Electronic structure of an atom, valence, oxidation state. Valency and oxidation state

Electronegativity (EO) is the ability of atoms to attract electrons when bonding with other atoms .

Electronegativity depends on the distance between the nucleus and the valence electrons, and how close the valence shell is to complete. The smaller the radius of an atom and the more valence electrons, the higher its EO.

Fluorine is the most electronegative element. Firstly, it has 7 electrons in its valence shell (only 1 electron is missing from the octet) and, secondly, this valence shell (...2s 2 2p 5) is located close to the nucleus.

The atoms of alkali and alkaline earth metals are the least electronegative. They have large radii and their outer electron shells are far from complete. It is much easier for them to give up their valence electrons to another atom (then the outer shell will become complete) than to “gain” electrons.

Electronegativity can be expressed quantitatively and the elements can be ranked in increasing order. The electronegativity scale proposed by the American chemist L. Pauling is most often used.

The difference in electronegativity of elements in a compound ( ΔX) will allow you to judge the type of chemical bond. If the value ΔX= 0 – connection covalent nonpolar.

When the electronegativity difference is up to 2.0, the bond is called covalent polar, For example: H-F connection in a hydrogen fluoride molecule HF: Δ X = (3.98 - 2.20) = 1.78

Bonds with an electronegativity difference greater than 2.0 are considered ionic. For example: Na-Cl bond in NaCl compound: Δ X = (3.16 - 0.93) = 2.23.

Oxidation state

Oxidation state (CO) - This conventional charge atom in a molecule, calculated under the assumption that the molecule consists of ions and is generally electrically neutral.

When an ionic bond is formed, an electron passes from a less electronegative atom to a more electronegative one, the atoms lose their electrical neutrality and turn into ions. integer charges arise. When a covalent polar bond is formed, the electron is not transferred completely, but partially, so partial charges arise (HCl in the figure below). Let's imagine that the electron has completely transferred from the hydrogen atom to chlorine, and a whole positive charge of +1 has appeared on hydrogen, and -1 on chlorine. Such conventional charges are called the oxidation state.

This figure shows the oxidation states characteristic of the first 20 elements.
Note. The highest CO is usually equal to the group number in the periodic table. Metals of the main subgroups have one characteristic CO, while non-metals, as a rule, have a scatter of CO. Therefore, nonmetals form a large number of compounds and have more “diverse” properties compared to metals.

Examples of determining the oxidation state

Let us determine the oxidation states of chlorine in the compounds:

The rules that we have considered do not always allow us to calculate the CO of all elements, such as in a given aminopropane molecule.

Here it is convenient to use the following technique:

1) We depict structural formula molecules, a dash is a bond, a pair of electrons.

2) We turn the dash into an arrow directed towards the more EO atom. This arrow symbolizes the transition of an electron to an atom. If two identical atoms are connected, we leave the line as it is - there is no transfer of electrons.

3) We count how many electrons “came” and “left”.

For example, let's calculate the charge of the first carbon atom. Three arrows are directed towards the atom, which means 3 electrons have arrived, charge -3.

Second carbon atom: hydrogen gave it an electron, and nitrogen took one electron. The charge has not changed, it is zero. Etc.

Valence

Valence(from Latin valēns “having strength”) - the ability of atoms to form a certain number of chemical bonds with atoms of other elements.

Basically, valence means the ability of atoms to form a certain number of covalent bonds. If an atom has n unpaired electrons and m lone electron pairs, then this atom can form n+m covalent bonds with other atoms, i.e. its valency will be equal n+m. When estimating the maximum valency, one should proceed from the electronic configuration of the “excited” state. For example, the maximum valency of a beryllium, boron and nitrogen atom is 4 (for example, in Be(OH) 4 2-, BF 4 - and NH 4 +), phosphorus - 5 (PCl 5), sulfur - 6 (H 2 SO 4) , chlorine - 7 (Cl 2 O 7).

In some cases, the valency may numerically coincide with the oxidation state, but in no way are they identical to each other. For example, in N2 and CO molecules a triple bond is realized (that is, the valence of each atom is 3), but the oxidation state of nitrogen is 0, carbon +2, oxygen -2.

