Electronegativity, oxidation state and valence of chemical elements
Electronegativity
The concept is widely used in chemistry electronegativity (EO).
The property of atoms of a given element to attract electrons from atoms of other elements in compounds is called electronegativity.
The electronegativity of lithium is conventionally taken as unity, the EO of other elements is calculated accordingly. There is a scale of values for EO elements.
The numerical values of EO elements have approximate values: it is a dimensionless quantity. The higher the EO of an element, the more clearly its non-metallic properties appear. According to EO, the elements can be written as follows:
$F > O > Cl > Br > S > P > C > H > Si > Al > Mg > Ca > Na > K > Cs$. Fluorine has the greatest EO value.
Comparing the EO values of elements from francium $(0.86)$ to fluorine $(4.1)$, it is easy to notice that EO obeys the Periodic Law.
In the Periodic Table of Elements, EO in a period increases with the element number (from left to right), and in the main subgroups it decreases (from top to bottom).
In periods, as the charges of the atomic nuclei increase, the number of electrons on the outer layer increases, the radius of the atoms decreases, therefore the ease of electron loss decreases, the EO increases, and therefore the non-metallic properties increase.
Oxidation state
Complex substances consisting of two chemical elements are called binary(from lat. bi - two), or two-element.
Let us recall the typical binary compounds that were given as an example to consider the mechanisms of formation of ionic and covalent polar bonds: $NaCl$ - sodium chloride and $HCl$ - hydrogen chloride. In the first case, the bond is ionic: the sodium atom transferred its outer electron to the chlorine atom and turned into an ion with a charge of $+1$, and the chlorine atom accepted an electron and turned into an ion with a charge of $-1$. Schematically, the process of converting atoms into ions can be depicted as follows:
$(Na)↖(0)+(Cl)↖(0)→(Na)↖(+1)(Cl)↖(-1)$.
In the $HCl$ molecule, the bond is formed due to the pairing of unpaired outer electrons and the formation of a common electron pair of hydrogen and chlorine atoms.
It is more correct to imagine the formation of a covalent bond in a hydrogen chloride molecule as the overlap of the one-electron $s$-cloud of the hydrogen atom with the one-electron $p$-cloud of the chlorine atom:
During a chemical interaction, the shared electron pair is shifted towards the more electronegative chlorine atom: $(H)↖(δ+)→(Cl)↖(δ−)$, i.e. the electron will not completely transfer from the hydrogen atom to the chlorine atom, but partially, thereby determining the partial charge of the atoms $δ$: $H^(+0.18)Cl^(-0.18)$. If we imagine that in the $HCl$ molecule, as well as in the $NaCl$ chloride, the electron has completely transferred from the hydrogen atom to the chlorine atom, then they would receive charges of $+1$ and $-1$: $(H)↖ (+1)(Cl)↖(−1). Such conditional charges are called degree of oxidation. When defining this concept, it is conventionally assumed that in covalent polar compounds the bonding electrons are completely transferred to a more electronegative atom, and therefore the compounds consist only of positively and negatively charged atoms.
The oxidation state is the conditional charge of the atoms of a chemical element in a compound, calculated based on the assumption that all compounds (both ionic and covalently polar) consist only of ions.
The oxidation number can have a negative, positive or zero value, which is usually placed above the element symbol at the top, for example:
$(Na_2)↖(+1)(S)↖(-2), (Mg_3)↖(+2)(N_2)↖(-3), (H_3)↖(-1)(N)↖(-3 ), (Cl_2)↖(0)$.
Those atoms that have accepted electrons from other atoms or to which common electron pairs are displaced have a negative oxidation state value, i.e. atoms of more electronegative elements.
The oxidation state has a positive value for those atoms that donate their electrons to other atoms or from which common electron pairs are drawn, i.e. atoms of less electronegative elements.
Atoms in molecules of simple substances and atoms in a free state have a zero oxidation state.
