Relative atomic mass of an element in chemistry and the history of its determination. Chemical encyclopedia What is atomic mass, what does it mean and how to spell it correctly

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 the calculations. Although it is incorrect from a formal point of view, this term is often used to refer to the average atomic mass of all isotopes of an element. In fact, this is the relative atomic mass, also called atomic weight element. Atomic weight is the average of the atomic masses of all naturally occurring isotopes of an element. Chemists must distinguish between these two types of atomic mass when doing their job - an incorrect value for atomic mass can, for example, lead to an incorrect result for the yield of a reaction product.

Steps

Finding the atomic mass according to the periodic table of elements

    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, due to the fact that atomic masses expressed in these units are extremely small, they are often written in unified atomic mass units, or a.u.m. for short. are atomic mass units. One atomic mass unit is equal to 1/12 the mass of the standard carbon-12 isotope.

    • The atomic mass unit characterizes the mass one mole of the 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.

  1. Find the atomic mass in Mendeleev's periodic table. Most standard periodic tables contain the atomic masses (atomic weights) of each element. As a rule, they are given as a number at the bottom of the cell with the element, under the letters denoting the chemical element. This is usually not an integer, but a decimal.

    Remember that the periodic table shows the average atomic masses of the elements. As noted earlier, the relative atomic masses given for each element in the periodic table are the averages of the masses of all the isotopes of an 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 given 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.

Calculation of the atomic mass of an individual atom

  1. Find the atomic number of a given element or its isotope. The atomic number is the number of protons in an element's atoms and never changes. For example, all hydrogen atoms, and only they have one proton. Sodium has an atomic number of 11 because it has eleven protons, while oxygen has an atomic number of eight because it has eight protons. You can find the atomic number of any element in the periodic table of Mendeleev - in almost all of 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 a carbon atom. There are always six protons in carbon atoms, 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, the atomic mass of an element may appear to be twice its atomic number, it is never calculated by multiplying the atomic number by two.

  2. 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 particular element can often change, so the average atomic mass of an element is written as a decimal fraction 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 little to no effect on your calculations.

Calculating the relative atomic mass (atomic weight) of an element

  1. Determine which isotopes are in the sample. Chemists often determine the ratio of isotopes in a particular sample using a special instrument called a mass spectrometer. However, during training, this data will be provided to you in the conditions of tasks, control, 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.

  2. 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 a percentage. 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.

    • 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.

  3. Multiply the atomic mass of each isotope by its concentration in the sample. Multiply the atomic mass of each isotope by its percentage (expressed as a decimal). To convert percentages to decimals, 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 is 99% of the sample and carbon-13 is 1%, then multiply 12 (atomic mass of carbon-12) by 0.99 and 13 (atomic mass of carbon-13) by 0.01.
    • Reference books give percentages based on the known amounts of all the isotopes of an element. Most chemistry textbooks include this information in a table at the end of the book. For the sample under study, the relative concentrations of isotopes can also be determined using a mass spectrometer.

  4. Add up the results. Sum 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 an element is considered as a whole, and not a specific isotope of a given element, it is this value that is 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 decay 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.

From the lesson materials, you will learn that the atoms of some chemical elements differ from the atoms of other chemical elements in mass. The teacher will tell you how chemists measured the mass of atoms, which are so small that you can't even see them with an electron microscope.

Topic: Initial chemical ideas

Lesson: Relative atomic mass of chemical elements

At the beginning of the 19th century (150 years after the work of Robert Boyle), the English scientist John Dalton proposed a method for determining the mass of atoms of chemical elements. Let's consider the essence of this method.

Dalton proposed a model according to which a molecule of a complex substance contains only one atom of various chemical elements. For example, he believed that a water molecule consists of 1 hydrogen atom and 1 oxygen atom. The composition of simple substances according to Dalton also includes only one atom of a chemical element. Those. An oxygen molecule must consist of one oxygen atom.

And then, knowing the mass fractions of elements in a substance, it is easy to determine how many times the mass of an atom of one element differs from the mass of an atom of another element. Thus, Dalton believed that the mass fraction of an element in a substance is determined by the mass of its atom.

It is known that the mass fraction of magnesium in magnesium oxide is 60%, and the mass fraction of oxygen is 40%. Following the path of Dalton's reasoning, we can say that the mass of a magnesium atom is 1.5 times greater than the mass of an oxygen atom (60/40 = 1.5):

The scientist noticed that the mass of the hydrogen atom is the smallest, because. there is no complex substance in which the mass fraction of hydrogen would be greater than the mass fraction of another element. Therefore, he proposed to compare the masses of the atoms of the elements with the mass of the hydrogen atom. And in this way he calculated the first values ​​of the relative (relative to the hydrogen atom) atomic masses of chemical elements.

