Electronic configuration of an atom. Structure of the oxygen atom Quantum cells of chemical elements table
Beryllium Electronic configuration - 2s2. Electronic diagram of the outer quantum layer:
Bor Electronic configuration - 2s22р1. The boron atom can go into an excited state. Electronic diagram of the outer quantum layer:
An unexcited carbon atom can form two covalent bonds due to electron pairing and one through the donor-acceptor mechanism. An example of such a compound is carbon monoxide (II), which has the formula CO and is called carbon monoxide. Its structure will be discussed in more detail in section 2.1.2. An excited carbon atom is unique: all orbitals of its outer quantum layer are filled with unpaired electrons, i.e. It has the same number of valence orbitals and valence electrons. Its ideal partner is the hydrogen atom, which has one electron in its only orbital. This explains their ability to form hydrocarbons. Having four unpaired electrons, the carbon atom forms four chemical bonds: CH4, CF4, CO2. In molecules of organic compounds, the carbon atom is always in an excited state:
The nitrogen atom cannot be excited because there is no free orbital in its outer quantum layer. It forms three covalent bonds due to electron pairing:
Having two unpaired electrons in the outer layer, the oxygen atom forms two covalent bonds:
Neon
Electronic configuration - 2s22р6. Lewis symbol: Electron diagram of the outer quantum layer:
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The neon atom has a complete external energy level and does not form chemical bonds with any atoms. This is the second noble gas. THIRD PERIOD Atoms of all elements of the third period have three quantum layers. The electronic configuration of the two internal energy levels can be depicted as . The outer electronic layer contains nine orbitals, which are populated by electrons, obeying general laws. So, for a sodium atom the electronic configuration is: 3s1, for calcium - 3s2 (in an excited state - 3s13р1), for aluminum - 3s23р1 (in an excited state - 3s13р2). Unlike elements of the second period, atoms of elements of groups V – VII of the third period can exist both in the ground and in excited states. Phosphorus Phosphorus is a group 5 element. Its electronic configuration is 3s23р3. Like nitrogen, it has three unpaired electrons in its outermost energy level and forms three covalent bonds. An example is phosphine, which has the formula PH3 (compare with ammonia). But phosphorus, unlike nitrogen, contains free d-orbitals in the outer quantum layer and can go into an excited state - 3s13р3d1:
This gives it the opportunity to form five covalent bonds in compounds such as P2O5 and H3PO4.
However, it can be excited by transferring an electron first from R- on d-orbital (first excited state), and then with s- on d-orbital (second excited state):
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In the first excited state, the sulfur atom forms four chemical bonds in compounds such as SO2 and H2SO3. The second excited state of the sulfur atom can be depicted using an electron diagram:
This sulfur atom forms six chemical bonds in the compounds SO3 and H2SO4.
1.3.3. Electronic configurations of atoms of large elements periods THE FOURTH PERIODThe period begins with potassium (19K) electron configuration: 1s22s22p63s23p64s1 or 4s1 and calcium (20Ca): 1s22s22p63s23p64s2 or 4s2. Thus, in accordance with the Klechkovsky rule, after the p-orbitals of Ar, the outer 4s sublevel is filled, which has lower energy, because The 4s orbital penetrates closer to the nucleus; The 3d sublevel remains empty (3d0). Starting from scandium, the orbitals of the 3d sublevel are populated in 10 elements. They're called d-elements.
In accordance with the principle of sequential filling of orbitals, the chromium atom should have an electronic configuration of 4s23d4, but it exhibits an electron “leap”, which consists in the transition of a 4s electron to a 3d orbital that is close in energy (Fig. 11).
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It has been experimentally established that atomic states in which the p-, d-, f-orbitals are half filled (p3, d5, f7), completely (p6, d10, f14) or free (p0, d0, f0) have increased stability. Therefore, if an atom lacks one electron before half-completion or completion of a sublevel, its “leap” from a previously filled orbital (in this case, 4s) is observed.