In compounds, the total oxidation state is always zero. Knowing this and the oxidation state of one of the elements, you can always find the oxidation state of another element using the formula of a binary compound. For example, let's find the oxidation state of chlorine: $Cl_2O_7$. Let us denote the oxidation state of oxygen: $(Cl_2)(O_7)↖(-2)$. Therefore, seven oxygen atoms will have a total negative charge of $(-2)·7=-14$. Then the total charge of two chlorine atoms is $+14$, and one chlorine atom is $(+14):2=+7$.
Similarly, knowing the oxidation states of elements, you can create a formula for a compound, for example, aluminum carbide (a compound of aluminum and carbon). Let's write the signs of aluminum and carbon side by side - $AlC$, with the sign of aluminum first, because it's metal. Using the periodic table of elements, we determine the number of outer electrons: $Al$ has $3$ electrons, $C$ has $4$. The aluminum atom will give up its three outer electrons to carbon and will receive an oxidation state of $+3$, equal to the charge of the ion. The carbon atom, on the contrary, will take the $4$ electrons missing to the “cherished eight” and receive an oxidation state of $-4$. Let's write these values into the formula $((Al)↖(+3)(C)↖(-4))$ and find the least common multiple for them, it is equal to $12$. Then we calculate the indices:
Valence
The concept is very important in describing the chemical structure of organic compounds valence.
Valence characterizes the ability of atoms of chemical elements to form chemical bonds; it determines the number of chemical bonds by which a given atom is connected to other atoms in the molecule.
The valency of an atom of a chemical element is determined, first of all, by the number of unpaired electrons participating in the formation of a chemical bond.
The valence capabilities of atoms are determined:
- the number of unpaired electrons (one-electron orbitals);
- the presence of free orbitals;
- presence of lone pairs of electrons.
In organic chemistry, the concept of “valence” replaces the concept of “oxidation state”, which is usually used in inorganic chemistry. However, this is not the same thing. Valence has no sign and cannot be zero, while the oxidation state is necessarily characterized by a sign and can have a value equal to zero.
Part 1. Task A5.
Checked elements: Electronegativity. Oxidation state and
valence of chemical elements.
Electronegativity-a value characterizing the ability of an atom to polarize covalent bonds. If in a diatomic molecule A - B the electrons forming the bond are attracted to atom B more strongly than to atom A, then atom B is considered more electronegative than A.
The electronegativity of an atom is the ability of an atom in a molecule (compound) to attract electrons that bind it to other atoms.
The concept of electronegativity (EO) was introduced by L. Pauling (USA, 1932). The quantitative characteristic of the electronegativity of an atom is very conditional and cannot be expressed in units of any physical quantities, therefore several scales have been proposed for the quantitative determination of EO. The scale of relative EO has received the greatest recognition and distribution:
Electronegativity values of elements according to Pauling
Electronegativity χ (Greek chi) is the ability of an atom to hold external (valence) electrons. It is determined by the degree of attraction of these electrons to the positively charged nucleus.
This property manifests itself in chemical bonds as a shift of bond electrons towards a more electronegative atom.
The electronegativity of the atoms involved in the formation of a chemical bond is one of the main factors that determines not only the TYPE, but also the PROPERTIES of this bond, and thereby affects the nature of the interaction between atoms during a chemical reaction.
In L. Pauling's scale of relative electronegativities of elements (compiled on the basis of the bond energies of diatomic molecules), metals and organogenic elements are arranged in the following row:
The electronegativity of elements obeys the periodic law: it increases from left to right in periods and from bottom to top in the main subgroups of the Periodic Table of Elements D.I. Mendeleev.
Electronegativity is not an absolute constant of an element. It depends on the effective charge of the atomic nucleus, which can change under the influence of neighboring atoms or groups of atoms, the type of atomic orbitals and the nature of their hybridization.
Oxidation state is the conditional charge of the atoms of a chemical element in a compound, calculated from the assumption that the compounds consist only of ions.
