El formulas of chemical elements. Electronic formula of the element

To learn how to compose electron graphic formulas, it is important to understand the theory of the structure of the atomic nucleus. The nucleus of an atom is made up of protons and neutrons. There are electrons around the nucleus of an atom in electron orbitals.

You will need

• - pen;
• - paper for notes;
• - periodic table of elements (periodic table).

Instructions

Electrons in an atom occupy vacant orbitals in a sequence called the energy scale: 1s/2s, 2p/3s, 3p/4s, 3d, 4p/5s, 4d, 5p/6s, 4d, 5d, 6p/7s, 5f, 6d, 7p . One orbital can contain two electrons with opposite spins - directions of rotation.

The structure of electron shells is expressed using graphical electronic formulas. Use a matrix to write the formula. One or two electrons with opposite spins can be located in one cell. Electrons are represented by arrows. The matrix clearly shows that two electrons can be located in the s orbital, 6 electrons in the p orbital, 10 in the d orbital, and -14 in the f orbital.

Consider the principle of drawing up an electronic graphic formula using manganese as an example. Find manganese in the periodic table. Its atomic number is 25, which means there are 25 electrons in the atom, it is an element of the fourth period.

Write down the serial number and symbol of the element next to the matrix. In accordance with the energy scale, fill the 1s, 2s, 2p, 3s, 3p, 4s levels in succession, writing two electrons per cell. You get 2+2+6+2+6+2=20 electrons. These levels are completely filled.

You still have five electrons left and an unfilled 3d level. Arrange the electrons in the d-sublevel cells, starting from the left. Place electrons with the same spins in the cells, one at a time. If all the cells are filled, starting from the left, add a second electron with the opposite spin. Manganese has five d electrons, one in each cell.

Electron graphic formulas clearly show the number of unpaired electrons that determine valence.

When creating theoretical and practical works in mathematics, physics, chemistry, a student or schoolchild is faced with the need to insert special characters and complex formulas. With the Word application from the Microsoft office suite, you can type an email formula of any complexity.

Instructions

Open a new document in Microsoft Word. Give it a name and save it in the same folder where you have your work so you don’t have to look for it in the future.

Go to the "Insert" tab. On the right, find the symbol?, and next to it is the inscription “Formula”. Click on the arrow. A window will appear where you can select a built-in formula, such as a quadratic formula.

Click on the arrow and a variety of symbols will appear on the top panel that you may need when writing this particular formula. After changing it the way you need, you can save it. From now on, it will appear in the list of built-in formulas.

If you need to transfer a formula into text that you later need to place on the site, then right-click on the active field with it and select not the professional, but the linear writing method. In particular, the formula of the same quadratic equation in this case will take the form: x=(-b±?(b^2-4ac))/2a.

Another option for writing an electronic formula in Word is through the constructor. Hold down the Alt and = keys at the same time. You will immediately have a field for writing a formula, and a constructor will open in the top panel. Here you can select all the signs that may be needed to write an equation and solve any problem.

Some linear notation symbols may not be clear to a reader unfamiliar with computer symbology. In this case, it makes sense to save the most complex formulas or equations in graphical form. To do this, open the simplest graphic editor Paint: “Start” - “Programs” - “Paint”. Then zoom in on your formula document until it fills the entire screen. This is necessary so that the saved image has the highest resolution. Press PrtScr on your keyboard, go to Paint and press Ctrl+V.

Trim off any excess. As a result, you will get a high-quality image with the desired formula.

Please note

Remember that chemistry is a science of exceptions. In atoms of side subgroups of the Periodic Table, electron “leakage” occurs. For example, in chromium with atomic number 24, one of the electrons from the 4s level goes into the d-level cell. A similar effect occurs in molybdenum, niobium, etc. In addition, there is the concept of an excited state of an atom, when paired electrons are paired and transferred to neighboring orbitals. Therefore, when compiling electronic graphic formulas for the elements of the fifth and subsequent periods of the secondary subgroup, check the reference book.