In nitric acid, the oxidation state of nitrogen is +5, while nitrogen cannot have a valence higher than 4, because it has only 4 orbitals at the outer level (and the bond can be considered as overlapping orbitals). And in general, any element of the second period for the same reason cannot have a valence greater than 4.

A few more “tricky” questions in which mistakes are often made.

Atoms of various chemical elements can attach different number other atoms, i.e., exhibit different valencies.

Valence characterizes the ability of atoms to combine with other atoms. Now, having studied the structure of the atom and the types of chemical bonds, we can consider this concept in more detail.

Valency is the number of single chemical bonds that an atom forms with other atoms in a molecule. The number of chemical bonds refers to the number of shared electron pairs. Since shared pairs of electrons are formed only in the case of a covalent bond, the valence of atoms can only be determined in covalent compounds.

In the structural formula of a molecule, chemical bonds are represented by dashes. The number of lines extending from the symbol of a given element is its valence. Valence always has a positive integer value from I to VIII.

As you remember, the highest valency of a chemical element in an oxide is usually equal to the number of the group in which it is found. To determine the valency of a nonmetal in a hydrogen compound, you need to subtract the group number from 8.

In the simplest cases, valence is equal to the number of unpaired electrons in the atom, so, for example, oxygen (contains two unpaired electrons) has valence II, and hydrogen (contains one unpaired electron) has valence I.

Ionic and metallic crystals do not have common pairs of electrons, so for these substances the concept of valence as the number of chemical bonds does not make sense. For all classes of compounds, regardless of the type of chemical bonds, a more universal concept is applicable, which is called the oxidation state.

Oxidation state

This is the conventional charge on an atom in a molecule or crystal. It is calculated by assuming that all covalent polar bonds are ionic in nature.

Unlike valency, oxidation number can be positive, negative, or zero. In the simplest ionic compounds, the oxidation states coincide with the charges of the ions.

For example, in potassium chloride KCl (K + Cl - ) potassium has an oxidation state of +1, and chlorine -1; in calcium oxide CaO (Ca +2 O -2), calcium exhibits an oxidation state of +2, and oxygen -2. This rule applies to all basic oxides: in them, the oxidation state of the metal is equal to the charge of the metal ion (sodium +1, barium +2, aluminum +3), and the oxidation state of oxygen is -2. The oxidation state is indicated by an Arabic numeral, which is placed above the symbol of the element, similar to valency:

Cu +2 Cl 2 -1 ; Fe +2 S -2

The oxidation state of an element in a simple substance is taken equal to zero:

Na 0 , O 2 0 , S 8 0 , Cu 0

Let's consider how oxidation states in covalent compounds are determined.

Hydrogen chloride HCl is a substance with a polar covalent bond. The common electron pair in the HCl molecule is shifted to the chlorine atom, which has a higher electronegativity. We mentally transform the H-Cl bond into an ionic one (this actually happens in an aqueous solution), completely shifting the electron pair to the chlorine atom. It will acquire a charge of -1, and hydrogen +1. Therefore, chlorine in this substance has an oxidation state of -1, and hydrogen +1:

Real charges and oxidation states of atoms in a hydrogen chloride molecule

Oxidation number and valence are related concepts. In many covalent compounds, the absolute value of the oxidation state of the elements is equal to their valence. There are, however, several cases where the valence is different from the oxidation state. This is typical, for example, for simple substances, where the oxidation state of atoms is zero, and the valence is equal to the number of common electron pairs:

O=O.

The valency of oxygen is II, and the oxidation state is 0.

In a hydrogen peroxide molecule

H-O-O-H

oxygen is divalent and hydrogen is monovalent. At the same time, the oxidation states of both elements are equal to 1 in absolute value:

H 2 +1 O 2 -1

The same element in different compounds can have both positive and negative oxidation states, depending on the electronegativity of the atoms associated with it. Consider, for example, two carbon compounds - methane CH 4 and carbon fluoride (IV) CF 4.

Carbon is more electronegative than hydrogen, so in methane the electron density of the C–H bonds is shifted from hydrogen to carbon, and each of the four hydrogen atoms has an oxidation state of +1, and the carbon atom is -4. In contrast, in the CF4 molecule, the electrons of all bonds are shifted from the carbon atom to the fluorine atoms, the oxidation state of which is -1, therefore, carbon is in the +4 oxidation state. Remember that the oxidation number of the most electronegative atom in a compound is always negative.