Oxidation states can have a positive, negative or zero value, and the sign is placed before the number: -1, -2, +3, in contrast to the charge of the ion, where the sign is placed after the number.
In molecules, the algebraic sum of the oxidation states of elements, taking into account the number of their atoms, is equal to 0.
The oxidation states of metals in compounds are always positive, the highest oxidation state corresponds to the number of the group of the periodic system where the element is located (excluding some elements: gold Au+3 (group I), Cu+2 (II), from group VIII the oxidation state +8 can only osmium Os and ruthenium Ru.
The degrees of non-metals can be both positive and negative, depending on which atom it is connected to: if with a metal atom it is always negative, if with a non-metal it can be both + and - (you will learn about this when studying a number of electronegativities) . The highest negative oxidation state of non-metals can be found by subtracting from 8 the number of the group in which the element is located, the highest positive is equal to the number of electrons in the outer layer (the number of electrons corresponds to the group number).
The oxidation states of simple substances are 0, regardless of whether it is a metal or a non-metal.
Table showing constant powers for the most commonly used elements:
The degree of oxidation (oxidation number, formal charge) is an auxiliary conventional value for recording the processes of oxidation, reduction and redox reactions, the numerical value of the electrical charge assigned to an atom in a molecule under the assumption that the electron pairs that carry out the bond are completely shifted towards more electronegative ones atoms.
Ideas about the degree of oxidation form the basis for the classification and nomenclature of inorganic compounds.
The degree of oxidation is a purely conventional value that has no physical meaning, but characterizes the formation of a chemical bond of interatomic interaction in a molecule.
Valency of chemical elements -(from Latin valens - having strength) - the ability of atoms of chemical elements to form a certain number of chemical bonds with atoms of other elements. In compounds formed by ionic bonds, the valency of the atoms is determined by the number of electrons added or given up. In compounds with covalent bonds, the valence of atoms is determined by the number of shared electron pairs formed.
Constant valence:
Remember:
The oxidation state is the conditional charge of the atoms of a chemical element in a compound, calculated on the assumption that all bonds are ionic in nature.
1. An element in a simple substance has a zero oxidation state. (Cu, H2)
2. The sum of the oxidation states of all atoms in a molecule of a substance is zero.
3. All metals have a positive oxidation state.
4. Boron and silicon in compounds have positive oxidation states.
5. Hydrogen has an oxidation state (+1) in compounds. Excluding hydrides
(hydrogen compounds with metals of the main subgroup of the first and second groups, oxidation state -1, for example Na + H -)
6. Oxygen has an oxidation state (-2), with the exception of the compound of oxygen with fluorine OF2, the oxidation state of oxygen (+2), the oxidation state of fluorine (-1). And in peroxides H 2 O 2 - the oxidation state of oxygen (-1);
7. Fluorine has an oxidation state (-1).
Electronegativity is the property of HeMe atoms to attract common electron pairs. Electronegativity has the same dependence as that of Nonmetallic properties: it increases along the period (from left to right), and decreases along the group (from above).
The most electronegative element is Fluorine, then Oxygen, Nitrogen...etc....
Algorithm for completing the task in the demo version:
Exercise:
The chlorine atom is located in group 7, so it can have a maximum oxidation state of +7.
The chlorine atom exhibits this degree of oxidation in the substance HClO4.
Let's check this: The two chemical elements hydrogen and oxygen have constant oxidation states and are equal to +1 and -2, respectively. The number of oxidation states for oxygen is (-2)·4=(-8), for hydrogen (+1)·1=(+1). The number of positive oxidation states is equal to the number of negative ones. Therefore (-8)+(+1)=(-7). This means that the chromium atom has 7 positive degrees, so we write down the oxidation states above the elements. The oxidation state of chlorine is +7 in the HClO4 compound.
Answer: Option 4. The oxidation state of chlorine is +7 in the HClO4 compound.
