To measure the mass of an atom, relative atomic mass is used, which is expressed in atomic mass units (amu). Relative molecular weight is made up of the relative atomic masses of substances.
Concepts
To understand what relative atomic mass is in chemistry, you should understand that the absolute mass of an atom is too small to be expressed in grams, much less in kilograms. Therefore, in modern chemistry, 1/12 of the mass of carbon is taken as an atomic mass unit (amu). Relative atomic mass is equal to the ratio of the absolute mass to 1/12 of the absolute mass of carbon. In other words relative mass reflects how many times the mass of an atom of a particular substance exceeds 1/12 the mass of a carbon atom. For example, the relative mass of nitrogen is 14, i.e. The nitrogen atom contains 14 a. e.m. or 14 times more than 1/12 of a carbon atom.
Rice. 1. Atoms and molecules.
Among all the elements, hydrogen is the lightest, its mass is 1 unit. The heaviest atoms have a mass of 300 a. e.m.
Molecular mass is a value indicating how many times the mass of a molecule exceeds 1/12 of the mass of carbon. Also expressed in a. e.m. The mass of a molecule is made up of the mass of atoms, therefore, to calculate the relative molecular mass it is necessary to add up the masses of the atoms of the substance. For example, the relative molecular weight of water is 18. This value is the sum of the relative atomic masses of two hydrogen atoms (2) and one oxygen atom (16).
Rice. 2. Carbon in the periodic table.
As you can see, these two concepts have several common characteristics:
- the relative atomic and molecular masses of a substance are dimensionless quantities;
- relative atomic mass is designated Ar, molecular mass - Mr;
- The unit of measurement is the same in both cases - a. e.m.
Molar and molecular masses are the same numerically, but differ in dimension. Molar mass is the ratio of the mass of a substance to the number of moles. It reflects the mass of one mole, which equal to the number Avogadro, i.e. 6.02 ⋅ 10 23 . For example, 1 mole of water weighs 18 g/mol, and M r (H 2 O) = 18 a. e.m. (18 times heavier than one atomic mass unit).
How to calculate
To express relative atomic mass mathematically, one should determine that 1/2 part of carbon or one atomic mass unit is equal to 1.66⋅10 −24 g. Therefore, the formula for relative atomic mass is as follows:
A r (X) = m a (X) / 1.66⋅10 −24,
where m a is the absolute atomic mass of the substance.
The relative atomic mass of chemical elements is indicated in the periodic table, so you do not need to calculate it yourself when solving problems. Relative atomic masses are usually rounded to whole numbers. The exception is chlorine. The mass of its atoms is 35.5.
It should be noted that when calculating the relative atomic mass of elements that have isotopes, their average value is taken into account. Atomic mass in this case is calculated as follows:
A r = ΣA r,i n i ,
where A r,i is the relative atomic mass of isotopes, n i is the content of isotopes in natural mixtures.
For example, oxygen has three isotopes - 16 O, 17 O, 18 O. Their relative mass is 15.995, 16.999, 17.999, and their content in natural mixtures is 99.759%, 0.037%, 0.204%, respectively. Dividing the percentages by 100 and substituting the values, we get:
A r = 15.995 ∙ 0.99759 + 16.999 ∙ 0.00037 + 17.999 ∙ 0.00204 = 15.999 amu
Referring to the periodic table, it is easy to find this value in the oxygen cell.
Rice. 3. Periodic table.
Relative molecular mass is the sum of the masses of the atoms of a substance:
When determining the relative molecular weight value, symbol indices are taken into account. For example, calculating the mass of H 2 CO 3 is as follows:
M r = 1 ∙ 2 + 12 + 16 ∙ 3 = 62 a. e.m.
Knowing the relative molecular mass, it is possible to calculate the relative density of one gas from the second, i.e. determine how many times one gaseous substance is heavier than the second. To do this, use the equation D (y) x = M r (x) / M r (y).
What have we learned?
From the 8th grade lesson we learned about relative atomic and molecular mass. The unit of relative atomic mass is taken to be 1/12 of the mass of carbon, equal to 1.66⋅10 −24 g. To calculate the mass, it is necessary to divide the absolute atomic mass of the substance by the atomic mass unit (amu). The value of the relative atomic mass is indicated in the periodic table of Mendeleev in each cell of the element. The molecular mass of a substance is the sum of the relative atomic masses of the elements.