Models of methane CH 4 and carbon(IV) fluoride CF 4 molecules. The polarity of bonds is indicated by arrows

Any molecule is electrically neutral, so the sum of the oxidation states of all atoms is zero. Using this rule, from the known oxidation state of one element in a compound, you can determine the oxidation state of another without resorting to reasoning about the displacement of electrons.

As an example, let’s take chlorine(I) oxide Cl 2 O. We proceed from the electrical neutrality of the particle. The oxygen atom in oxides has an oxidation state of –2, which means that both chlorine atoms carry a total charge of +2. It follows that each of them has a +1 charge, i.e. chlorine has an oxidation state of +1:

Cl 2 +1 O -2

In order to correctly place the signs of the oxidation state of different atoms, it is enough to compare their electronegativity. An atom with a higher electronegativity will have a negative oxidation state, and an atom with a lower electronegativity will have a positive oxidation state. According to established rules, the symbol of the most electronegative element is written in the last place in the compound formula:

I +1 Cl -1 , O +2 F 2 -1 , P +5 Cl 5 -1

Real charges and oxidation states of atoms in a water molecule

When determining the oxidation states of elements in compounds, the following rules are observed.

The oxidation state of an element in a simple substance is zero.

Fluorine is the most electronegative chemical element, therefore the oxidation state of fluorine in all substances except F2 is -1.

Oxygen is the most electronegative element after fluorine, therefore the oxidation state of oxygen in all compounds except fluorides is negative: in most cases it is -2, and in hydrogen peroxide H 2 O 2 -1.

The oxidation state of hydrogen is +1 in compounds with non-metals, -1 in compounds with metals (hydrides); zero in the simple substance H 2.

The oxidation states of metals in compounds are always positive. The oxidation state of metals of the main subgroups is usually equal to the group number. Metals of secondary subgroups often have several oxidation states.

The maximum possible positive oxidation state of a chemical element is equal to the group number (exception – Cu +2).

The minimum oxidation state of metals is zero, and that of non-metals is group number minus eight.

The sum of the oxidation states of all atoms in a molecule is zero.

Navigation

• Solving combined problems based on quantitative characteristics of a substance
• Problem solving. The law of constancy of the composition of substances. Calculations using the concepts of “molar mass” and “chemical amount” of a substance
• Solving calculation problems based on quantitative characteristics of matter and stoichiometric laws
• Solving calculation problems based on the laws of the gas state of matter
• Electronic configuration of atoms. The structure of the electron shells of atoms of the first three periods

Electronegativity, like other properties of atoms of chemical elements, changes with increasing serial number element periodically:

The graph above shows the periodicity of changes in the electronegativity of elements of the main subgroups depending on the atomic number of the element.

When moving down a subgroup of the periodic table, the electronegativity of chemical elements decreases, and when moving to the right along the period it increases.

Electronegativity reflects the non-metallicity of elements: the higher the electronegativity value, the more non-metallic properties the element has.

Oxidation state

How to calculate the oxidation state of an element in a compound?

1) The degree of oxidation of chemical elements in simple substances always equal to zero.

2) There are elements that exhibit a constant state of oxidation in complex substances:

3) There are chemical elements that exhibit a constant oxidation state in the vast majority of compounds. These elements include:

Element	Oxidation state in almost all compounds	Exceptions
hydrogen H	+1	Hydrides of alkali and alkaline earth metals, for example:
oxygen O	-2	Hydrogen and metal peroxides: Oxygen fluoride -

4) The algebraic sum of the oxidation states of all atoms in a molecule is always zero. The algebraic sum of the oxidation states of all atoms in an ion is equal to the charge of the ion.

5) The highest (maximum) oxidation state is equal to the group number. Exceptions that do not fall under this rule are elements of the secondary subgroup of group I, elements of the secondary subgroup of group VIII, as well as oxygen and fluorine.

Chemical elements whose group number does not coincide with their highest oxidation state (mandatory to remember)

6) The lowest oxidation state of metals is always zero, and the lowest oxidation state of non-metals is calculated by the formula:

lowest oxidation state of non-metal = group number − 8

Based on the rules presented above, you can establish the oxidation state of a chemical element in any substance.

Finding the oxidation states of elements in various compounds

Example 1

Determine the oxidation states of all elements in sulfuric acid.

