One of the most important characteristics of atoms is their weight.

The absolute mass is the mass of an atom, expressed in kilograms (grams).

The absolute mass of an atom ( m a volume) is extremely small. Thus, an atom of the light isotope of hydrogen (protium) has a mass of 1.66 · 10–27 kg.

m(N) \u003d 1.66 10 -27 kg, m(H) \u003d 1.66 10 -24 g,

an atom of one of the oxygen isotopes has a mass of 2.67 10–26 kg,

m(O) \u003d 2.67 10 -26 kg, m(O) \u003d 2.67 10 -23 g,

an atom of the carbon isotope 12 C has a mass of 1.99 10 -26 kg,

m(C) \u003d 1.99 10 -26 kg, m(C) \u003d 1.99 10 -23 g.

In practical calculations, it is extremely inconvenient to use such quantities. Therefore, they usually use the values ​​not of the absolute masses of atoms, but the values relative atomic masses.

Relative atomic mass is denoted Ar, index r is the initial letter of the English word relative, which means relative.

As a unit for measuring the masses of atoms and molecules, atomic mass unit (a.m.u.).

An atomic mass unit (a.m.u.) is 1/12 of the mass of an atom of the carbon isotope 12 C, i.e.

a.u.m. == 1.99 10 -26 kg = 1.99 10 -23 g.

The relative atomic mass shows how many times the mass of an atom of a given element is greater than 1/12 of the mass of an atom of the carbon isotope 12 C, i.e., an atomic mass unit.

Relative atomic mass is a dimensionless quantity, but the designation of its value in atomic mass units (a.m.u.) is allowed. For example:

Thus, the value of the relative atomic mass of the element hydrogen is 1.001 or, rounded off,

Ar(H) ≈ 1 amu, and oxygen - Ar(O) = 15.999 ≈ 16 amu.

The values ​​of the relative atomic masses of the elements are given in the periodic system of D.I. Mendeleev. These values ​​are the average value of the mass of an atom of an element, taking into account the isotopes of this element existing in nature and their number. For ordinary calculations, rounded values ​​of the relative atomic masses of the elements should be used. (see Table 4 of the appendix).

Similarly to the concepts of the absolute mass of an atom and the relative atomic mass, one can formulate the concepts the absolute mass of the molecule and the relative molecular mass.

Absolute mass of a molecule(m) mol. - the mass of a chemical molecule, expressed in kilograms (grams).

Relative molecular weight(Mr) (or just molecular weight) - the mass of a molecule, expressed in atomic mass units.

Knowing the chemical formula of a compound, one can easily determine the value of its molecular weight, which is defined as the sum of the values ​​of the atomic masses of all the elements that make up the molecule of the substance.

For example, the relative molecular weight of sulfuric acid Mr(H 2 SO 4) will be the sum of two relative atomic masses of the hydrogen element, one relative atomic mass of the sulfur element, and four relative atomic masses of the oxygen element:

Mr (H 2 SO 4) \u003d 2Ar (H) + Ar (S) + 4Ar (O) \u003d 2 1 + 32 + 4 16 \u003d 98.

Thus, the value of the molecular weight of sulfuric acid is 98 or 98 amu.

Molecular weight (relative molecular weight) shows how many times the mass of a molecule of a given substance is greater than 1/12 of the mass of a carbon atom 12 C.

In the example above, the molecular weight of sulfuric acid is 98 amu, i.e. the mass of sulfuric acid is 98 times greater than 1/12 of the mass of the carbon atom 12 C .

Relative atomic and relative molecular weight. Moth. Avogadro's number

Modern research methods make it possible to determine extremely small masses of atoms 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 atomic mass unit (abbreviated a.m.u.) was adopted as a unit of atomic mass, which is 1/12 of the mass of an atom of the carbon isotope "C". Most chemical elements have atoms with different masses. Therefore, the relative atomic mass A, 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 the carbon atom 12C. The relative atomic masses of the elements are denoted A, where the index r is the initial letter of the English word relative - relative. The entries Ar(H), Ar(0), Ar(C) mean: the atomic mass of hydrogen, the atomic mass of oxygen, and the atomic mass of carbon. For example, Ar(H) = 1.6747x 10-27 = 1.0079; 1/12 x 1.993 x 10 -26

Relative atomic mass is one of the main characteristics of a chemical element. The relative molecular weight 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 12C carbon atom. Instead of the term "attributes atomic mass", the term "atomic mass" can be used. The relative molecular weight 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 by 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) \u003d 18.01534 This means that the ratio of the molecular weight of water is 18.01534, rounded, 18. The ratio of the molecular weight shows how much the mass of a molecule of a given substance is more than 1/12 of the mass of an atom C +12. So, 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 a C +12 atom. Molecular weight refers to one of the main characteristics of a substance. Moth. Molar mass. The International System of Units (SI) uses the mole as the unit of quantity of a substance. 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 the carbon isotope C +12. Knowing the mass of one carbon atom (1.993 10-26 kg), you can calculate the number of NA atoms in 0.012 kg of carbon: NA \u003d 0.012 kg / mol \u003d 1.993 x 10-26 kg 6.02 x 1023 units / mol.