With the exception of Cr and Cu, all elements from Ca to Zn have the same number of electrons in their outer shell - two. This explains the relatively small change in properties in the series of transition metals. However, for the listed elements, both the 4s electrons of the outer and the 3d electrons of the pre-external sublevel are valence electrons (with the exception of the zinc atom, in which the third energy level is completely completed).
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The 4d and 4f orbitals remained free, although the fourth period was completed.
FIFTH PERIOD
The sequence of filling the orbitals is the same as in the previous period: first the 5s orbital is filled ( 37Rb 5s1), then 4d and 5p ( 54Xe 5s24d105p6). The 5s and 4d orbitals are even closer in energy, so most 4d elements (Mo, Tc, Ru, Rh, Pd, Ag) experience an electron transition from the 5s to the 4d sublevel.
SIXTH AND SEVENTH PERIODS
Unlike the previous one, the sixth period includes 32 elements. Cesium and barium are 6s elements. The next energetically favorable states are 6p, 4f and 5d. Contrary to Klechkovsky's rule, in lanthanum it is not the 4f but the 5d orbital that is filled ( 57La 6s25d1), however, for the elements following it, the 4f-sublevel is filled ( 58Ce 6s24f2), on which there are fourteen possible electronic states. Atoms from cerium (Ce) to lutetium (Lu) are called lanthanides - these are f-elements. In the series of lanthanides, sometimes an electron “leak” occurs, just as in the series of d-elements. When the 4f-sublevel is completed, the 5d-sublevel (nine elements) continues to be filled and the sixth period, like any other except the first, is completed by six p-elements.
The first two s elements in the seventh period are francium and radium, followed by one 6d element, actinium ( 89Ac 7s26d1). Actinium is followed by fourteen 5f elements - actinides. The actinides should be followed by nine 6d elements and six p elements should complete the period. The seventh period is incomplete.
The considered pattern of the formation of periods of a system by elements and the filling of atomic orbitals with electrons shows the periodic dependence of the electronic structures of atoms on the charge of the nucleus.
Period is a set of elements arranged in order of increasing charges of atomic nuclei and characterized by the same value of the principal quantum number of outer electrons. At the beginning of the period are filled ns -, and at the end - n.p. -orbitals (except for the first period). These elements form eight main (A) subgroups of the periodic system of D.I. Mendeleev.
Main subgroup is a set of chemical elements arranged vertically and having the same number of electrons at the outer energy level.
Within the period, with an increase in the charge of the nucleus and an increasing force of attraction of external electrons to it from left to right, the radii of atoms decrease, which in turn causes a weakening of metallic properties and an increase in non-metallic properties. Behind atomic radius take the theoretically calculated distance from the nucleus to the maximum electron density of the outer quantum layer. In groups, from top to bottom, the number of energy levels increases, and, consequently, the atomic radius. At the same time, the metallic properties are enhanced. Important properties of atoms that change periodically depending on the charges of the atomic nuclei also include ionization energy and electron affinity, which will be discussed in section 2.2.
The filling of orbitals in a non-excited atom is carried out in such a way that the energy of the atom is minimal (the principle of minimum energy). First, the orbitals of the first energy level are filled, then the second, and the orbital of the s-sublevel is filled first and only then the orbitals of the p-sublevel. In 1925, the Swiss physicist W. Pauli established the fundamental quantum mechanical principle of natural science (the Pauli principle, also called the exclusion principle or the exclusion principle). According to the Pauli principle:
The electronic configuration of an atom is expressed by a formula in which the filled orbitals are indicated by a combination of a number equal to the principal quantum number and a letter corresponding to the orbital quantum number. The superscript indicates the number of electrons in these orbitals.An atom cannot have two electrons that have the same set of all four quantum numbers.