Solution:

Let's write the formula of sulfuric acid:

The oxidation state of hydrogen in all complex substances is +1 (except metal hydrides).

The oxidation state of oxygen in all complex substances is -2 (except for peroxides and oxygen fluoride OF 2). Let us arrange the known oxidation states:

Let us denote the oxidation state of sulfur as x:

The sulfuric acid molecule, like the molecule of any substance, is generally electrically neutral, because the sum of the oxidation states of all atoms in a molecule is zero. Schematically this can be depicted as follows:

Those. we got the following equation:

Let's solve it:

Thus, the oxidation state of sulfur in sulfuric acid is +6.

Example 2

Determine the oxidation state of all elements in ammonium dichromate.

Solution:

Let's write the formula of ammonium dichromate:

As in the previous case, we can arrange the oxidation states of hydrogen and oxygen:

However, we see that the oxidation states of two chemical elements at once are unknown - nitrogen and chromium. Therefore, we cannot find oxidation states similarly to the previous example (one equation with two variables does not have a single solution).

Let us draw attention to the fact that this substance belongs to the class of salts and, accordingly, has an ionic structure. Then we can rightly say that the composition of ammonium dichromate includes NH 4 + cations (the charge of this cation can be seen in the solubility table). Consequently, since the formula unit of ammonium dichromate contains two positive singly charged NH 4 + cations, the charge of the dichromate ion is equal to -2, since the substance as a whole is electrically neutral. Those. the substance is formed by NH 4 + cations and Cr 2 O 7 2- anions.

We know the oxidation states of hydrogen and oxygen. Knowing that the sum of the oxidation states of the atoms of all elements in an ion is equal to the charge, and denoting the oxidation states of nitrogen and chromium as x And y accordingly, we can write:

Those. we get two independent equations:

Solving which, we find x And y:

Thus, in ammonium dichromate the oxidation states of nitrogen are -3, hydrogen +1, chromium +6, and oxygen -2.

How to determine the oxidation states of elements in organic matter you can read it.

Valence

The valence of atoms is indicated by Roman numerals: I, II, III, etc.

The valence capabilities of an atom depend on the quantity:

1) unpaired electrons

2) lone electron pairs in the orbitals of valence levels

3) empty electron orbitals of the valence level

Valence possibilities of the hydrogen atom

Let us depict the electron graphic formula of the hydrogen atom:

It has been said that three factors can influence the valence possibilities - the presence of unpaired electrons, the presence of lone electron pairs in the outer level, and the presence of vacant (empty) orbitals in the outer level. We see one unpaired electron at the outer (and only) energy level. Based on this, hydrogen can definitely have a valence of I. However, in the first energy level there is only one sublevel - s, those. The hydrogen atom at the outer level has neither lone electron pairs nor empty orbitals.

Thus, the only valence that a hydrogen atom can exhibit is I.

Valence possibilities of the carbon atom

Let's consider the electronic structure of the carbon atom. In the ground state, the electronic configuration of its outer level is as follows:

Those. in the ground state at the outer energy level of the unexcited carbon atom there are 2 unpaired electrons. In this state it can exhibit a valence of II. However, the carbon atom very easily goes into an excited state when energy is imparted to it, and the electronic configuration of the outer layer in this case takes the form:

Despite the fact that a certain amount of energy is spent on the process of excitation of the carbon atom, the expenditure is more than compensated for by the formation of four covalent bonds. For this reason, valency IV is much more characteristic of the carbon atom. So, for example, carbon has valence IV in carbon dioxide molecules, carbonic acid and absolutely all organic substances.

In addition to unpaired electrons and lone electron pairs, the presence of vacant ()valence level orbitals also affects the valence possibilities. The presence of such orbitals at the filled level leads to the fact that the atom can act as an electron pair acceptor, i.e. form additional covalent bonds through a donor-acceptor mechanism. For example, contrary to expectations, in the carbon monoxide CO molecule the bond is not double, but triple, as is clearly shown in the following illustration:

Valence possibilities of the nitrogen atom

Let us write the electronic graphic formula for the external energy level of the nitrogen atom:

As can be seen from the illustration above, the nitrogen atom in its normal state has 3 unpaired electrons, and therefore it is logical to assume that it is capable of exhibiting a valence of III. Indeed, a valence of three is observed in the molecules of ammonia (NH 3), nitrous acid (HNO 2), nitrogen trichloride (NCl 3), etc.