The electronic structure of an atom can be shown by an electronic formula and an electron graphic diagram. In electronic formulas, energy levels and sublevels are written sequentially in the order in which they are filled and the total number of electrons in the sublevel. In this case, the state of an individual electron, in particular its magnetic and spin quantum numbers, is not reflected in the electronic formula. In electronic graphic circuits, each electron is “visible” completely, i.e. it can be characterized by all four quantum numbers. Electron graphic diagrams are usually given for external electrons.

Example 1. Write the electronic formula of fluorine, express the state of the outer electrons with an electronic graphic diagram. How many unpaired electrons are there in an atom of this element?

Solution. The atomic number of fluorine is nine, therefore, its atom has nine electrons. In accordance with the principle of least energy, using Fig. 7 and taking into account the consequences of the Pauli principle, we write the electronic formula of fluorine: 1s 2 2s 2 2p 5. For the outer electrons (second energy level), we draw up an electron graphic diagram (Fig. 8), from which it follows that the fluorine atom has one unpaired electron.

Rice. 8. Electron graphic diagram of valence electrons of a fluorine atom

Example 2. Make electronic graphic diagrams of possible states of the nitrogen atom. Which of them reflect a normal state, and which ones reflect an excited state?

Solution. The electronic formula of nitrogen is 1s 2 s 2 2p 3, the formula of outer electrons is: 2s 2 2p 3. Sublevel 2p is incomplete because the number of electrons on it is less than six. Possible options The distributions of three electrons at the 2p sublevel are shown in Fig. 9.

Rice. 9. Electron graphic diagrams of possible states of the 2p sublevel in the nitrogen atom.

Maximum (by absolute value) the spin value (3 / 2) corresponds to states 1 and 2, therefore, they are ground, and the rest are excited.

Example 3. Determine the quantum numbers that determine the state of the last electron in the vanadium atom?

Solution. The atomic number of vanadium is Z = 23, therefore, the complete electronic formula of the element is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3. The electronic graphic diagram of external electrons (4s 2 3d 3) is as follows (Fig. 10):

Rice. 10. Electron graphic diagram of the valence electrons of the vanadium atom

Principal quantum number of the last electron n = 3 (third energy level), orbital l= 2 (sublevel d). The magnetic quantum number for each of the three d-electrons is different: for the first it is –2, for the second –1, for the third – 0. The spin quantum number for all three electrons is the same: m s = + 1 / 2. Thus, the state of the last electron in a vanadium atom is characterized by quantum numbers: n = 3; l= 2; m = 0; m s = + 1 / 2 .

7. Paired and unpaired electrons

Electrons that fill orbitals in pairs are called paired, and single electrons are called unpaired. Unpaired electrons provide chemical bonds between an atom and other atoms. The presence of unpaired electrons is established experimentally by studying magnetic properties. Substances with unpaired electrons paramagnetic(they are drawn into a magnetic field due to the interaction of electron spins, like elementary magnets, with an external magnetic field). Substances that have only paired electrons diamagnetic(external magnetic field does not affect them). Unpaired electrons are found only at the outer energy level of the atom and their number can be determined from its electron-graphic diagram.

Example 4. Determine the number of unpaired electrons in a sulfur atom.

Solution. The atomic number of sulfur is Z = 16, therefore the full electronic formula of the element is: 1s 2 2s 2 2p 6 3s 2 3p 4. The electronic graphic diagram of external electrons is as follows (Fig. 11).

Rice. 11. Electron graphic diagram of valence electrons of a sulfur atom

From the electron graphic diagram it follows that the sulfur atom has two unpaired electrons.

Electronic configuration of 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

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.

• 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.

1. 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).

2. 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.

3. 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.

4. 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 section 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.

5. 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

   1. 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 the 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.

   2. Find your atom in the ADOMAH table. To write the electron 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.

   3. 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 .

   4. 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.

   5. 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 don't obey general rule only in 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 occurs because each orbital is occupied by one electron, so repulsion between electrons is minimized.
   • 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 ascending 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.