Hydrogen and helium
The electronic configuration of the hydrogen atom is 1s 1, and the helium atom is 1s 2. A hydrogen atom has one unpaired electron, and a helium atom has two paired electrons. Paired electrons have the same values of all quantum numbers except the spin one. A hydrogen atom can give up its electron and turn into a positively charged ion - the H + cation (proton), which has no electrons (electronic configuration 1s 0). A hydrogen atom can add one electron and become a negatively charged H - ion (hydride ion) with the electron configuration 1s 2.Lithium
The three electrons in a lithium atom are distributed as follows: 1s 2 1s 1. Only electrons from the outer energy level, called valence electrons, participate in the formation of a chemical bond. In a lithium atom, the valence electron is the 2s sublevel electron, and the two electrons of the 1s sublevel are internal electrons. The lithium atom quite easily loses its valence electron, transforming into the Li + ion, which has the 1s 2 2s 0 configuration. Note that the hydride ion, helium atom, and lithium cation have the same number of electrons. Such particles are called isoelectronic. They have similar electronic configurations but different nuclear charges. The helium atom is very chemically inert, which is due to the special stability of the 1s 2 electronic configuration. Orbitals that are not filled with electrons are called vacant. In the lithium atom, three orbitals of the 2p sublevel are vacant.Beryllium
The electronic configuration of the beryllium atom is 1s 2 2s 2. When an atom is excited, electrons from a lower energy sublevel move to vacant orbitals of a higher energy sublevel. The process of excitation of a beryllium atom can be conveyed by the following diagram:1s 2 2s 2 (ground state) + hν→ 1s 2 2s 1 2p 1 (excited state).
A comparison of the ground and excited states of the beryllium atom shows that they differ in the number of unpaired electrons. In the ground state of the beryllium atom there are no unpaired electrons; in the excited state there are two. Despite the fact that when an atom is excited, in principle, any electrons from lower energy orbitals can move to higher orbitals, for the consideration of chemical processes only transitions between energy sublevels with similar energies are significant.
This is explained as follows. When a chemical bond is formed, energy is always released, i.e., the combination of two atoms goes into an energetically more favorable state. The process of excitation requires energy expenditure. When pairing electrons within the same energy level, the excitation costs are compensated by the formation of a chemical bond. When pairing electrons within different levels, the excitation costs are so high that they cannot be compensated by the formation of a chemical bond. In the absence of a partner in a possible chemical reaction, the excited atom releases a quantum of energy and returns to the ground state - this process is called relaxation.
Bor
The electronic configurations of atoms of elements of the 3rd period of the Periodic Table of Elements will be to a certain extent similar to those given above (the subscript indicates the atomic number):
11 Na 3s 1
12 Mg 3s 2
13 Al 3s 2 3p 1
14 Si 2s 2 2p2
15P 2s 2 3p 3
However, the analogy is not complete, since the third energy level is split into three sublevels and all of the listed elements have vacant d-orbitals to which electrons can transfer upon excitation, increasing multiplicity. This is especially important for elements such as phosphorus, sulfur and chlorine.
The maximum number of unpaired electrons in a phosphorus atom can reach five:
This explains the possibility of the existence of compounds in which the valency of phosphorus is 5. The nitrogen atom, which has the same configuration of valence electrons in the ground state as the phosphorus atom, cannot form five covalent bonds.
A similar situation arises when comparing the valence capabilities of oxygen and sulfur, fluorine and chlorine. The pairing of electrons in a sulfur atom results in the appearance of six unpaired electrons:
3s 2 3p 4 (ground state) → 3s 1 3p 3 3d 2 (excited state).
This corresponds to the six-valence state, which is unattainable for oxygen. The maximum valency of nitrogen (4) and oxygen (3) requires a more detailed explanation, which will be given later.
The maximum valency of chlorine is 7, which corresponds to the configuration of the excited state of the atom 3s 1 3p 3 d 3.