It was said above that the valence of an atom of a chemical element depends not only on the number of unpaired electrons, but also on the presence of lone electron pairs. This is due to the fact that a covalent chemical bond can be formed not only when two atoms provide each other with one electron, but also when one atom with a lone pair of electrons - donor () provides it to another atom with a vacant () orbital valence level (acceptor). Those. For the nitrogen atom, valence IV is also possible due to an additional covalent bond formed by the donor-acceptor mechanism. For example, four covalent bonds, one of which is formed by a donor-acceptor mechanism, are observed during the formation of an ammonium cation:

Despite the fact that one of the covalent bonds is formed according to the donor-acceptor mechanism, all N-H connections in the ammonium cation are absolutely identical and do not differ from each other in any way.

The nitrogen atom is not capable of exhibiting a valency equal to V. This is due to the fact that it is impossible for a nitrogen atom to transition to an excited state, in which two electrons are paired with the transition of one of them to a free orbital that is closest in energy level. The nitrogen atom has no d-sublevel, and the transition to the 3s orbital is energetically so expensive that the energy costs are not covered by the formation of new bonds. Many may wonder, what is the valency of nitrogen, for example, in molecules of nitric acid HNO 3 or nitric oxide N 2 O 5? Oddly enough, the valency there is also IV, as can be seen from the following structural formulas:

The dotted line in the illustration shows the so-called delocalized π -connection. For this reason, terminal NO bonds can be called “one and a half bonds.” Similar one-and-a-half bonds are also present in the molecule of ozone O 3, benzene C 6 H 6, etc.

Valence possibilities of phosphorus

Let us depict the electronic graphic formula of the external energy level of the phosphorus atom:

As we see, the structure of the outer layer of the phosphorus atom in the ground state and the nitrogen atom is the same, and therefore it is logical to expect for the phosphorus atom, as well as for the nitrogen atom, possible valences equal to I, II, III and IV, as observed in practice.

However, unlike nitrogen, the phosphorus atom also has d-sublevel with 5 vacant orbitals.

In this regard, it is capable of transitioning to an excited state, steaming electrons 3 s-orbitals:

Thus, the valence V for the phosphorus atom, which is inaccessible to nitrogen, is possible. For example, the phosphorus atom has a valency of five in molecules of compounds such as phosphoric acid, phosphorus (V) halides, phosphorus (V) oxide, etc.

Valence possibilities of the oxygen atom

The electron graphic formula for the external energy level of an oxygen atom has the form:

We see two unpaired electrons at the 2nd level, and therefore valence II is possible for oxygen. It should be noted that this valence of the oxygen atom is observed in almost all compounds. Above, when considering the valency capabilities of the carbon atom, we discussed the formation of the carbon monoxide molecule. The bond in the CO molecule is triple, therefore, the oxygen there is trivalent (oxygen is an electron pair donor).

Due to the fact that the oxygen atom does not have an external d-sublevel, electron pairing s And p- orbitals is impossible, which is why the valence capabilities of the oxygen atom are limited compared to other elements of its subgroup, for example, sulfur.

Valence possibilities of the sulfur atom

External energy level of a sulfur atom in an unexcited state:

The sulfur atom, like the oxygen atom, normally has two unpaired electrons, so we can conclude that a valence of two is possible for sulfur. Indeed, sulfur has valency II, for example, in the hydrogen sulfide molecule H 2 S.

As we see, the sulfur atom appears at the external level d-sublevel with vacant orbitals. For this reason, the sulfur atom is able to expand its valence capabilities, unlike oxygen, due to the transition to excited states. Thus, when pairing a lone electron pair 3 p-sublevel, the sulfur atom acquires the electronic configuration of the outer level of the following form:

In this state, the sulfur atom has 4 unpaired electrons, which tells us that sulfur atoms can exhibit a valence of IV. Indeed, sulfur has valency IV in molecules SO 2, SF 4, SOCl 2, etc.

When pairing the second lone electron pair located at 3 s-sublevel, the external energy level acquires the configuration:

In this state, the manifestation of valency VI becomes possible. Examples of compounds with VI-valent sulfur are SO 3, H 2 SO 4, SO 2 Cl 2, etc.