According to the ideas of Heitler and London, the valence of elements is determined by the number of unpaired electrons. Let's consider the electronic graphic formulas of some elements, in which the orbitals are represented in the form of square cells, and the electron in the form of arrows + ½; -1/2.

From these formulas it follows that in the normal (unpaired) state, carbon has valence II, Sc – I. Atoms can go into an excited state, in which the lower lying sublevels can go from the lower lying sublevels to the higher lying empty sublevels (within one sublevel).

6. Periodic law and periodic system D.I. Mendeleev Structure of the periodic system (period, group, subgroup). The meaning of the periodic law and the periodic system.

Periodic law D.I. Mendeleev:Properties simple tel, A Also forms And properties connectopinions elements are V periodic dependencies from quantities atomic scales elements.(The properties of elements are periodically dependent on the charge of the atoms of their nuclei).

Periodic table of elements. Series of elements within which properties change sequentially, such as the series of eight elements from lithium to neon or from sodium to argon, Mendeleev called periods. If we write these two periods one below the other so that sodium is under lithium and argon is under neon, we get the following arrangement of elements:

With this arrangement, the vertical columns contain elements that are similar in their properties and have the same valency, for example, lithium and sodium, beryllium and magnesium, etc.

Having divided all the elements into periods and placing one period under another so that elements similar in properties and type of compounds formed were located under each other, Mendeleev compiled a table that he called the periodic system of elements by groups and series.

The meaning of the periodic systemWe. The periodic table of elements had a great influence on the subsequent development of chemistry. Not only was it the first natural classification chemical elements, which showed that they form a coherent system and are in close connection with each other, but also became a powerful tool for further research.

7. Periodic changes in the properties of chemical elements. Atomic and ionic radii. Ionization energy. Electron affinity. Electronegativity.

The dependence of atomic radii on the charge of the nucleus of an atom Z is periodic. Within one period, as Z increases, there is a tendency for the size of the atom to decrease, which is especially clearly observed in short periods

With the beginning of the construction of a new electronic layer, more distant from the nucleus, i.e., during the transition to the next period, atomic radii increase (compare, for example, the radii of fluorine and sodium atoms). As a result, within a subgroup, with increasing nuclear charge, the sizes of atoms increase.

The loss of electron atoms leads to a decrease in its effective size, and the addition of excess electrons leads to an increase. Therefore, the radius of a positively charged ion (cation) is always smaller, and the radius of a negatively charged non (anion) is always greater than the radius of the corresponding electrically neutral atom.

Within one subgroup, the radii of ions of the same charge increase with increasing nuclear charge. This pattern is explained by an increase in the number of electronic layers and the growing distance of outer electrons from the nucleus.

The most characteristic chemical property of metals is the ability of their atoms to easily give up external electrons and transform into positively charged ions, while non-metals, on the contrary, are characterized by the ability to gain electrons to form negative ions. To remove an electron from an atom and transform the latter into a positive ion, it is necessary to expend some energy, called ionization energy.

Ionization energy can be determined by bombarding atoms with electrons accelerated in an electric field. The lowest field voltage at which the electron speed becomes sufficient to ionize atoms is called the ionization potential of the atoms of a given element and is expressed in volts.

With the expenditure of sufficient energy, two, three or more electrons can be removed from an atom. Therefore, they speak of the first ionization potential (the energy of the removal of the first electron from the atom) and the second ionization potential (the energy of the removal of the second electron)

As noted above, atoms can not only donate, but also gain electrons. The energy released when an electron attaches to a free atom is called the atom's electron affinity. Electron affinity, like ionization energy, is usually expressed in electron volts. Thus, the electron affinity of the hydrogen atom is 0.75 eV, oxygen - 1.47 eV, fluorine - 3.52 eV.

The electron affinities of metal atoms are typically close to zero or negative; From this it follows that for atoms of most metals the addition of electrons is energetically unfavorable. The electron affinity of nonmetal atoms is always positive and the greater, the closer the nonmetal is located to the noble gas in the periodic table; this indicates an increase in non-metallic properties as the end of the period approaches.

