Atomic-molecular science
The idea of atoms as the smallest indivisible particles originated in ancient Greece. The foundations of modern atomic-molecular science were first formulated by M.V. Lomonosov (1748), but his ideas, set out in a private letter, were unknown to most scientists. Therefore, the founder of modern atomic-molecular science is considered to be the English scientist J. Dalton, who formulated (1803–1807) its main postulates.
1. Each element consists of very small particles - atoms.
2. All atoms of one element are the same.
3. Atoms of different elements have different masses and have different properties.
4. Atoms of one element do not transform into atoms of other elements as a result of chemical reactions.
5. Chemical compounds are formed by the combination of atoms of two or more elements.
6. In a given compound, the relative amounts of atoms of different elements are always constant.
These postulates were initially indirectly proven by a set of stoichiometric laws. Stoichiometry - part of chemistry that studies the composition of substances and its changes during chemical transformations. This word is derived from the Greek words “stoechion” - element and “metron” - measure. The laws of stoichiometry include the laws of conservation of mass, constancy of composition, multiple ratios, volume ratios, Avogadro's law and the law of equivalents.
1.3. Stoichiometric laws
The laws of stoichiometry are considered components of the AMU. Based on these laws, the concept of chemical formulas, chemical equations and valence was introduced.
The establishment of stoichiometric laws made it possible to assign a strictly defined mass to the atoms of chemical elements. The masses of atoms are extremely small. Thus, the mass of a hydrogen atom is 1.67∙10 -27 kg, oxygen - 26.60∙10 -27 kg, carbon - 19.93∙10 -27 kg. It is very inconvenient to use such numbers for various calculations. Therefore, since 1961, 1/12 of the mass of the carbon isotope 12 C has been accepted as a unit of atomic mass - atomic mass unit (a.m.u.). Previously, it was called a carbon unit (cu), but now this name is not recommended.
Mass a.m.u. is 1.66. 10 –27 kg or 1.66. 10–24 years
Relative atomic mass of the element (Ar) is called the ratio of the absolute mass of an atom to 1/12 of the absolute mass of an atom of the carbon isotope 12 C. In other words, A r shows how many times the mass of an atom of a given element is heavier than 1/12 of the mass of an atom of 12 C. For example, the A r value of oxygen rounded to a whole number is 16; this means that the mass of one oxygen atom is 16 times greater than 1/12 the mass of a 12 C atom.
The relative atomic masses of elements (Ar) are given in the Periodic Table of Chemical Elements by D.I. Mendeleev.
Relative molecular weight (Mr) a substance is called the mass of its molecule, expressed in amu. It is equal to the sum of the atomic masses of all atoms that make up the molecule of the substance and is calculated using the formula of the substance. For example, the relative molecular weight of sulfuric acid H 2 SO 4 is composed of the atomic masses of two hydrogen atoms (1∙2 = 2), the atomic mass of one sulfur atom (32) and the atomic mass of four oxygen atoms (4∙16 = 64). It is equal to 98.
This means that the mass of a sulfuric acid molecule is 98 times greater than 1/12 the mass of a 12 C atom.
Relative atomic and molecular masses are relative quantities, and therefore dimensionless.
Chemists around the world reflect the composition of simple and complex substances very beautifully and concisely in the form of chemical formulas. Chemical formulas are analogues of words that are written using letters - symbols of chemical elements.
Let us express, using chemical symbols, the composition of the most common substance on Earth - water. A water molecule contains two hydrogen atoms and one oxygen atom. Now let's translate this sentence into a chemical formula using chemical symbols (hydrogen - H and oxygen - O). We write the number of atoms in the formula using indices - numbers located at the bottom right of the chemical symbol (index 1 is not written for oxygen): H 2 0 (read “ash-two-o”).
The formulas of simple substances hydrogen and oxygen, whose molecules consist of two identical atoms, are written as follows: H 2 (read “ash-two”) and 0 2 (read “o-two”) (Fig. 26).
Rice. 26.