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Relative atomic and relative molecular mass. Mol. Avogadro's number
Modern research methods make it possible to determine extremely small atomic masses with great accuracy. So, for example, the mass of a hydrogen atom is 1.674 10 27 kg, oxygen - 2.667 x 10 -26 kg, carbon - 1.993 x 10 26 kg. In chemistry, not absolute values of atomic masses are traditionally used, but relative ones. In 1961, the unit of atomic mass was adopted as the atomic mass unit (abbreviated a.m.u.), which is ‘/12th part of the mass of an atom of the carbon isotope “C.” Most chemical elements have atoms with different masses. Therefore, the relative atomic mass of a chemical element is a value equal to the ratio of the average mass of an atom of the natural isotopic composition of the element to 1/12 of the mass of a carbon atom 12C. The relative atomic masses of elements are denoted by A, where the subscript r is the initial letter English word relative - relative. The entries Ar(H), Ar(0), Ar(C) mean: relative atomic mass of hydrogen, relative atomic mass of oxygen, relative atomic mass of carbon. For example, Ar(H) = 1.6747x 10-27 = 1.0079; 1/12 x 1.993 x 10 -26Relative atomic mass is one of the main characteristics chemical element. The relative molecular mass M of a substance is a value equal to the ratio of the average mass of a molecule of the natural isotopic composition of a substance to 1/12 of the mass of a carbon atom 12C. Instead of the term "relates atomic mass", the term "atomic mass" can be used. The relative molecular mass is numerically equal to the sum of the relative atomic masses of all the atoms that make up the molecule of the substance. It is easily calculated using the formula of the substance. For example, Mg(H2O) is composed of 2Ar(H) = 2 1.00797 = 2.01594 Ar(0) = 1x15, 9994 = 15.9994
Mr (H2O) = 18.01534 So the relative molecular weight of water is 18.01534, rounded to 18. The relative molecular weight shows how much the mass of the molecule is of this substance more than 1/12 of the mass of the C +12 atom. Thus, the molecular weight of water is 18. This means that the mass of a water molecule is 18 times greater than 1/12 of the mass of the C +12 atom. Molecular mass is one of the main characteristics of a substance. Mol. Molar mass. In the International System of Units (SI), the unit of quantity of a substance is the mole. A mole is the amount of a substance containing as many structural units (molecules, atoms, ions, electrons and others) as there are atoms in 0.012 kg of carbon isotope C +12. Knowing the mass of one carbon atom (1.993 10-26 kg), we can calculate the number of NA atoms in 0.012 kg of carbon: NA = 0.012 kg/mol = 1.993 x10-26 kg 6.02 x 1023 units/mol.
This number is called Avogadro's constant (designation HA dimension 1/mol), shows the number of structural units in a mole of any substance. Molar mass is a value equal to the ratio of the mass of a substance to the amount of substance. It has the dimension kg/mol or g/mol; it is usually denoted by the letter M. The molar mass of a substance is easy to calculate if you know the mass of the molecule. So, if the mass of a water molecule is 2.99x10-26, kg, then the molar mass of Mr (H2O) = 2.99 10-26 kg 6.02 1023 1/mol = 0.018 kg/mol, or 18 g/mol. IN general case the molar mass of a substance, expressed in g/mol, is numerically equal to the relative atomic or relative molecular mass of this substance. -For example, the relative atomic and molecular masses of C, Fe, O, H 2O are respectively 12, 56, 32.18, and their molar masses are respectively 12 g/mol, 56 g/mol, 32 g/mol, 18 g/ mole. Molar mass can be calculated for substances in both molecular and atomic states. For example, the relative molecular mass of hydrogen is Mr (H 2) = 2, and the relative atomic mass of hydrogen is A (H) = 1. The amount of the substance, determined by the number of structural units (H A), is the same in both cases - 1 mol. However, the molar mass of molecular hydrogen is 2 g/mol, and the molar mass of atomic hydrogen is 1 g/mol. One mole of atoms, molecules or ions contains a number of these particles equal to Avogadro's constant, for example
1 mole of C +12 atoms = 6.02 1023 C +12 atoms
1 mole of H 2 O molecules = 6.02 1023 H 2 O molecules
1 mole of S0 4 2- ions = 6.02 1023 S0 4 2- ions
Mass and quantity of a substance are different concepts. Mass is expressed in kilograms (grams), and the amount of substance is expressed in moles. There are simple relationships between the mass of a substance (t, g), the amount of substance (n, mol) and the molar mass (M, g/mol): m=nM, n=m/M M=m/n Using these formulas it is easy to calculate the mass of a certain amount of a substance, or determine the amount of a substance in a known quantity of it, or find the molar mass of a substance.