Practical work

1. Basic provisions

Periodic table of chemical elements and atomic structure

Modern definition of the Periodic Law

The properties of chemical elements and the substances they form periodically depend on the charges of their atomic nuclei

The periodic table of chemical elements graphically displays the periodic law.

Each number in it characterizes some feature in the structure of atoms:

A)ordinal The (atomic) number of a chemical element indicates the charge of its atomic nucleus, that is, the number of protons contained in it, and since the atom is electrically neutral, the number of electrons located around the atomic nucleus.

Number of neutrons determined by the formula:N = A - Z ,

WhereA - mass number (atomic mass),Z - serial number of the element;

b) the period number corresponds to the number of energy levels (elector layers) in the atoms of elements of a given period;

c) the group number corresponds to the number of electrons on the outer level for elements of the main subgroups and the maximum number of valence electrons for elements of side subgroups.

Changes in metallic and non-metallic properties of elements

in periods and groups

1. Within one period With increasing serial number, the metallic properties of elements weaken, and non-metallic properties increase, since:

1) the number ē at the external level of atoms increases (it is equal to the group number);

2) the number of energy levels within a period does not change (it is equal to the period number);

3) the radius of atoms decreases.

2. Within the same group (main subgroup) As the serial number increases, the metallic properties of the elements increase, and the non-metallic ones weaken, since:

1) the number of electrons on the outer level of atoms is the same (it is equal to the group number);

2) the number of energy levels in atoms increases (it is equal to the period number);

3) the radius of the atoms increases.

Evidence of the complexity of the atomic structure

1. Irish physicist Stoney introduced the concept of “electron” to designate particles (for example, electrification of an ebonite stick), the appearance of static electricity on clothing.

2. Cathode rays - a stream of electrons from the metal atoms from which the cathode is made - caused the glass to glow (Thomson and Perrin). The negative charge of the electron was established. This smallest charge is taken as one = -1.

Thomson also established its mass equal to 1/1840 of the mass of a hydrogen atom.

3. Radioactivity is a phenomenon discovered by A. Becquerel. There are 3 types of radioactive rays:

a) α – rays consisting of α – particles with charge +2 and mass 4;

b) β – rays – electron flow; c) γ – rays – electromagnetic waves.

Therefore, the atom is divisible and has a complex structure.

Table 1Planetary model of the atom (Rutherford)

Equal to the number of nucleons (sum of protons and neutrons)

1) p + (have mass = 1 and charge = +1)

Their number is equal to the element number;

2) n 0 (have mass = 1 and charge = 0)

Their numberN = A r – Z. ( Z– number of protons)

Electronic shell

Consists of electrons

(mass tends to zero and charge = -1);

Their number is equal to the element number.

All the mass of an atom is concentrated in the nucleus

The atom is electrically neutral

Atom - electrically neutral system of interacting elementary particles, consisting of a nucleus (formed by protons and neutrons) and electrons

Structure of electronic shells of atoms

The concept of the electron shell of an atom and energy levels

1. Electronic shell – the collection of electrons surrounding an atomic nucleus.

2. In the electron shell, there are layers on which electrons with different amounts of energy are located, they are calledenergy levels . The number of these levels is equal to the period number in the periodic table.

3. The space around the nucleus in which the electron is most likely to be found (about 90%) is calledorbital .

Size and shape of orbitals

1) s 2 - electrons; spherical, symmetrical about the nucleus and has no direction.

2) p 6 – electrons; dumbbell-shaped, located mutually perpendicular in the atom

There are orbitals of more complex shapes:d 10 - orbitals andf 14 - orbitals.

The number of energy levels (electronic layers) in an atom is equal to the period number in the D.I. system. Mendeleev, to which a chemical element belongs: atoms of elements of the first period have one energy level, the second period has two, the third period has three, the seventh period has seven.

The largest number of electrons at an energy level is determined by the formula:

N = 2 n 2 , WhereN- maximum number of electrons;

n- level number or main quantum number. (Integern, indicating the energy level number, is calledprincipal quantum number ).