Models of molecules and formulas of oxygen, hydrogen and water
To reflect the number of molecules, coefficients are used that are written before chemical formulas: for example, the entry 2CO 2 (read “two-ce-o-two”) means two molecules of carbon dioxide, each of which consists of one carbon atom and two oxygen atoms.
Coefficients are written similarly when indicating the number of free atoms of a chemical element. For example, we need to write down the expression: five iron atoms and seven oxygen atoms. This is done as follows: 5Fe and 7O.
The sizes of molecules, and even more so of atoms, are so small that they cannot be seen even in the best optical microscopes, which provide a magnification of 5-6 thousand times. They cannot be seen even in electron microscopes, which provide a magnification of 40 thousand times. Naturally, the negligible size of molecules and atoms corresponds to their negligible masses. Scientists have calculated, for example, that the mass of a hydrogen atom is 0.000 000 000 000 000 000 000 001 674 g, which can be represented as 1.674 10 -24 g, the mass of an oxygen atom is 0.000 000 000 000 000 000 000 026 667 g , or 2.6667 10 -23 g, the mass of a carbon atom is 1.993 10 -23 g, and the mass of a water molecule is 3.002 10 -23 g.
Let's calculate how many times the mass of an oxygen atom is greater than the mass of a hydrogen atom, the lightest element:
Similarly, the mass of a carbon atom is 12 times greater than the mass of a hydrogen atom:
Rice. 27. The mass of a carbon atom is equal to the mass of 12 hydrogen atoms
The mass of a water molecule is 18 times greater than the mass of a hydrogen atom (Fig. 28). These values show how many times the mass of an atom of a given chemical element is greater than the mass of a hydrogen atom, i.e. they are relative.
Rice. 27. The mass of a water atom is equal to the mass of 18 hydrogen atoms
Currently, physicists and chemists are of the opinion that the relative atomic mass of an element is a value that shows how many times the mass of its atom is greater than 1/12 the mass of a carbon atom. Relative atomic mass is denoted by Ar, where r is the initial letter of the English word relative, which means “relative”. For example, A r (0) = 16, A r (C) = 12, A r (H) = 1.
Each chemical element has its own relative atomic mass value (Fig. 29). The values of the relative atomic masses of chemical elements are indicated in the corresponding cells of D.I. Mendeleev’s table.
Rice. 29.
Each element has its own relative atomic mass value
Similarly, the relative molecular weight of a substance is denoted by M r, for example M r (H 2 0) = 18.
The relative atomic mass of an element A r and the relative molecular mass of a substance M r are quantities that do not have units of measurement.
To find out the relative molecular mass of a substance, it is not necessary to divide the mass of its molecule by the mass of the hydrogen atom. You just need to add up the relative atomic masses of the elements that form the substance, taking into account the number of atoms, for example:
A chemical formula contains important information about a substance. For example, the formula C0 2 shows the following information:
Let's calculate the mass fractions of the elements carbon and oxygen in carbon dioxide CO 2 .
Key words and phrases
- Chemical formula.
- Indices and coefficients.
- Relative atomic mass (A r).
- Relative molecular weight (Mr).
- Mass fraction of an element in a substance.
Working with a computer
- Refer to the electronic application. Study the lesson material and complete the assigned tasks.
- Find email addresses on the Internet that can serve as additional sources that reveal the content of keywords and phrases in the paragraph. Offer your help to the teacher in preparing a new lesson - make a report on the key words and phrases of the next paragraph.
Questions and tasks
- What do the entries mean: 3H; 2H 2 O; 5O2?
- Write down the formula for sucrose if you know that its molecule contains twelve carbon atoms, twenty-two hydrogen atoms and eleven oxygen atoms.
- Using Figure 2, write down the formulas of the substances and calculate their relative molecular weights.
- Which form of existence of the chemical element oxygen corresponds to each of the following entries: 3O; 5O2; 4CO2?
- Why do the relative atomic mass of an element and the relative molecular mass of a substance not have units of measurement?
- Which of the substances whose formulas are SO 2 and SO 3 has a greater mass fraction of sulfur? Confirm your answer with calculations.