The atomic mass of hydrogen was taken as a unit. And the value of the relative mass of sulfur turned out to be 17. But all the values ​​\u200b\u200bobtained were either approximate or incorrect, because. the technique of the experiment of that time was far from perfect, and Dalton's installation on the composition of matter was incorrect.

In 1807 - 1817. Swedish chemist Jöns Jakob Berzelius did a great deal of research to refine the relative atomic masses of elements. He managed to get results close to modern ones.

Much later than the work of Berzelius, the masses of atoms of chemical elements began to be compared with 1/12 of the mass of a carbon atom (Fig. 2).

Rice. 1. Model for calculating the relative atomic mass of a chemical element

The relative atomic mass of a chemical element shows how many times the mass of an atom of a chemical element is greater than 1/12 of the mass of a carbon atom.

Relative atomic mass is denoted A r , it has no units of measurement, as it shows the ratio of the masses of atoms.

For example: A r (S) = 32, i.e. a sulfur atom is 32 times heavier than 1/12 the mass of a carbon atom.

The absolute mass of 1/12 of a carbon atom is a reference unit, the value of which is calculated with high accuracy and is 1.66 * 10 -24 g or 1.66 * 10 -27 kg. This reference mass is called atomic mass unit (a.u.m.).

The values ​​of the relative atomic masses of chemical elements do not need to be memorized, they are given in any textbook or reference book on chemistry, as well as in the periodic table of D.I. Mendeleev.

When calculating the values ​​of relative atomic masses, it is customary to round up to integers.

An exception is the relative atomic mass of chlorine - for chlorine, a value of 35.5 is used.

1. Collection of tasks and exercises in chemistry: 8th grade: to the textbook by P.A. Orzhekovsky and others. "Chemistry, Grade 8" / P.A. Orzhekovsky, N.A. Titov, F.F. Hegel. – M.: AST: Astrel, 2006.

2. Ushakova O.V. Chemistry workbook: 8th grade: to the textbook by P.A. Orzhekovsky and others. “Chemistry. Grade 8” / O.V. Ushakova, P.I. Bespalov, P.A. Orzhekovsky; under. ed. prof. P.A. Orzhekovsky - M .: AST: Astrel: Profizdat, 2006. (p. 24-25)

3. Chemistry: 8th grade: textbook. for general institutions / P.A. Orzhekovsky, L.M. Meshcheryakova, L.S. Pontak. M.: AST: Astrel, 2005.(§10)

4. Chemistry: inorg. chemistry: textbook. for 8 cells. general institutions / G.E. Rudzitis, FuGyu Feldman. - M .: Education, JSC "Moscow textbooks", 2009. (§§8,9)

5. Encyclopedia for children. Volume 17. Chemistry / Chapter. edited by V.A. Volodin, leading. scientific ed. I. Leenson. – M.: Avanta+, 2003.

Additional web resources

1. A single collection of digital educational resources ().

2. Electronic version of the journal "Chemistry and Life" ().

Homework

p.24-25 Nos. 1-7 from the Workbook in Chemistry: 8th grade: to the textbook by P.A. Orzhekovsky and others. “Chemistry. Grade 8” / O.V. Ushakova, P.I. Bespalov, P.A. Orzhekovsky; under. ed. prof. P.A. Orzhekovsky - M.: AST: Astrel: Profizdat, 2006.

What is "atomic mass"? What is the correct spelling of this word. Concept and interpretation.

Atomic mass The concept of this quantity underwent long-term changes in accordance with the change in the idea 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 some standard element. However, by about 1920 it became clear that the elements found in nature are of two types: some are actually identical atoms, while others have the same nuclear charge but different masses; such varieties 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 value 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 in 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 mass of the 16O isotope, and the corresponding scale was called physical. In the 1920s, natural oxygen was found to be a mixture of three isotopes: 16O, 17O, and 18O. In this regard, two problems arose. First, it turned out that the relative abundance of natural oxygen isotopes varies slightly, which means that the chemical scale is based on a quantity that is not an absolute constant. Secondly, physicists and chemists obtained different values ​​of 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 carbon isotope 12С (carbon unit). (1 amu, or 1D (dalton), in SI mass units is 1.66057×10-27 kg.) Natural carbon also consists of two isotopes: 12C - 99% and 13C - 1%, but the new values 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 12C 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 are distinguished by the fact that at one of the stages 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 is usually found by indirect methods. Let's consider both of these approaches. The atomic mass of Al has recently been determined as follows. Known amounts of Al were converted to nitrate, sulfate or hydroxide and then calcined to alumina (Al2O3) which was accurately quantified. From the ratio between the two known masses and the atomic masses of aluminum and oxygen (15.9)

Currently, the atomic mass unit is taken equal to 1/12 of the mass of a neutral atom of the most common isotope of carbon 12 C, so the atomic mass of this isotope is, by definition, exactly 12. The difference between the atomic mass of an isotope and its mass number is called the mass excess (usually expressed in MeV ). It can be both positive and 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 of the isotope 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.