Various formulations of task A5:
3. Oxidation state of chlorine in Ca(ClO 2) 2
1) 0 2) -3 3) +3 4) +5
4.The element has the lowest electronegativity
5. Manganese has the lowest oxidation state in the compound
1)MnSO 4 2)MnO 2 3)K 2 MnO 4 4)Mn 2 O 3
6. Nitrogen exhibits an oxidation state of +3 in each of the two compounds
1)N 2 O 3 NH 3 2)NH 4 Cl N 2 O 3)HNO 2 N 2 H 4 4)NaNO 2 N 2 O 3
7.The valency of the element is
1) the number of σ bonds it forms
2) the number of connections it forms
3) the number of covalent bonds it forms
4) oxidation states with the opposite sign
8. Nitrogen exhibits its maximum oxidation state in the compound
1)NH 4 Cl 2)NO 2 3)NH 4 NO 3 4)NOF
Electronegativity, like other properties of atoms of chemical elements, changes periodically with increasing atomic number of the element:
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 oxidation state of chemical elements in simple substances is always 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.
You can read how to determine the oxidation states of elements in organic substances.
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 electronic 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 valency 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. For example, carbon has valency IV in the molecules of carbon dioxide, 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 molecule CO 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 bonds in the ammonium cation are absolutely identical and do not differ from each other.
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 so energetically 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 valence 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 valence 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 valence 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.
DEFINITION
The ability of an atom to form chemical bonds is called valence. A quantitative measure of valence is considered to be the number of different atoms in a molecule with which a given element forms bonds.
According to the exchange mechanism of the valence bond method, the valence of chemical elements is determined by the number of unpaired electrons contained in an atom. For s- and p-elements, these are electrons of the outer level; for d-elements, these are electrons of the outer and pre-external levels.
The values of the highest and lowest valencies of a chemical element can be determined using the Periodic Table D.I. Mendeleev. The highest valence of an element coincides with the number of the group in which it is located, and the lowest is the difference between the number 8 and the group number. For example, bromine is located in group VIIA, which means its highest valence is VII, and its lowest is I.
Paired electrons (located two at a time in atomic orbitals) upon excitation can be separated in the presence of free cells of the same level (the separation of electrons into any level is impossible). Let's look at the example of elements of groups I and II. For example, the valence of elements of the main subgroup of group I is equal to one, since at the outer level the atoms of these elements have one electron:
3 Li 1s 2 2s 1
The valence of elements of the main subgroup of group II in the ground (unexcited) state is zero, since there are no unpaired electrons at the outer energy level:
4 Be 1s 2 2 s 2
When these atoms are excited, the paired s-electrons are separated into free cells of the p-sublevel of the same level and the valence becomes equal to two (II):
Oxidation state
To characterize the state of elements in compounds, the concept of oxidation state was introduced.
DEFINITION
The number of electrons displaced from an atom of a given element or to an atom of a given element in a compound is called oxidation state.
A positive oxidation state indicates the number of electrons that are displaced from a given atom, and a negative oxidation state indicates the number of electrons that are displaced toward a given atom.
From this definition it follows that in compounds with non-polar bonds the oxidation state of elements is zero. Examples of such compounds are molecules consisting of identical atoms (N 2, H 2, Cl 2).
The oxidation state of metals in the elemental state is zero, since the distribution of electron density in them is uniform.
In simple ionic compounds, the oxidation state of the elements included in them is equal to the electric charge, since during the formation of these compounds there is an almost complete transition of electrons from one atom to another: Na +1 I -1, Mg +2 Cl -1 2, Al +3 F - 1 3 , Zr +4 Br -1 4 .
When determining the oxidation state of elements in compounds with polar covalent bonds, their electronegativity values are compared. Since during the formation of a chemical bond, electrons are displaced to the atoms of more electronegative elements, the latter have a negative oxidation state in compounds.
The concept of oxidation state for most compounds is conditional, since it does not reflect the real charge of the atom. However, this concept is very widely used in chemistry.