This number is called the Avogadro 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 a substance. It has the units of kg/mol or g/mol; usually it is denoted by the letter M. The molar mass of a substance is easy to calculate, knowing the mass of the molecule. So, if the mass of a water molecule is 2.99x10-26, kg, then the molar mass Mr (H2O) \u003d 2.99 10-26 kg 6.02 1023 1 / mol \u003d 0.018 kg / mol, or 18 g / mol. In general, the molar mass of a substance, expressed in g/mol, is numerically equal to the relative atomic or relative molecular mass of that 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 / mol. Molar mass can be calculated for substances in both molecular and atomic states. For example, the relative molecular mass of hydrogen Mr (H 2) \u003d 2, and the atomic mass of hydrogen A (H) \u003d 1 refers. The amount of 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 the number of these particles equal to the Avogadro constant, for example

1 mole of C atoms +12 = 6.02 1023 C atoms +12

1 mol of H 2 O molecules \u003d 6.02 1023 H 2 O molecules

1 mol 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 a substance is expressed in moles. There are simple relationships between the mass of a substance (t, g), the amount of a substance (p, 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 the amount of a substance, or to determine the amount of a substance in a known assay of it, or to find the molar mass of a substance.

Atomic-molecular theory. Atom, molecule. Chemical element. Simple and complex matter. Allotropy.

Chemistry- the science of substances, the patterns of their transformations (physical and chemical properties) and applications. 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 composition and properties, while the composition of the nuclei of atoms does not change.

Physical phenomena: the physical state of substances changes (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 its chemical properties. Different elements correspond to different atoms.

3. Molecules and atoms are in continuous motion; between them there are forces of attraction and repulsion.

Chemical element- this is a type of atom, characterized by certain charges of the nuclei and the structure of the electron 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 conditional record of the composition of a substance using chemical signs (proposed in 1814 by J. Berzelius) and indices (the index is the number to the right below the symbol. It indicates the number of atoms in the molecule). The chemical formula shows which atoms of which elements and in what relation are interconnected in a molecule.

Allotropy- the phenomenon of the formation by a chemical element of several simple substances that differ in structure and properties.

Simple substances Molecules are made up of atoms of the same element.

Compound substances Molecules are made up of atoms of various chemical elements.

The international unit of atomic mass is equal to 1/12 of the mass of the 12 C isotope - the main isotope of natural carbon: 1 amu \u003d 1/12 m (12 C) \u003d 1.66057 10 -24 g

Relative atomic mass (ar)- a dimensionless value equal to the ratio of the average mass of an element atom (taking into account the percentage of isotopes in nature) to 1/12 of the mass of a 12 C atom.

Average absolute mass of an atom (m) is equal to the relative atomic mass times the a.m.u. (1 a.m.u. = 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 of the mass of a carbon atom 12 C.

Mr = mr / (1/12 ma(12 C))

mr is the mass of the molecule of the 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 the 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 - moth.

Amount of substance, mol. Means a certain number of structural elements (molecules, atoms, ions). Denoted n, measured in moles. A mole is the amount of a substance that contains as many particles as there are atoms in 12 g of carbon.

Avogadro di Quaregna number(NA). The number of particles in 1 mole of any substance is the same and equal to 6.02 10 23. (The Avogadro 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 value has the dimension g/mol, and the second is dimensionless: M = N A m(1 molecule) = N A Mr 1 a.m.u. = (N A 1 amu) Mr = Mr

Equivalent is a real or conditional particle of matter, which is equivalent to:
a) one ion H + or OH - in this acid-base reaction;

b) one electron in a given OVR (redox reaction);

c) one unit of charge in a given exchange reaction,

d) the number of monodentate ligands involved in the reaction of complex formation.

Atoms are very small and have very little mass. If we express the mass of an atom of any chemical element in grams, then this will be a number preceded by more than twenty zeros after the decimal point. Therefore, it is inconvenient to measure the mass of atoms in grams.

However, if we take any very small mass as a unit, then all other small masses can be expressed as a ratio to this unit. 1/12 of the mass of a carbon atom was chosen as the unit for measuring the mass of an atom.

1/12 of the mass of a carbon atom is called atomic mass unit(a.e.m.).

Relative atomic mass is a value equal to the ratio of the real mass of an atom of a particular chemical element to 1/12 of the real mass of a carbon atom. This is a dimensionless quantity, since two masses are divided.

A r = m at. / (1/12)m arc.

However absolute atomic mass is relative in value and has the unit a.u.m.