Similarly, we can consider the valence possibilities of other chemical elements.

Video tutorial 2: Oxidation state of chemical elements

Video tutorial 3: Valence. Determination of valency

Lecture: Electronegativity. Oxidation state and valency of chemical elements

Electronegativity

Electronegativity is the ability of atoms to attract electrons from other atoms to join them.

It is easy to judge the electronegativity of a particular chemical element using the table. Remember, in one of our lessons it was said that it increases when moving from left to right through periods in the periodic table and when moving from bottom to top through groups.

For example, the task was given to determine which element from the proposed series is the most electronegative: C (carbon), N (nitrogen), O (oxygen), S (sulfur)? We look at the table and find that this is O, because he is to the right and higher than the others.

What factors influence electronegativity? This:

• The radius of an atom, the smaller it is, the higher the electronegativity.
• The valence shell is filled with electrons; the more electrons there are, the higher the electronegativity.

Of all the chemical elements, fluorine is the most electronegative because it has a small atomic radius and 7 electrons in its valence shell.

Elements with low electronegativity include alkali and alkaline earth metals. They have large radii and very few electrons in the outer shell.

The electronegativity values ​​of an atom cannot be constant, because it depends on many factors, including those listed above, as well as the degree of oxidation, which can be different for the same element. Therefore, it is customary to talk about the relativity of electronegativity values. You can use the following scales:

You will need electronegativity values ​​when writing formulas for binary compounds consisting of two elements. For example, the formula of copper oxide Cu 2 O - the first element should be written down the one whose electronegativity is lower.

At the moment of formation of a chemical bond, if the electronegativity difference between the elements is greater than 2.0, a covalent polar bond is formed; if less, an ionic bond is formed.

Oxidation state

Oxidation state (CO)- this is the conditional or real charge of an atom in a compound: conditional - if the bond is polar covalent, real - if the bond is ionic.

An atom acquires a positive charge when it gives up electrons, and a negative charge when it accepts electrons.

Oxidation states are written above the symbols with a sign «+»/«-» . There are also intermediate COs. The maximum CO of an element is positive and equal to group number, and the minimum negative for metals is zero, for non-metals = (Group No. – 8). Elements with maximum CO only accept electrons, and elements with minimum CO only give up electrons. Elements that have intermediate COs can both give and receive electrons.

Let's look at some rules that should be followed to determine CO:

The CO of all simple substances is zero.

The sum of all CO atoms in a molecule is also equal to zero, since any molecule is electrically neutral.

In compounds with a covalent nonpolar bond, CO is equal to zero (O 2 0), and with an ionic bond it is equal to the charges of the ions (Na + Cl - sodium CO +1, chlorine -1). CO elements of compounds with a covalent polar bond are considered as with an ionic bond (H:Cl = H + Cl -, which means H +1 Cl -1).

Elements in a compound that have the greatest electronegativity have negative oxidation states, while those with the least electronegativity have positive oxidation states. Based on this, we can conclude that metals have only a “+” oxidation state.

Constant oxidation states:

Alkali metals +1.

All metals of the second group +2. Exception: Hg +1, +2.

Aluminum +3.

• Hydrogen +1. Exception: hydrides of active metals NaH, CaH 2, etc., where the oxidation state of hydrogen is –1.

  Oxygen –2. Exception: F 2 -1 O +2 and peroxides that contain the –O–O– group, in which the oxidation state of oxygen is –1.

When is it formed ionic bond, a certain transition of an electron occurs, from a less electronegative atom to an atom of greater electronegativity. Also, in this process, atoms always lose electrical neutrality and subsequently turn into ions. Integer charges are also formed. When a polar covalent bond is formed, the electron is transferred only partially, so partial charges arise.

Valence

Valenceis the ability of atoms to form n - the number of chemical bonds with atoms of other elements.

Valence is also the ability of an atom to hold other atoms near itself. As you know from your school chemistry course, different atoms are bonded to each other by electrons from the outer energy level. An unpaired electron seeks a pair from another atom. These outer level electrons are called valence electrons. This means that valence can also be defined as the number of electron pairs connecting atoms to each other. Look at the structural formula of water: H – O – H. Each dash is an electron pair, which means it shows the valency, i.e. oxygen here has two lines, which means it is divalent, hydrogen molecules come from one line each, which means hydrogen is monovalent. When writing, valence is indicated by Roman numerals: O (II), H (I). Can also be indicated above the element.