- Calculate the mass fractions of elements in nitric acid HNO 3.
- Give a complete description of glucose C 6 H 12 0 6, using the example of describing carbon dioxide C0 2.
The most important method for determining the molecular masses of substances in a gaseous state is based on Avogadro's law. But before talking about this method, it should be said in what units molecular and atomic masses are expressed.
When calculating atomic masses, initially the mass of a hydrogen atom as the lightest element was taken as a unit of mass, and the masses of atoms of other elements were calculated in relation to it. But since the atomic masses of most elements are determined based on the composition of their oxygen compounds, the calculations were actually made in relation to the atomic mass of oxygen, which was considered equal to 16; the ratio between the atomic masses of oxygen and hydrogen was assumed to be equal. Subsequently, more accurate measurements showed that this ratio is equal to or. A change in the atomic mass of oxygen would entail a change in the atomic masses of most elements. Therefore, it was decided to leave the atomic mass of oxygen at 16, taking the atomic mass of hydrogen equal to 1.0079.
Thus, the unit of atomic mass was taken to be part of the mass of the oxygen atom, which was called the oxygen unit. Later it was found that natural oxygen is a mixture of isotopes (see § 35), so the oxygen mass unit characterizes the average mass of the atoms of natural oxygen isotopes . For atomic physics, such a unit turned out to be unacceptable, and in this branch of science, part of the mass of an oxygen atom was accepted as a unit of atomic mass. As a result, two scales of atomic masses took shape - chemical and physical. The presence of two atomic mass scales created great inconvenience.
In 1961, a unified scale of relative atomic masses was adopted, which is based on part of the mass of an atom of a carbon isotope, called the atomic mass unit. In accordance with this, at present the relative atomic mass (abbreviated as atomic mass) of an element is the ratio of the mass of its atom to the part of the mass of the atom. On the modern scale, the relative atomic masses of oxygen and hydrogen are 15.9994 and 1.00794, respectively.
Similarly, the relative molecular weight (abbreviated as molecular weight) of a simple or complex substance is the ratio of the mass of its molecule to part of the mass. Since the mass of any molecule is equal to the sum of the masses of its constituent atoms, the relative molecular mass is equal to the sum of the corresponding relative atomic masses.
For example, the molecular weight of water, the molecule of which contains two hydrogen atoms and one oxygen atom, is equal to: . (Until recently, the terms “atomic weight” and “molecular weight” were used instead of the terms “atomic weight” and “molecular weight.”)
Along with units of mass and volume, chemistry also uses a unit of quantity of a substance called a mole (abbreviated as “mole”).
Mole - an amount of substance containing as many molecules, atoms, ions, electrons or other structural units as there are atoms in an isotope of carbon.
When using the concept of “mole”, it is necessary in each specific case to indicate exactly which structural units are meant. For example, one should distinguish between moles of H atoms, moles of molecules, and moles of ions.
Currently, the number of structural units contained in one mole of a substance (Avogadro's constant) has been determined with great accuracy. In practical calculations it is taken equal to .
The ratio of the mass m of a substance to its quantity is called the molar mass of the substance
Molar mass is usually expressed in g/mol. Since one mole of any substance contains the same number of structural units, the molar mass of the substance (g/mol) is proportional to the mass of the corresponding structural unit, i.e., the relative molecular (or atomic) mass of the substance (Motn)
where K is the proportionality coefficient, the same for all substances.
It is easy to see that K=1. In fact, for the carbon isotope Motn = 12, and the molar mass (by definition of the concept “mole”) is 12 g/mol. Consequently, the numerical values of M (g/mol) and Motn coincide, which means K = 1. It follows that the molar mass of a substance, expressed in grams per mole, has the same numerical value as its relative molecular (atomic) mass. Thus, the molar mass of atomic hydrogen is 1.0079 g/mol, molecular hydrogen is 2.0158 g/mol, and molecular oxygen is 31.9988 g/mol.