In chemical literature, a situation has historically developed when the masses of atoms and molecules are expressed through the concepts of atomic weight and molecular weight.
As is known, if a body with mass m (as well as an atom or molecule) moves under the influence of the Earth’s gravity with acceleration g, then the force of gravity of this body is equal, i.e., the force of gravity is proportional to the mass of the body on which it acts.
If the body is at rest, then the weight of the body is equal to the force of gravity acting on it, and in the formula we can consider P as the weight of the body. Consequently, for bodies at rest, their weights are proportional to their masses. However, the acceleration at various points earth's surface different, therefore the weight of the same body (atom, molecule!) will be different here. The body's weight will also decrease as it rises above the Earth's surface.
In conclusion, let us ask ourselves the question: Is the weight of a body (and, accordingly, the weight of an atom, molecule) the same on Earth, in space? orbital station, on the surface of the Moon?
If necessary, you can repeat it physical concepts“weight”, “mass”, etc..
Relative units to express the weight of atoms were first used by Dalton, who defined atomic weight as a number showing how many times an atom of an element is heavier than an atom of another element. As a unit of atomic weights, he proposed the weight of the lightest atom - hydrogen.
More correctly, as shown above, we need to talk about a unit of atomic or molecular masses, therefore, in further presentation, the authors tried to use these concepts everywhere instead of “atomic weight”, “molecular weight”.
Since the atomic masses of elements were calculated from experimental data on the weight ratios in various compounds, and oxygen forms compounds with much a large number elements compared to hydrogen, then in subsequent years, until 1961, part of the mass of the oxygen atom was adopted as a unit of atomic mass. This relative unit of measure of the mass of atoms was called the oxygen unit (o.u.).
However, by 1930 it was discovered that in addition to oxygen atoms with a mass of 16 k.u., there are isotopes of oxygen that differ in mass (0.039%) and (0.204%). The chemical properties of oxygen isotopes are the same, but the physical properties, although not very much, differ, therefore the isotopic composition of oxygen in different natural compounds is not the same. For example, the average atomic mass atmospheric oxygen 0.00011 atomic units less than the average atomic mass of oxygen from seawater.
A physical and chemical system of atomic mass units emerged. Physicists took part of the mass of an isotope as a unit of atomic mass, while chemists took part of the average mass of an oxygen atom of natural isotopic composition. This led to different values of atomic masses and made it difficult to compare physical and chemical atomic masses, which was ultimately the main reason for the abandonment of the oxygen atomic unit.
In 1961, the International Union of Pure and Applied Chemistry decided to choose a standard unit of atomic mass and move to a unified atomic mass scale. The carbon unit (cu) was chosen as the new standard unit of atomic mass - part of the mass of the carbon isotope. Atomic masses based on the new unit (cu) are equal to the old ones (cu) multiplied by 0.99996, so that changes in previous atomic masses are so small, and this should be especially emphasized, that they do not affect almost all chemical calculations.
Thus, the mass of an atom expressed in carbon units is called atomic mass. Atomic mass shows how many times the mass of an atom of a given element is heavier than the mass of a C10 carbon atom. The mass of molecules is also expressed in carbon units (cu).
The molecular mass of a substance is the mass of its molecule, expressed in carbon units. Molecular mass shows how many times the mass of a molecule of a given substance is heavier than the mass of carbon C12. Therefore, both atomic and molecular masses are relative units of measurement. When writing, they usually do not indicate the dimensions of atomic and molecular masses, remembering that they are expressed in carbon units.
For quantitative calculations, it is convenient to use the following characteristics - gram-atom and gram-molecule.
A gram atom is the number of grams of a substance that is numerically equal to the atomic mass of that element. For example, the atomic mass of sodium is 23 cu. i.e., therefore, G-sodium atom has a mass of 23 g.
The number of grams of a substance, numerically equal to its molecular weight, is called a gram molecule of this substance, or mole. For example, the molecular weight of potassium permanganate is 158 c.u. e., therefore, constitute 1 gram molecule.
The concepts of atomic and molecular masses are fundamentally different from the concepts of gram-atomic and gram-molecular masses. If the values of atomic and molecular masses are relative numbers and show how many times the mass of an atom or molecule is larger than part of an atom of a carbon isotope, then gram-atom and gram-molecule are absolute numbers, showing the number of grams of a substance.
