possible to establish with high accuracy the relative abundance of two isotopes of silver, 107 Ag and 109 Ag. Measurements were performed at the US National Bureau of Standards. Using these new data and earlier measurements of the masses of silver isotopes, the atomic mass of natural silver was clarified. This value is now considered to be 107.8731 (chemical scale).
Nuclear reactions.
To determine the atomic masses of some elements, we can use the relationship between mass and energy obtained by Einstein. Let us consider the reaction of bombardment of 14 N nuclei by fast deuterium nuclei with the formation of the 15 N isotope and ordinary hydrogen 1 H:
14 N + 2 H = 15 N + 1 H + Q
The reaction releases energy Q= 8,615,000 eV, which, according to Einstein’s equation, is equivalent to 0.00948 amu. This means that the mass of 14 N + 2 H exceeds the mass of 15 N + 1 H by 0.00948 amu, and if we know the mass numbers of any three isotopes participating in the reaction, we can find the mass of the fourth. The method allows you to determine the difference in the mass numbers of two isotopes with greater accuracy than mass spectrography.
Radiography.
This physical method can be used to determine the atomic masses of substances that form a regular crystal lattice at ordinary temperatures. The method is based on the relationship between the atomic (or molecular) mass of a crystalline substance, its density, Avogadro's number and a certain coefficient, which is determined from the distances between atoms in the crystal lattice. It is necessary to carry out precision measurements of two quantities: lattice constant using radiographic methods and density using pycnometry. The application of the method is limited by the difficulties of obtaining pure perfect crystals (without vacancies and defects of any kind).
Clarification of atomic masses.
All measurements of atomic masses that were carried out more than 20 years ago were carried out using chemical methods or a method based on determining the density of gases. Recently, data obtained by mass spectrometric and isotope methods coincide with such high accuracy that the International Atomic Mass Commission decided to correct the atomic masses of 36 elements, 18 of which do not have isotopes.
See also