According to Avogadro's law, the same number of molecules of any gas occupies the same volume under the same conditions. On the other hand, 1 mole of any substance contains (by definition) the same number of particles. It follows that at a certain temperature and pressure, 1 mole of any substance in the gaseous state occupies the same volume.
It is not difficult to calculate how much volume one mole of gas occupies under normal conditions, i.e. at normal atmospheric pressure or) and temperature. For example, it has been experimentally established that the mass of 1 liter of oxygen under normal conditions is 1.43 grams. Consequently, the volume occupied by one mole of oxygen (32 grams) under the same conditions will be 32:1.43 = 22.4 liters. We get the same number by calculating the volume of one mole of hydrogen, carbon dioxide, etc.
The ratio of the volume occupied by a substance to its quantity is called the molar volume of the substance. As follows from the above, under normal conditions the molar volume of any gas is 22.4 l/mol.
§ 1 What makes up the mass of matter
Any body has mass. Let's take a body such as, for example, 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 own mass, which is determined by adding up the mass of all the rice grains, although they are very small and light.
All bodies are made of substances. 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, particles of matter also have mass.
§ 2 Atomic mass unit
If we express the mass of the lightest hydrogen atom in grams, we get a very difficult number for further work
1.66 ∙10-24g.
The mass of the oxygen atom is approximately sixteen times greater and amounts to 2.66∙10-23 g, the mass of the carbon atom is 1.99∙10-23 g. The mass of an atom is denoted by ma.
It is inconvenient to make calculations with such numbers.
To measure atomic (and molecular) masses, the atomic mass unit (amu) is used.
An atomic mass unit is 1/12 the mass of a carbon atom.
In this case, the mass of a hydrogen atom will be equal to 1 amu, the mass of an oxygen atom will be 16 amu, and the mass of a carbon atom will be 12 amu.
For a long time, chemists did not have the slightest idea of how much one atom of any element weighs in the units of mass that are familiar and convenient to us (grams, kilograms, etc.).
Therefore, initially the task of determining atomic masses was changed.
Attempts have been 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 to this problem was also fraught with great difficulties, and above all with the choice of a standard, that is, the chemical element against which the atomic masses of other elements should be compared.
§ 3 Relative atomic mass
Scientists of the 19th century solved this problem on the basis of experimental data on determining the composition of substances. The lightest atom, the hydrogen atom, was taken as a 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.
They agreed to denote this quantity with the letters Ar (the index “r” is from the initial letter of the English word “relative”). Thus, the recording 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.
Relative atomic mass shows how many times the mass of an atom of one chemical element is greater than the mass of the atom that 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 became 1/12 of the mass of a carbon atom (a carbon atom is 12 times heavier than a hydrogen atom).
The relative atomic mass of an element (Ar) is the ratio of the mass of an atom of a chemical element to 1/12 the mass of a carbon atom.
The values of the atomic masses of chemical elements are given in the Periodic Table of Chemical Elements by D.I. Mendeleev. Take a look at the periodic table and look at any of its cells, for example, number 8.
Under the chemical sign and name in the bottom line, the atomic mass of the chemical element is indicated: the relative atomic mass of oxygen is 15.9994. Please note: the relative atomic masses of almost all chemical elements have fractional values. The reason for this is the existence of isotopes. Let me remind you that isotopes are atoms of the same chemical element that differ slightly in mass.
At school, calculations usually use relative atomic masses, rounded to whole numbers. But in several cases, fractional values are used, for example: the relative atomic mass of chlorine is 35.5.
§ 4 Relative molecular weight
The mass of a molecule is made up of the masses of atoms.
The relative molecular mass of a substance is a number indicating how many times the mass of a molecule of this substance is greater than 1/12 the mass of a carbon atom.
Relative molecular weight is designated - Mr
The relative molecular weight of substances is calculated using chemical formulas expressing the composition of substances. To find the relative molecular mass, it is necessary to sum up the values of the relative atomic masses of the elements that make up the molecule of the 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 with the formula H2O is equal to the sum of two relative values
atomic mass of hydrogen and one value of the relative atomic mass of oxygen:
The relative molecular weight of sulfuric acid having the formula H2SO4 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.