That is, the relative atomic mass shows how many times the mass of a particular atom is greater than 1/12 of a carbon atom. If the A atom has r = 12, then its mass is 12 times greater than 1/12 of the mass of a carbon atom, or, in other words, it has 12 atomic mass units. This can only happen to carbon itself (C). The hydrogen atom (H) has Ar = 1. This means that its mass is equal to the mass of 1/12 of the mass of the carbon atom. Oxygen (O) has a relative atomic mass of 16 amu. This means that an oxygen atom is 16 times more massive than 1/12 of a carbon atom, it has 16 atomic mass units.

The lightest element is hydrogen. Its mass is approximately equal to 1 amu. The heaviest atoms have a mass approaching 300 amu.

Usually for each chemical element its value is the absolute mass of atoms, expressed in terms of a. e. m. are rounded up.

The value of atomic mass units is recorded in the periodic table.

For molecules, the concept is used relative molecular weight (Mr). Relative molecular weight shows how many times the mass of a molecule is greater than 1/12 of the mass of a carbon atom. But since the mass of a molecule is equal to the sum of the masses of its constituent atoms, the relative molecular mass can be found by simply adding the relative masses of these atoms. For example, a water molecule (H 2 O) contains two hydrogen atoms with Ar = 1 and one oxygen atom with Ar = 16. Therefore, Mr(H 2 O) = 18.

A number of substances have a non-molecular structure, such as metals. In such a case, their relative molecular weight is considered equal to their relative atomic weight.

In chemistry, an important quantity is called mass fraction of a chemical element in a molecule or substance. It shows what part of the relative molecular weight is accounted for by a given element. For example, in water, hydrogen accounts for 2 shares (since there are two atoms), and oxygen for 16. That is, if you mix hydrogen with a mass of 1 kg and oxygen with a mass of 8 kg, they will react without residue. The mass fraction of hydrogen is 2/18 = 1/9, and the mass fraction of oxygen is 16/18 = 8/9.

Atomic and molecular weights

BASIC CHEMICAL CONCEPTS AND LAWS. AGGREGATE STATES OF SUBSTANCES

Chemistry - the science of substances and their transformations

Substance- a type of matter consisting of discrete particles with a rest mass (atoms, molecules, ions). The mode of existence of matter traffic .

The fundamental law of nature is law of indestructibility of matter and motion results law of conservation of mass opened by M.V. Lomonosov in 1748 and published in 1760: the mass of the substances that have entered into the reaction is equal to the mass of the substances formed as a result of the reaction.

Atomic-molecular doctrine

M.V. Lomonosov is also the creator of the atomic and molecular doctrine, which he formulated in 1741.

The main provisions of the atomic and molecular doctrine:

1) All substances consist of molecules, between which there are gaps. Molecule - the smallest particle of a substance that has its chemical properties.

2) Molecules are made up of atoms that are connected to each other in certain ratios.

Atom- the smallest particle of a chemical element in the composition of simple and complex substances, chemically indivisible.

3) Molecules and atoms are in continuous motion.

4) Atoms are characterized by a certain mass and size.

5) Different elements correspond to different atoms ( element is the type of atom).

6) Molecules of simple substances consist of identical atoms, and molecules of complex substances consist of different ones.

Law of constancy of composition

The discovery of the law of conservation of mass marked the transition of chemistry to quantitative research methods. The composition of many substances was studied and the law of composition constancy was established in 1799-1807. J. Proust : any pure substance, regardless of the methods of its production and presence in nature, has a constant qualitative and quantitative composition.

Law of simple multiple ratios

It follows from the law of composition constancy that when a complex substance is formed, the elements combine with each other in certain weight ratios. Many elements can combine with each other in several different weight ratios and in this case different substances are formed (CO, CO 2). In CO and CO 2, N 2 O, NO and NO 2 molecules, the composition changes in jumps, and not gradually, which indicates a discrete structure of the substance. This law, confirmed by experience, was the first proof reality of the existence of atoms.

Atomic and molecular weights

Atoms and molecules have absolute masses of the order of 10–24–10–21 g, which are inconvenient for comparison, so chemists use the relative masses of atoms. The concept of relative atomic mass was introduced by J. Dalton in 1803. He took the mass of the lightest atom, hydrogen, as a unit of mass. At present, the mass of 1/12 of the mass of the carbon atom of the 12 C isotope, equal to 1.66043 × 10 -24 g, is accepted as a standard.

Relative atomic (BUT r) weight shows how many times a given atom is heavier than 1/12 of the mass of an atom of the carbon isotope 12 C.

Using the value specific heat, which is easily determined experimentally ( the ratio of the amount of heat received or given 1 g of a substance to the corresponding change in temperature) you can find an approximate value of the atomic mass. The exceptions are light elements, especially non-metals, which have a much lower heat capacity (beryllium, boron, silicon, diamond).

At present, the atomic masses of elements are determined by mass spectroscopy. The masses of atoms are determined by the deviation of the trajectory of their ions moving in a magnetic field, since the magnitude of the deviation depends on the ratio of the mass of the ion to its charge.

Relative molecular weight (M r) shows how many times a given molecule is heavier than 1/12 of the mass of a 12 C atom.

,

(1.4)

where m m is the mass of the molecule.