The presence of vacant 3d orbitals in all elements of the third period is explained by the fact that, starting from the 3rd energy level, partial overlap of sublevels of different levels occurs when filled with electrons. Thus, the 3d sublevel begins to fill only after the 4s sublevel is filled. The energy reserve of electrons in atomic orbitals of different sublevels and, consequently, the order of their filling increases in the following order:
Orbitals for which the sum of the first two quantum numbers (n + l) is smaller are filled earlier; if these sums are equal, the orbitals with the lower principal quantum number are filled first.
This pattern was formulated by V. M. Klechkovsky in 1951.
Elements in whose atoms the s-sublevel is filled with electrons are called s-elements. These include the first two elements of each period: hydrogen. However, already in the next d-element - chromium - there is some “deviation” in the arrangement of electrons in energy levels in the ground state: instead of the expected four unpaired electrons on the 3d sublevel, the chromium atom has five unpaired electrons in the 3d sublevel and one unpaired electron in the s sublevel: 24 Cr 4s 1 3d 5 .
The phenomenon of the transition of one s-electron to the d-sublevel is often called “leakthrough” of an electron. This can be explained by the fact that the orbitals of the d-sublevel filled by electrons become closer to the nucleus due to increased electrostatic attraction between electrons and the nucleus. As a result, the state 4s 1 3d 5 becomes energetically more favorable than 4s 2 3d 4. Thus, the half-filled d-sublevel (d 5) has increased stability compared to other possible electron distribution options. The electronic configuration corresponding to the existence of the maximum possible number of paired electrons, achievable in previous d-elements only as a result of excitation, is characteristic of the ground state of the chromium atom. The electronic configuration d 5 is also characteristic of the manganese atom: 4s 2 3d 5. For the following d-elements, each energy cell of the d-sublevel is filled with a second electron: 26 Fe 4s 2 3d 6 ; 27 Co 4s 2 3d 7 ; 28 Ni 4s 2 3d 8 .
In the copper atom, the state of a completely filled d-sublevel (d 10) becomes achievable due to the transition of one electron from the 4s sub-level to the 3d sublevel: 29 Cu 4s 1 3d 10. The last element of the first row of d-elements has the electronic configuration 30 Zn 4s 23 d 10.
The general trend, manifested in the stability of the d 5 and d 10 configurations, is also observed in elements of lower periods. Molybdenum has an electronic configuration similar to chromium: 42 Mo 5s 1 4d 5, and silver to copper: 47 Ag5s 0 d 10. Moreover, the d 10 configuration is already achieved in palladium due to the transition of both electrons from the 5s orbital to the 4d orbital: 46Pd 5s 0 d 10. There are other deviations from the monotonic filling of d- and f-orbitals.
The electronic configuration of an element is a record of the distribution of electrons in its atoms across shells, subshells and orbitals. Electronic configuration is usually written for atoms in their ground state. The electronic configuration of an atom in which one or more electrons are in an excited state is called the excited configuration. To determine the specific electronic configuration of an element in the ground state, the following three rules exist: Rule 1: filling principle. According to the filling principle, electrons in the ground state of an atom fill orbitals in a sequence of increasing orbital energy levels. The lowest energy orbitals are always filled first.
Hydrogen; atomic number = 1; number of electrons = 1
This single electron in the hydrogen atom must occupy the s orbital of the K-shell, since it has the lowest energy of all possible orbitals (see Fig. 1.21). The electron in this s orbital is called an ls electron. Hydrogen in its ground state has an electronic configuration of Is1.
Rule 2: Pauli's exclusion principle. According to this principle, any orbital can contain no more than two electrons, and then only if they have opposite spins (unequal spin numbers).