List of used literature:
- NOT. Kuznetsova. Chemistry. 8th grade. Textbook for general education institutions. – M. Ventana-Graf, 2012.
Images used:
The international unit of atomic mass is equal to 1/12 of the mass of the 12C isotope, the main isotope of natural carbon.
1 amu = 1/12 m (12C) = 1.66057 10-24 g
Relative atomic mass (Ar) is 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 12C atom.
The average absolute mass of an atom (m) is equal to the relative atomic mass times the amu.
(Mg) = 24.312 1.66057 10-24 = 4.037 10-23 g
Relative molecular mass (Mr) is a dimensionless quantity that shows how many times the mass of a molecule of a given substance is greater than 1/12 the mass of a 12C carbon atom.
Mg = mg / (1/12 ma(12C))
mr is the mass of a molecule of a given substance;
ma(12C) is the mass of the 12C carbon atom.
Mg = Σ 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 indices.
Mg(B2O3) = 2 Ar(B) + 3 Ar(O) = 2 11 + 3 16 = 70
Mg(KAl(SO4)2) = 1 Ar(K) + 1 Ar(Al) + 1 2 Ar(S) + 2 4 Ar(O) =
1 39 + 1 27 + 1 2 32 + 2 4 16 = 258
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 - the mole.
Amount of substance, mol. Means a certain number of structural elements (molecules, atoms, ions). Denoted by ν, 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 diQuaregna number (NA). The number of particles in 1 mole of any substance is the same and equals 6.02 1023. (Avogadro’s constant has the dimension mol-1).
How many molecules are there in 6.4 g of sulfur? The molecular weight of sulfur is 32 g/mol. We determine the amount of g/mol of substance in 6.4 g of sulfur:
ν(s) = m(s) / M(s) = 6.4 g / 32 g/mol = 0.2 mol
Let's determine the number of structural units (molecules) using Avogadro's constant NA
N(s) = ν(s) NA = 0.2 6.02 1023 = 1.2 1023
Molar mass shows the mass of 1 mole of a substance (denoted M).
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 = NA m(1 molecule) = NA Mg 1 a.m.u. = (NA 1 amu) Mg = Mg
This means that if the mass of a certain molecule is, for example, 80 amu. (SO3), then the mass of one mole of molecules is equal to 80 g. Avogadro’s constant is a proportionality coefficient that ensures the transition from molecular relationships to molar ones. All statements regarding molecules remain valid for moles (with replacement, if necessary, of amu by g). For example, the reaction equation: 2Na + Cl2 → 2NaCl, means that two sodium atoms react with one chlorine molecule or that the same thing, two moles of sodium react with one mole of chlorine.
Stoichiometry. Law of conservation of mass of substances. The law of constancy of the composition of substances of molecular structure. Avogadro's law and consequences from it.
Stoichiometry(from Old Greekστοιχειον “element” + μετρειν “measure”) - section chemistry about the ratios of reagents in chemical reactions.
Allows you to theoretically calculate the required volumes reagents.
Law of Constancy of Composition was discovered by the French scientist Louis Jeanne Prousteau in 1799 and is formulated:
Any pure substance has a constant qualitative and quantitative composition, regardless of its location in nature and the method of production in industry.
For example: H 2 O a) qualitative composition - elements H and O
b) quantitative composition – two hydrogen atoms H, one oxygen atom O.
Water can be obtained:
1. 2H 2 + O 2 = 2H 2 O - reaction of the compound.
2. Cu(OH) 2 t°C H 2 O + CuO – decomposition reaction.
3. HCl + NaOH = H 2 O + NaCl – neutralization reaction.
The meaning of the law of constancy of composition:
· Based on the law, the concepts of “chemical compound” and “mixture of substances” were differentiated
· Based on the law, various practical calculations can be made.
Law of conservation of mass of matter was discovered by M.V. Lomonosov in 1748 and is formulated.