Most elements can exhibit varying degrees of oxidation in compounds. When determining their oxidation state, they use the rule according to which the sum of the oxidation states of elements in electrically neutral molecules is equal to zero, and in complex ions - the charge of these ions. As an example, let's calculate the degree of oxidation of nitrogen in compounds of the composition KNO 2 and HNO 3. The oxidation state of hydrogen and alkali metals in compounds is (+), and the oxidation state of oxygen is (-2). Accordingly, the oxidation degree of nitrogen is equal to:
KNO 2 1+ x + 2 × (-2) = 0, x=+3.
HNO 3 1+x+ x + 3 × (-2) = 0, x=+5.
Examples of problem solving
EXAMPLE 1
| Exercise | Valence IV is characteristic of: a) Ca; b) P; c) O; d)Si? |
| Solution | In order to give the correct answer to the question posed, we will consider each of the proposed options separately. a) Calcium is a metal. It is characterized by the only possible valency value, coinciding with the group number in the Periodic Table D.I. Mendeleev, in which it is located, i.e. The valency of calcium is II. The answer is incorrect. b) Phosphorus is a non-metal. Refers to a group of chemical elements with variable valence: the highest is determined by the group number in the Periodic Table D.I. Mendeleev, in which it is located, i.e. is equal to V, and the lowest is the difference between the number 8 and the group number, i.e. equal to III. The answer is incorrect. c) Oxygen is a non-metal. It is characterized by the only possible valency value equal to II. The answer is incorrect. d) Silicon is a non-metal. It is characterized by the only possible valency value, coinciding with the group number in the Periodic Table D.I. Mendeleev, in which it is located, i.e. The valency of silicon is IV. This is the correct answer. |
| Answer | Option (d) |
EXAMPLE 2
| Exercise | What is the valence of iron in the compound that is formed when it reacts with hydrochloric acid: a) I; b) II; c) III; d) VIII? |
| Solution | Let us write the equation for the interaction of iron with hydrochloric acid: Fe + HCl = FeCl 2 + H 2. As a result of the interaction, ferric chloride is formed and hydrogen is released. To determine the valency of iron using the chemical formula, we first count the number of chlorine atoms: We calculate the total number of chlorine valence units: We determine the number of iron atoms: it is equal to 1. Then the valence of iron in its chloride will be equal to: |
| Answer | The valency of iron in the compound formed during its interaction with hydrochloric acid is II. |
Among chemical reactions, including in nature, redox reactions are the most common. These include, for example, photosynthesis, metabolism, biological processes, as well as the combustion of fuel, the production of metals and many other reactions. Redox reactions have long been successfully used by humanity for various purposes, but the electronic theory of redox processes itself appeared quite recently - at the beginning of the 20th century.
In order to move on to the modern theory of oxidation-reduction, it is necessary to introduce several concepts - these are valence, oxidation state and structure of electronic shells of atoms. While studying sections such as , elements and , we have already encountered these concepts. Next, let's look at them in more detail.
Valency and oxidation state
Valence- a complex concept that arose together with the concept of a chemical bond and is defined as the property of atoms to attach or replace a certain number of atoms of another element, i.e. is the ability of atoms to form chemical bonds in compounds. Initially, valency was determined by hydrogen (its valency was taken to be 1) or oxygen (valency was taken to be 2). Later they began to distinguish between positive and negative valence. Quantitatively, positive valency is characterized by the number of electrons donated by an atom, and negative valency is characterized by the number of electrons that must be added to the atom to implement the octet rule (i.e., completion of the external energy level). Later, the concept of valence also began to combine the nature of the chemical bonds that arise between atoms in their connection.
As a rule, the highest valence of elements corresponds to the group number in the periodic table. But, as with all rules, there are exceptions: for example, copper and gold are in the first group of the periodic table and their valency must be equal to the group number, i.e. 1, but in reality the highest valence of copper is 2, and gold is 3.