Lithium; atomic number = 3; number of electrons = 3
The lowest energy orbital is the 1s orbital. It can only accept two electrons. These electrons must have unequal spins. If we denote spin +1/2 with an arrow pointing up, and spin -1/2 with an arrow pointing down, then two electrons with opposite (antiparallel) spins in the same orbital can be schematically represented by the notation (Fig. 1.27)
Two electrons with identical (parallel) spins cannot exist in one orbital:
The third electron in a lithium atom must occupy the orbital next in energy to the lowest orbital, i.e. 2b-orbital. Thus, lithium has an electronic configuration of Is22s1.
Rule 3: Hund's rule. According to this rule, the filling of the orbitals of one subshell begins with single electrons with parallel (equal sign) spins, and only after single electrons occupy all the orbitals can the final filling of the orbitals with pairs of electrons with opposite spins occur.
Nitrogen; atomic number = 7; number of electrons = 7 Nitrogen has an electron configuration of ls22s22p3. The three electrons located on the 2p subshell must be located singly in each of the three 2p orbitals. In this case, all three electrons must have parallel spins (Fig. 1.22).
In table Figure 1.6 shows the electronic configurations of elements with atomic numbers from 1 to 20.
Table 1.6. Ground state electronic configurations for elements with atomic number 1 to 20
Initially, the elements in the Periodic Table of Chemical Elements by D.I. Mendeleev's atoms were arranged in accordance with their atomic masses and chemical properties, but in fact it turned out that the decisive role is played not by the mass of the atom, but by the charge of the nucleus and, accordingly, the number of electrons in a neutral atom.
The most stable state of an electron in an atom of a chemical element corresponds to the minimum of its energy, and any other state is called excited, in which the electron can spontaneously move to a level with a lower energy.
Let's consider how electrons in an atom are distributed among orbitals, i.e. electronic configuration of a multielectron atom in the ground state. To construct the electronic configuration, the following principles are used for filling orbitals with electrons:
- Pauli principle (prohibition) - in an atom there cannot be two electrons with the same set of all 4 quantum numbers;
- the principle of least energy (Klechkovsky's rules) - the orbitals are filled with electrons in order of increasing energy of the orbitals (Fig. 1).
Rice. 1. Energy distribution of orbitals of a hydrogen-like atom; n is the principal quantum number.
The energy of the orbital depends on the sum (n + l). The orbitals are filled with electrons in order of increasing sum (n + l) for these orbitals. Thus, for the 3d and 4s sublevels, the sums (n + l) will be equal to 5 and 4, respectively, as a result of which the 4s orbital will be filled first. If the sum (n + l) is the same for two orbitals, then the orbital with the smaller n value is filled first. So, for 3d and 4p orbitals, the sum (n + l) will be equal to 5 for each orbital, but the 3d orbital is filled first. According to these rules, the order of filling the orbitals will be as follows:
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<5d<4f<6p<7s<6d<5f<7p
An element's family is determined by the last orbital to be filled by electrons, according to energy. However, it is impossible to write electronic formulas in accordance with the energy series.
41 Nb 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 3 5s 2 correct notation of electronic configuration
41 Nb 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 3 incorrect electronic configuration entry
For the first five d - elements, the valence (i.e., electrons responsible for the formation of a chemical bond) is the sum of the electrons on d and s, the last ones filled with electrons. For p-elements, the valence is the sum of the electrons located in the s and p sublevels. For s elements, the valence electrons are the electrons located in the s sublevel of the outer energy level.
- Hund's rule - at one value of l, electrons fill the orbitals in such a way that the total spin is maximum (Fig. 2)
Rice. 2. Change in energy in the 1s -, 2s – 2p – orbitals of atoms of the 2nd period of the Periodic Table.
Examples of constructing electronic configurations of atoms
Examples of constructing electronic configurations of atoms are given in Table 1.