Valence can be constant or variable. For example, in metal alkalis it is constant and equals I. But chlorine in various compounds exhibits valencies I, III, V, VII.

How to determine the valency of an element?

Let's look again at the Periodic Table. Metals of the main subgroups have a constant valency, so metals of the first group have valency I, the second - II. And metals of side subgroups have variable valency. It is also variable for non-metals. The highest valency of an atom is equal to group number, the lowest is equal to = group number - 8. A familiar formulation. Doesn't this mean that the valency coincides with the oxidation state? Remember, valence may coincide with the oxidation state, but these indicators are not identical to each other. Valency cannot have a =/- sign, and also cannot be zero.

The second method is to determine valency using a chemical formula, if the constant valency of one of the elements is known. For example, take the formula of copper oxide: CuO. Oxygen valency II. We see that for one oxygen atom in this formula there is one copper atom, which means that the valence of copper is equal to II. Now let's take a more complicated formula: Fe 2 O 3. The valency of the oxygen atom is II. There are three such atoms here, multiply 2*3 =6. We found that there are 6 valences per two iron atoms. Let's find out the valence of one iron atom: 6:2=3. This means that the valence of iron is III.

In addition, when it is necessary to estimate the "maximum valence", one should always start from the electronic configuration that is present in the "excited" state.

Chapter 3. CHEMICAL BOND

The ability of an atom of a chemical element to attach or replace a certain number of atoms of another element to form a chemical bond is called the valency of the element.

Valence is expressed as a positive integer ranging from I to VIII. Valence equal to 0 or greater VIII no. Constant valence is exhibited by hydrogen (I), oxygen (II), alkali metals - elements of the first group of the main subgroup (I), alkaline earth elements - elements of the second group of the main subgroup (II). Atoms of other chemical elements exhibit variable valence. Thus, transition metals - elements of all secondary subgroups - exhibit from I to III. For example, iron in compounds can be di- or trivalent, copper - mono- and divalent. The atoms of other elements can exhibit a valence in compounds equal to the group number and intermediate valences. For example, the highest valency of sulfur is IV, the lowest is II, and the intermediate ones are I, III and IV.

Valency is equal to the number of chemical bonds by which an atom of a chemical element is connected to atoms of other elements in chemical compound. A chemical bond is indicated by a dash (–). Formulas that show the order of connection of atoms in a molecule and the valence of each element are called graphical.

Oxidation state is the conditional charge of an atom in a molecule, calculated under the assumption that all bonds are ionic in nature. This means that a more electronegative atom, by displacing one electron pair completely towards itself, acquires a charge of 1–. Non-polar covalent bond between identical atoms does not contribute to the oxidation state.

To calculate the oxidation state of an element in a compound, one should proceed from the following provisions:

1) the oxidation states of elements in simple substances are assumed to be zero (Na 0; O 2 0);

2) the algebraic sum of the oxidation states of all atoms that make up the molecule is equal to zero, and in a complex ion this sum is equal to the charge of the ion;

3) atoms have a constant oxidation state: alkali metals(+1), alkaline earth metals, zinc, cadmium (+2);

4) the oxidation state of hydrogen in compounds is +1, except for metal hydrides (NaH, etc.), where the oxidation state of hydrogen is –1;

5) the oxidation state of oxygen in compounds is –2, except for peroxides (–1) and oxygen fluoride OF2 (+2).

The maximum positive oxidation state of an element usually coincides with its group number in periodic table. The maximum negative oxidation state of an element is equal to the maximum positive oxidation state minus eight.

The exceptions are fluorine, oxygen, iron: their highest oxidation state is expressed by a number whose value is lower than the number of the group to which they belong. Elements of the copper subgroup, on the contrary, have a highest oxidation state greater than one, although they belong to group I.

Atoms of chemical elements (except noble gases) can interact with each other or with atoms of other elements forming b.m. complex particles - molecules, molecular ions and free radicals. The chemical bond is due electrostatic forces between atoms , those. forces of interaction between electrons and atomic nuclei. The main role in the formation of chemical bonds between atoms is played by valence electrons, i.e. electrons located in the outer shell.