Oxidation state sometimes called oxidation number, electrochemical valence or oxidation state and is a relative concept. Thus, when calculating the oxidation state, it is assumed that the molecule consists only of ions, although most compounds are not ionic at all. Quantitatively, the degree of oxidation of the atoms of an element in a compound is determined by the number of electrons attached to the atom or displaced from the atom. Thus, in the absence of electron displacement, the oxidation state will be zero, when electrons are displaced towards a given atom, it will be negative, and when electrons are displaced from a given atom, it will be positive.
Defining oxidation state of atoms the following rules must be followed:
- In molecules of simple substances and metals, the oxidation state of atoms is 0.
- Hydrogen in almost all compounds has an oxidation state equal to +1 (and only in hydrides of active metals equal to -1).
- For oxygen atoms in its compounds, the typical oxidation state is -2 (exceptions: OF 2 and metal peroxides, the oxidation state of oxygen is +2 and -1, respectively).
- The atoms of alkali (+1) and alkaline earth (+2) metals, as well as fluorine (-1) also have a constant oxidation state
- In simple ionic compounds, the oxidation state is equal in magnitude and sign to its electric charge.
- For a covalent compound, the more electronegative atom has an oxidation state with a “-” sign, and the less electronegative one has a “+” sign.
- For complex compounds, the oxidation state of the central atom is indicated.
- The sum of the oxidation states of atoms in a molecule is zero.
For example, let's determine the oxidation state of Se in the compound H 2 SeO 3
So, the oxidation state of hydrogen is +1, oxygen -2, and the sum of all oxidation states is 0, let’s create an expression, taking into account the number of atoms in the compound H 2 + Se x O 3 -2:
(+1)2+x+(-2)3=0, whence
those. H 2 + Se +4 O 3 -2
Knowing what the oxidation state of an element in a compound is, it is possible to predict its chemical properties and reactivity towards other compounds, as well as whether this compound is reducing agent or oxidizing agent. These concepts are fully revealed in oxidation-reduction theories:
- Oxidation is the process of loss of electrons by an atom, ion or molecule, which leads to an increase in the oxidation state.
Al 0 -3e - = Al +3 ;
2O -2 -4e - = O 2 ;
2Cl - -2e - = Cl 2
- Recovery - This is the process by which an atom, ion or molecule gains electrons, resulting in a decrease in oxidation state.
Ca +2 +2e - = Ca 0 ;
2H + +2e - =H 2
- Oxidizing agents– compounds that accept electrons during a chemical reaction, and reducing agents– electron donating compounds. Reducing agents are oxidized during a reaction, and oxidizing agents are reduced.
- The essence of redox reactions– movement of electrons (or displacement of electron pairs) from one substance to another, accompanied by a change in the oxidation states of atoms or ions. In such reactions, one element cannot be oxidized without reducing the other, because The transfer of electrons always causes both oxidation and reduction. Thus, the total number of electrons taken away from one element during oxidation is the same as the number of electrons gained by another element during reduction.
So, if the elements in compounds are in their highest oxidation states, then they will exhibit only oxidizing properties, due to the fact that they can no longer give up electrons. On the contrary, if the elements in the compounds are in their lowest oxidation states, then they exhibit only reducing properties, because they can no longer add electrons. Atoms of elements in an intermediate oxidation state, depending on the reaction conditions, can be both oxidizing agents and reducing agents. Let's give an example: sulfur in its highest oxidation state +6 in the compound H 2 SO 4 can exhibit only oxidizing properties, in the compound H 2 S - sulfur is in its lowest oxidation state -2 and will exhibit only reducing properties, and in the compound H 2 SO 3 being in the intermediate oxidation state +4, sulfur can be both an oxidizing agent and a reducing agent.
