Any body has mass. Let's take a body like a bag of apples. This body has mass. Its mass will be the sum of the mass of each apple in the bag. A bag of rice also has its mass, which is determined by adding the mass of all the rice grains, although they are very small and light.

All bodies are made up of matter. The mass of a body is made up of the mass of its constituent substances. Substances, in turn, consist of particles, molecules or atoms, therefore, the particles of a substance also have mass.

§ 2. Atomic mass unit

The mass of an oxygen atom is about sixteen times greater and is 2.66∙ 10 -23 g, mass of a carbon atom 1.99∙ 10 -23 g. The mass of an atom is denoted -m a.

In this case, the mass of a hydrogen atom will be 1 amu, the mass of an oxygen atom will be 16 amu, and the mass of a carbon atom will be 12 amu.

Chemists for a long time had no idea how much one atom of any element weighs in the usual and convenient units for us to measure mass (grams, kilograms, etc.).

Attempts were made to determine how many times the atoms of some elements are heavier than others. Thus, scientists sought to compare the mass of an atom of one element with the mass of an atom of another element.

The solution of this problem was also fraught with great difficulties, and above all with the choice of a standard, i.e. chemical element, relative to which the comparison of the atomic masses of the remaining elements should be carried out.

§3 . Relative atomic mass

Scientists of the 19th century solved this problem on the basis of experimental data to determine the composition of substances. The lightest atom, the hydrogen atom, was taken as the standard. Experimentally, it was found that the oxygen atom is 16 times heavier than the hydrogen atom, i.e., its relative mass(relative to the mass of the hydrogen atom) is 16.

This value was denoted by the letters A r (index " r " - from the initial letter English word"relative" - ​​relative). Thus, the record of the value of the relative atomic masses of chemical elements should look like this: the relative atomic mass of hydrogen is 1, the relative atomic mass of oxygen is 16, the relative atomic mass of carbon is 12.

The relative atomic mass shows how many times the mass of an atom of one chemical element is greater than the mass of an atom, which is the standard, therefore this value has no dimension.

As already mentioned, initially the values ​​of atomic masses were determined in relation to the mass of the hydrogen atom. Later, the standard for determining atomic masses was 1/12 of the mass of a carbon atom (a carbon atom is 12 times heavier than a hydrogen atom).

Relative atomic mass of an element ( A r ) is the ratio of the mass of an atom of a chemical element to 1/12 of the mass of a carbon atom.

The values ​​of the atomic masses of chemical elements are given in the Periodic system of chemical elements by D.I. Mendeleev. Take a look at the periodic table and consider any of its cells, for example, at number 8.

Under chemical sign and the name in the bottom line indicates the value atomic mass chemical element:the relative atomic mass of oxygen is 15.9994. Please note: the relative atomic masses of almost all chemical elements have a fractional value. The reason for this is the existence of isotopes. Let me remind you that isotopes are called atoms of the same chemical element, slightly different in mass.

At school, calculations usually use the values ​​\u200b\u200bof relative atomic masses, rounded to whole numbers. But in a few cases fractional values ​​are used, for example: the relative atomic mass of chlorine is 35.5.

§ 4. Relative molecular weight

The masses of atoms add up to the mass of a molecule.

The relative molecular weight of a substance is a number showing how many times the mass of a molecule of this substance is greater than 1/12 of the mass of a carbon atom.

Relative molecular weight is denoted by −M r

The relative molecular weight of substances is calculated from chemical formulas expressing the composition of substances. To find the relative molecular weight, it is necessary to sum up the values ​​of the relative atomic masses of the elements that make up the molecule of a substance, taking into account the quantitative composition, i.e., the number of atoms of each element (in chemical formulas it is expressed using indices). For example, the relative molecular weight of water, which has the formula H2O , is equal to the sum of two values ​​of the relative

atomic mass of hydrogen and one value of the relative atomic mass of oxygen:

Relative molecular weight of sulfuric acid having the formula H 2 SO 4 is equal to the sum

two values ​​of the relative atomic mass of hydrogen, one value of the relative atomic mass of sulfur, and four values ​​of the relative atomic mass of oxygen:.

Relative molecular weight is a dimensionless quantity. It should not be confused with the true mass of molecules, expressed in atomic mass units.

Literature:

NOT. Kuznetsova. Chemistry. 8th grade. Tutorial for educational institutions. – M. Ventana-Graf, 2012.

Sources used for visual design:

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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 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, because atomic masses expressed in these units are extremely small, they are often written in unified atomic mass units, or a.m.u. 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 tables Mendeleev contains 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 decimal, just 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 atomic nucleus. Such isotopes are called radioactive.

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 the atomic theory Dalton did, who compiled a table of atomic masses, taking the mass of a hydrogen atom as a unit.

Until 1961, in physics, 1/16 of the mass of an oxygen atom 16 O was taken as an atomic mass unit (abbreviated amu), and in chemistry - 1/16 of the average atomic mass of natural oxygen, which is a mixture of three isotopes. The chemical mass unit was 0.03% larger than the physical one.

At present, a unified measurement system has been adopted in physics and chemistry. 1/12 of the mass of the carbon atom 12 C is chosen as the standard unit of atomic mass.

1 amu \u003d 1/12 m (12 C) \u003d 1.66057 × 10 -27 kg \u003d 1.66057 × 10 -24 g.

DEFINITION

Relative atomic mass of an element (A r)- this is a dimensionless quantity equal to the ratio of the average mass of an element atom to 1/12 of the mass of an atom 12 C.

When calculating the relative atomic mass, the abundance of isotopes of elements in the earth's crust is taken into account. For example, chlorine has two isotopes 35 Cl (75.5%) and 37 Cl (24.5%). The relative atomic mass of chlorine is:

A r (Cl) \u003d (0.755 × m (35 Cl) + 0.245 × m (37 Cl)) / (1/12 × m (12 C) = 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 times the amu:

m(Cl) = 35.5 × 1.66057 × 10 -24 = 5.89 × 10 -23 g.

Examples of problem solving

EXAMPLE 1

EXAMPLE 2