Table 1. Examples of constructing electronic configurations of atoms
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Electronic configuration |
Applicable rules |
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Pauli principle, Kleczkowski rules |
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Hund's rule |
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1s 2 2s 2 2p 6 4s 1 |
Klechkovsky's rules |
Electronic configuration an atom is a numerical representation of its electron orbitals. Electron orbitals are regions of various shapes located around the atomic nucleus in which it is mathematically probable that an electron will be found. Electronic configuration helps quickly and easily tell the reader how many electron orbitals an atom has, as well as determine the number of electrons in each orbital. After reading this article, you will master the method of drawing up electronic configurations.
Steps
Distribution of electrons using the periodic system of D. I. Mendeleev
- For example, a sodium atom with charge -1 will have an extra electron in addition to its base atomic number 11. In other words, the atom will have a total of 12 electrons.
- If we are talking about a sodium atom with a charge of +1, one electron must be subtracted from the base atomic number 11. Thus, the atom will have 10 electrons.
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Remember the basic list of orbitals. As the number of electrons in an atom increases, they fill the various sublevels of the atom's electron shell according to a specific sequence. Each sublevel of the electron shell, when filled, contains an even number of electrons. The following sublevels are available:
Understand electronic configuration notation. Electron configurations are written to clearly show the number of electrons in each orbital. Orbitals are written sequentially, with the number of atoms in each orbital written as a superscript to the right of the orbital name. The completed electronic configuration takes the form of a sequence of sublevel designations and superscripts.
- Here, for example, is the simplest electronic configuration: 1s 2 2s 2 2p 6 . This configuration shows that there are two electrons in the 1s sublevel, two electrons in the 2s sublevel, and six electrons in the 2p sublevel. 2 + 2 + 6 = 10 electrons in total. This is the electronic configuration of a neutral neon atom (neon's atomic number is 10).
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Remember the order of the orbitals. Keep in mind that electron orbitals are numbered in order of increasing electron shell number, but arranged in increasing order of energy. For example, a filled 4s 2 orbital has lower energy (or less mobility) than a partially filled or filled 3d 10 orbital, so the 4s orbital is written first. Once you know the order of the orbitals, you can easily fill them according to the number of electrons in the atom. The order of filling the orbitals is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
- The electronic configuration of an atom in which all orbitals are filled will be as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6
- Note that the above entry, when all orbitals are filled, is the electron configuration of element Uuo (ununoctium) 118, the highest numbered atom in the periodic table. Therefore, this electronic configuration contains all the currently known electronic sublevels of a neutrally charged atom.
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Fill the orbitals according to the number of electrons in your atom. For example, if we want to write down the electron configuration of a neutral calcium atom, we must start by looking up its atomic number in the periodic table. Its atomic number is 20, so we will write the configuration of an atom with 20 electrons according to the above order.
- Fill the orbitals according to the order above until you reach the twentieth electron. The first 1s orbital will have two electrons, the 2s orbital will also have two, the 2p will have six, the 3s will have two, the 3p will have 6, and the 4s will have 2 (2 + 2 + 6 +2 +6 + 2 = 20 .) In other words, the electronic configuration of calcium has the form: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 .
- Note that the orbitals are arranged in order of increasing energy. For example, when you are ready to move to the 4th energy level, first write down the 4s orbital, and then 3d. After the fourth energy level, you move to the fifth, where the same order is repeated. This happens only after the third energy level.
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Use the periodic table as a visual cue. You've probably already noticed that the shape of the periodic table corresponds to the order of the electron sublevels in the electron configurations. For example, the atoms in the second column from the left always end in "s 2", and the atoms on the right edge of the thin middle part always end in "d 10", etc. Use the periodic table as a visual guide to writing configurations - how the order in which you add to the orbitals corresponds to your position in the table. See below:
- Specifically, the leftmost two columns contain atoms whose electronic configurations end in s orbitals, the right block of the table contains atoms whose configurations end in p orbitals, and the bottom half contains atoms that end in f orbitals.