Story

Until the 1960s, atomic mass was determined so that the nuclide oxygen-16 had an atomic mass of 16 (oxygen scale). However, the ratio of oxygen-17 to 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 a natural mixture of oxygen isotopes should have an atomic mass of 16, while physicists assigned the same number of 16 to the atomic mass of the most abundant oxygen isotope (having eight protons and eight neutrons).

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See what "Atomic mass" is in other dictionaries:

    The mass of an atom, expressed in atomic mass units. The atomic mass is less than the sum of the masses of the particles that make up the atom (protons, neutrons, electrons) by an amount determined by the energy of their interaction (see, for example, Mass defect) ... Big Encyclopedic Dictionary

    Atomic mass is the mass of an atom of a chemical element, expressed in atomic mass units (a.m.u.). For 1 amu 1/12 of the mass of a carbon isotope with an atomic mass of 12 is adopted. 1 amu = 1.6605655 10 27 kg. Atomic mass is made up of the masses of all protons and... Nuclear power terms

    atomic mass- is the mass of the element's atoms, expressed in atomic mass units. The mass of that amount of an element that contains the same number of atoms as 12 g of the 12C isotope. General chemistry: textbook / A. V. Zholnin ... Chemical terms

    ATOMIC MASS is a dimensionless quantity. A. m. atom mass chem. element, expressed in atomic units (see) ... Great Polytechnic Encyclopedia

    - (obsolete term atomic weight), the relative value of the mass of an atom, expressed in atomic mass units (amu). A. m. is less than the sum of the masses of the constituent atom h q per mass defect. A. m. was taken by D. I. Mendeleev for the main. characteristic of the element at ... ... Physical Encyclopedia

    atomic mass- — [Ya.N. Luginsky, M.S. Fezi Zhilinskaya, Yu.S. Kabirov. English-Russian Dictionary of Electrical Engineering and Power Industry, Moscow, 1999] Electrical engineering topics, basic concepts EN atomic weight ... Technical Translator's Handbook

    The mass of an atom, expressed in atomic mass units. For the atomic mass of a chemical element consisting of a mixture of isotopes, take the average value of the atomic mass of isotopes, taking into account their percentage (this value is given in the periodic ... ... encyclopedic Dictionary

    The concept of this quantity underwent long-term changes in accordance with the change in the idea 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 ... ... Collier Encyclopedia

    atomic mass- santykinė atominė masė statusas T sritis Standartizacija ir metrologija apibrėžtis Cheminio elemento vidutinės masės ir nuklido ¹²C atomo masės 1/12 dalies dalmuo. atitikmenys: engl. atomic mass; atomic weight; relative atomic mass vok. Atomasse …

    atomic mass- santykinė atominė masė statusas T sritis Standartizacija ir metrologija apibrėžtis Vidutinės elemento atomų masės ir 1/12 nuklido ¹²C atomo masės dalmuo. atitikmenys: engl. atomic mass; atomic weight; relative atomic mass vok. Atomasse, f;… … Penkiakalbis aiskinamasis metrologijos terminų žodynas


One of the fundamental properties of atoms is their mass. Absolute (true) mass of an atom- is extremely small. It is impossible to weigh atoms on a scale, because such exact scales do not exist. Their masses were determined by calculations.

For example, the mass of one hydrogen atom is 0.000,000,000,000,000,000,000,001,663 grams! The mass of an atom of uranium, one of the heaviest atoms, is approximately 0.000,000,000,000,000,000,000 4 grams.

The exact value of the 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 make 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 (to find the ratio of their masses). Since the determination of the relative atomic masses of the elements, different atoms have been used as a comparison. At one time, hydrogen and oxygen atoms were original 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 date, the benchmark 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 (a.m.u.) - the mass of 1/12 of a carbon atom

Let's compare how many times the absolute mass of a hydrogen atom and uranium differs from 1 amu, for this we divide these numbers one by one:

The values ​​obtained in the calculations and are the relative atomic masses of the elements - relatively 1/12 of the mass of a carbon atom.

So, the relative atomic mass of hydrogen is approximately equal to 1, and uranium - 238. Note that the relative atomic mass does not have units, as absolute mass units (grams) are canceled out when divided.

The relative atomic masses of all elements are indicated in the Periodic Table of chemical elements by D.I. Mendeleev. The symbol used to represent relative atomic mass is Ar (the letter r is an abbreviation for the word relative, which means relative).

Values ​​for the relative atomic masses of elements are used in many calculations. As a general rule, values ​​given in the Periodic System are rounded to whole numbers. Note that the elements in the Periodic Table are listed in order of increasing relative atomic masses.

For example, using the Periodic System, 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 the 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 integers, with the exception of chlorine, for which Ar = 35.5
  • Relative atomic mass has no units