After the discovery of Avogadro's law (see § 5, Chapter IV) "it was proven that the number of molecules (atoms) contained in one gram-molecule (gram-atom) of any substance is the same and equal (Avogadro's number), i.e. . and the mass of a gram molecule is equal to the mass of molecules of a given substance. It is worth emphasizing that there are molecules (atoms) in 1 mole.
(1 g-atom) - any substance in any state of aggregation - solid, liquid, gaseous.
Atomic-molecular theory. Atom, molecule. Chemical element. Simple and compound. Allotropy.
Chemistry- the science of substances, the laws of their transformations (physical and chemical properties) and application. Currently, more than 100 thousand inorganic and more than 4 million organic compounds are known.
Chemical phenomena: Some substances are transformed into others that differ from the original ones in composition and properties, while the composition of the atomic nuclei does not change.
Physical phenomena: is changing physical condition substances (vaporization, melting, electrical conductivity, release of heat and light, malleability, etc.) or new substances are formed with a change in the composition of atomic nuclei.
1. All substances are made up of molecules. Molecule- the smallest particle of a substance that has its chemical properties.
2. Molecules are made up of atoms. Atom- the smallest particle of a chemical element that retains all of it chemical properties. Various elements different atoms correspond.
3. Molecules and atoms are in continuous motion; there are forces of attraction and repulsion between them.
Chemical element- this is a type of atoms characterized by certain nuclear charges and the structure of electronic shells. Currently, 117 elements are known: 89 of them are found in nature (on Earth), the rest are obtained artificially. Atoms exist in a free state, in compounds with atoms of the same or other elements, forming molecules. The ability of atoms to interact with other atoms and form chemical compounds is determined by its structure. Atoms consist of a positively charged nucleus and negatively charged electrons moving around it, forming an electrically neutral system that obeys the laws characteristic of microsystems.
Chemical formula- this is a conventional notation of the composition of a substance using chemical symbols (proposed in 1814 by J. Berzelius) and indices (index is the number at the bottom right of the symbol. Indicates the number of atoms in the molecule). The chemical formula shows which atoms of which elements and in what ratio are connected to each other in a molecule.
Allotropy- the phenomenon of formation by a chemical element of several simple substances, differing in structure and properties.
Simple substances- molecules consist of atoms of the same element.
Complex substances- molecules consist of atoms of various chemical elements.
The international unit of atomic mass is equal to 1/12 of the mass of the isotope 12 C - the main isotope of natural carbon: 1 amu = 1/12 m (12 C) = 1.66057 10 -24 g
Relative atomic mass (Ar)- a dimensionless quantity equal to the ratio of the average mass of an atom of an element (taking into account the percentage of isotopes in nature) to 1/12 of the mass of a 12 C atom.
Average absolute atomic mass (m) equal to the relative atomic mass times the amu. (1 amu=1.66*10 -24)
Relative molecular weight (Mr)- a dimensionless quantity showing how many times the mass of a molecule of a given substance is greater than 1/12 the mass of a carbon atom 12 C.
Mr = mr / (1/12 mа(12 C))
mr is the mass of a molecule of a given substance;
ma(12 C) - mass of carbon atom 12 C.
Mr = S Ar(e). The relative molecular mass of a substance is equal to the sum of the relative atomic masses of all elements, taking into account formula indices.
The absolute mass of a molecule is equal to the relative molecular mass times the amu. The number of atoms and molecules in ordinary samples of substances is very large, therefore, when characterizing the amount of a substance, a special unit of measurement is used - mole.
Amount of substance, mol. Means a certain number of structural elements (molecules, atoms, ions). Denoted n and measured in moles. A mole is the amount of a substance containing as many particles as there are atoms in 12 g of carbon.
Avogadro di Quaregna number(N A). The number of particles in 1 mole of any substance is the same and equals 6.02 10 23. (Avogadro's constant has the dimension - mol -1).
Molar mass shows the mass of 1 mole of a substance (denoted by M): M = m/n
The molar mass of a substance is equal to the ratio of the mass of the substance to the corresponding amount of the substance.
The molar mass of a substance is numerically equal to its relative molecular mass, however, the first quantity has the dimension g/mol, and the second is dimensionless: M = N A m(1 molecule) = N A Mr 1 amu. = (N A 1 amu) Mr = Mr
Equivalent- is a real or conditional particle of a substance that is equivalent to:
a) one H + or OH - ion in a given acid-base reaction;
b) one electron in a given ORR (redox reaction);
c) one unit of charge in a given exchange reaction,
d) the number of monodentate ligands participating in the complex formation reaction.