- For example, when you write down the electronic configuration of chlorine, think like this: "This atom is located in the third row (or "period") of the periodic table. It is also located in the fifth group of the p orbital block of the periodic table. Therefore, its electronic configuration will end with. ..3p 5
- Note that elements in the d and f orbital region of the table are characterized by energy levels that do not correspond to the period in which they are located. For example, the first row of a block of elements with d-orbitals corresponds to 3d orbitals, although it is located in the 4th period, and the first row of elements with f-orbitals corresponds to a 4f orbital, despite being in the 6th period.
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Learn abbreviations for writing long electron configurations. The atoms on the right edge of the periodic table are called noble gases. These elements are chemically very stable. To shorten the process of writing long electron configurations, simply write the chemical symbol of the nearest noble gas with fewer electrons than your atom in square brackets, and then continue writing the electron configuration of subsequent orbital levels. See below:
- To understand this concept, it will be helpful to write an example configuration. Let's write the configuration of zinc (atomic number 30) using the abbreviation that includes the noble gas. The complete configuration of zinc looks like this: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10. However, we see that 1s 2 2s 2 2p 6 3s 2 3p 6 is the electron configuration of argon, a noble gas. Simply replace part of the electronic configuration for zinc with the chemical symbol for argon in square brackets (.)
- So, the electronic configuration of zinc, written in abbreviated form, has the form: 4s 2 3d 10 .
- Please note that if you are writing the electronic configuration of a noble gas, say argon, you cannot write it! One must use the abbreviation for the noble gas preceding this element; for argon it will be neon ().
Using the periodic table ADOMAH
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Master the periodic table ADOMAH. This method of recording the electronic configuration does not require memorization, but requires a modified periodic table, since in the traditional periodic table, starting from the fourth period, the period number does not correspond to the electron shell. Find the periodic table ADOMAH - a special type of periodic table developed by scientist Valery Zimmerman. It is easy to find with a short internet search.
- In the ADOMAH periodic table, the horizontal rows represent groups of elements such as halogens, noble gases, alkali metals, alkaline earth metals, etc. Vertical columns correspond to electronic levels, and so-called "cascades" (diagonal lines connecting blocks s, p, d and f) correspond to periods.
- Helium is moved towards hydrogen because both of these elements are characterized by a 1s orbital. The period blocks (s,p,d and f) are shown on the right side, and the level numbers are given at the bottom. Elements are represented in boxes numbered 1 to 120. These numbers are ordinary atomic numbers, which represent the total number of electrons in a neutral atom.
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Find your atom in the ADOMAH table. To write the electronic configuration of an element, look up its symbol on the periodic table ADOMAH and cross out all elements with a higher atomic number. For example, if you need to write the electron configuration of erbium (68), cross out all elements from 69 to 120.
- Note the numbers 1 through 8 at the bottom of the table. These are numbers of electronic levels, or numbers of columns. Ignore columns that contain only crossed out items. For erbium, columns numbered 1,2,3,4,5 and 6 remain.
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Count the orbital sublevels up to your element. Looking at the block symbols shown to the right of the table (s, p, d, and f) and the column numbers shown at the base, ignore the diagonal lines between the blocks and break the columns into column blocks, listing them in order from bottom to top. Again, ignore blocks that have all the elements crossed out. Write column blocks starting from the column number followed by the block symbol, thus: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s (for erbium).
- Please note: The above electron configuration of Er is written in ascending order of electron sublevel number. It can also be written in order of filling the orbitals. To do this, follow the cascades from bottom to top, rather than columns, when you write column blocks: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 12 .
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Count the electrons for each electron sublevel. Count the elements in each column block that have not been crossed out, attaching one electron from each element, and write their number next to the block symbol for each column block thus: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 12 5s 2 5p 6 6s 2 . In our example, this is the electronic configuration of erbium.