Based on the oxidation states of elements, the likelihood of a reaction between substances can be predicted. It is clear that if both elements in their compounds are in higher or lower oxidation states, then a reaction between them is impossible. A reaction is possible if one of the compounds can exhibit oxidizing properties, and the other – reducing properties. For example, in HI and H 2 S, both iodine and sulfur are in their lowest oxidation states (-1 and -2) and can only be reducing agents, therefore, they will not react with each other. But they will interact well with H 2 SO 4, which is characterized by reducing properties, because sulfur here is in its highest state of oxidation.
The most important reducing and oxidizing agents are presented in the following table.
| Restorers | |
| Neutral atoms | General scheme M—ne →Mn+ All metals, as well as hydrogen and carbon. The most powerful reducing agents are alkali and alkaline earth metals, as well as lanthanides and actinides. Weak reducing agents are noble metals - Au, Ag, Pt, Ir, Os, Pd, Ru, Rh. In the main subgroups of the periodic table, the reducing ability of neutral atoms increases with increasing atomic number. |
| negatively charged nonmetal ions | General scheme E +ne - → En- Negatively charged ions are strong reducing agents due to the fact that they can donate both excess electrons and their outer electrons. The reducing power, with the same charge, increases with increasing atomic radius. For example, I is a stronger reducing agent than Br - and Cl -. Reducing agents can also be S 2-, Se 2-, Te 2- and others. |
| positively charged metal ions of the lowest oxidation state | Metal ions of lower oxidation states can exhibit reducing properties if they are characterized by states with a higher oxidation state. For example, Sn 2+ -2e — → Sn 4+ Cr 2+ -e — → Cr 3+ Cu + -e — → Cu 2+ |
| Complex ions and molecules containing atoms in intermediate oxidation states | Complex or complex ions, as well as molecules, can exhibit reducing properties if their constituent atoms are in an intermediate oxidation state. For example, SO 3 2-, NO 2 -, AsO 3 3-, 4-, SO 2, CO, NO and others. |
| Carbon, Carbon monoxide (II), Iron, Zinc, Aluminum, Tin, Sulfurous acid, Sodium sulfite and bisulfite, Sodium sulfide, Sodium thiosulfate, Hydrogen, Electric current | |
| Oxidizing agents | |
| Neutral atoms | General scheme E + ne- → E n- Oxidizing agents are atoms of p-elements. Typical nonmetals are fluorine, oxygen, chlorine. The strongest oxidizing agents are halogens and oxygen. In the main subgroups of groups 7, 6, 5 and 4, the oxidative activity of atoms decreases from top to bottom |
| positively charged metal ions | All positively charged metal ions exhibit oxidizing properties to varying degrees. Of these, the most powerful oxidizing agents are ions with a high oxidation state, for example, Sn 4+, Fe 3+, Cu 2+. Noble metal ions, even in low oxidation states, are strong oxidizing agents. |
| Complex ions and molecules containing metal atoms in the highest oxidation state | Typical oxidizing agents are substances that contain metal atoms in the state of the highest oxidation state. For example, KMnO4, K2Cr2O7, K2CrO4, HAuCl4. |
| Complex ions and molecules containing non-metal atoms in a state of positive oxidation state | These are mainly oxygen-containing acids, as well as their corresponding oxides and salts. For example, SO 3, H 2 SO 4, HClO, HClO 3, NaOBr and others. In a row H 2SO4 →H 2SeO4 →H 6TeO6 oxidizing activity increases from sulfuric to telluric acid. In a row HClO -HClO2 -HClO 3 -HClO4 HBrO - HBrO 3 - HIO - HIO 3 - HIO 4 , H5IO 6 oxidative activity increases from right to left, and acidic properties increase from left to right. |
| The most important reducing agents in technology and laboratory practice | Oxygen, Ozone, Potassium permanganate, Chromic and dichromic acids, Nitric acid, Nitrous acid, Sulfuric acid (conc.), Hydrogen peroxide, Electric current, Hypochlorous acid, Manganese dioxide, Lead dioxide, Bleach, Solutions of potassium and sodium hypochlorites, Potassium hypobromide , Potassium hexacyanoferrate (III). |
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