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Be aware of incorrect electronic configurations. There are eighteen typical exceptions that relate to the electronic configurations of atoms in the lowest energy state, also called the ground energy state. They do not obey the general rule only for the last two or three positions occupied by electrons. In this case, the actual electronic configuration assumes that the electrons are in a state with a lower energy compared to the standard configuration of the atom. Exception atoms include:
- Cr(..., 3d5, 4s1); Cu(..., 3d10, 4s1); Nb(..., 4d4, 5s1); Mo(..., 4d5, 5s1); Ru(..., 4d7, 5s1); Rh(..., 4d8, 5s1); Pd(..., 4d10, 5s0); Ag(..., 4d10, 5s1); La(..., 5d1, 6s2); Ce(..., 4f1, 5d1, 6s2); Gd(..., 4f7, 5d1, 6s2); Au(..., 5d10, 6s1); Ac(..., 6d1, 7s2); Th(..., 6d2, 7s2); Pa(..., 5f2, 6d1, 7s2); U(..., 5f3, 6d1, 7s2); Np(..., 5f4, 6d1, 7s2) and Cm(..., 5f7, 6d1, 7s2).
- To find the atomic number of an atom when it is written in electron configuration form, simply add up all the numbers that follow the letters (s, p, d, and f). This only works for neutral atoms, if you're dealing with an ion it won't work - you'll have to add or subtract the number of extra or lost electrons.
- The number following the letter is a superscript, do not make a mistake in the test.
- There is no "half-full" sublevel stability. This is a simplification. Any stability that is attributed to "half-filled" sublevels is due to the fact that each orbital is occupied by one electron, thus minimizing repulsion between electrons.
- Each atom tends to a stable state, and the most stable configurations have the s and p sublevels filled (s2 and p6). Noble gases have this configuration, so they rarely react and are located on the right in the periodic table. Therefore, if a configuration ends in 3p 4, then it needs two electrons to reach a stable state (to lose six, including the s-sublevel electrons, requires more energy, so losing four is easier). And if the configuration ends in 4d 3, then to achieve a stable state it needs to lose three electrons. In addition, half-filled sublevels (s1, p3, d5..) are more stable than, for example, p4 or p2; however, s2 and p6 will be even more stable.
- When you are dealing with an ion, this means that the number of protons is not equal to the number of electrons. The charge of the atom in this case will be depicted at the top right (usually) of the chemical symbol. Therefore, an antimony atom with charge +2 has the electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 . Note that 5p 3 has changed to 5p 1 . Be careful when the neutral atom configuration ends in sublevels other than s and p. When you take away electrons, you can only take them from the valence orbitals (s and p orbitals). Therefore, if the configuration ends with 4s 2 3d 7 and the atom receives a charge of +2, then the configuration will end with 4s 0 3d 7. Please note that 3d 7 Not changes, electrons from the s orbital are lost instead.
- There are conditions when an electron is forced to "move to a higher energy level." When a sublevel is one electron short of being half or full, take one electron from the nearest s or p sublevel and move it to the sublevel that needs the electron.
- There are two options for recording the electronic configuration. They can be written in increasing order of energy level numbers or in the order of filling electron orbitals, as was shown above for erbium.
- You can also write the electronic configuration of an element by writing only the valence configuration, which represents the last s and p sublevel. Thus, the valence configuration of antimony will be 5s 2 5p 3.
- Ions are not the same. It's much more difficult with them. Skip two levels and follow the same pattern depending on where you started and how large the number of electrons is.
Find the atomic number of your atom. Each atom has a certain number of electrons associated with it. Find your atom's symbol on the periodic table. The atomic number is a positive integer starting at 1 (for hydrogen) and increasing by one for each subsequent atom. Atomic number is the number of protons in an atom, and therefore it is also the number of electrons of an atom with zero charge.
Determine the charge of an atom. Neutral atoms will have the same number of electrons as shown on the periodic table. However, charged atoms will have more or less electrons, depending on the magnitude of their charge. If you are working with a charged atom, add or subtract electrons as follows: add one electron for each negative charge and subtract one for each positive charge.
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