Energy levels and electronic configuration of an atom

The atom has a complex structure. It consists of a nucleus, which contains protons and neutrons, and electrons orbiting around the nucleus of the atom. The charge of a proton is +1, and the mass is 1.u. Neutron is an electrically neutral particle, mass approximately 1.0. Electron - charge is -1, mass 5.5∙10 -4 c.u. In general, an atom is electrically neutral; the number of protons in the nucleus of an atom is equal to the number of electrons in the atom. Electrons in an atom are distributed at energy levels.

The number of energy levels in an atom is determined by the number of the period in which the element is located. When constructing electronic models of atoms, it should be remembered that the maximum number of electrons at an energy level is 2n 2 , Wheren– energy level number. In accordance with this, the first level closest to the nucleus can contain no more than 2 electrons, the second - no more than 8, the third - no more than 18, and the fourth - no more than 32. The outer energy level cannot have more than 8 electrons.

Atomic absorption and emission spectra clearly show that all atoms have a range of possible energy states, called ground and excited electronic states (Fig. 1).

Recording the distribution of electrons in an atom across electronic levels and sublevels is called its electronic configuration and can be made for both the ground and excited states of the atom. To determine the specific electronic configuration of an atom in the ground state, the following three positions exist:

Filling principle (lowest energy). Electrons in the ground state fill orbitals in a sequence of increasing orbital energy levels. The lowest energy orbitals are always filled first.

Pauli's principle. Any orbital can contain no more than two electrons, and with oppositely directed spins (spin is a special property of an electron that has no analogues in the macrocosm, which can be simplified as the rotation of an electron around its own axis).

Hund's rule. Degenerate (with the same energy) orbitals are filled with single electrons with identically directed spins, only after this the degenerate orbitals are filled with electrons with oppositely directed spins according to the Pauli principle.

Quantum numbers

Principal quantum number n is equivalent to the quantum number in Bohr's theory. It basically determines the energy of electrons in a given orbital.

.....

....

Orbital quantum number l determines the value of the orbital angular momentum of an electron in a given orbital. Valid values: 0, 1, 2, 3, ... , n-1.

This quantum number describes the behavior of an atomic orbital during rotation of the coordinate system centered on the atomic nucleus.

Orbital magnetic quantum number m l determines the value of the component of the projection of the angular momentum of the electron onto a selected direction in space. In the absence of an external magnetic field, electrons are in orbitals with the same orbital quantum numberl energetically equivalent (i.e. their energy levels are degenerate).

However, in a constant magnetic field, some spectral lines are split. This means that the electrons become energetically unequal. For example, p-states in a magnetic field take 3 values ​​instead of one, d-states take 5 values. Valid values ​​of m l for thisl : - l , ... -2, -1, 0, +1, +2, ... + l

Spin quantum number m s is associated with the presence of an electron’s own magnetic moment. IN general view The expression for the magnetic moment of momentum coincides with that for the orbital moment:

For electron m s takes only two values: +1/2 and -1/2. Sometimes, to more clearly explain the concept of spin, a rough analogy is used - an electron is represented as a flying top (a circular current that creates its own magnetic field). This analogy allows us to explain the presence of spin 1/2 for the electron and proton, but not for the neutron - particles with zero charge.

The concept of “spin” does not fit into our “macro-ideas” of space. In all ways of registering it, the spin is always directed along the axis that the observer chose as the initial one. A spin value of 1/2 means that an electron (proton, neutron) becomes identical to itself after a revolution of 720 0 , not 360 0 , as in our three-dimensional world. Spin is considered to be one of the fundamental properties of nature (i.e., it is irreducible, like gravity and electricity).

Each orbital is designated by a square cell, electrons – by oppositely directed arrows (see the solution to the exercises on this topic)

Electronic formula is a formula that shows the distribution of electrons on the electron layers in an atom.

Table 2

Main quantum number, types and number of orbitals, maximum number of electrons in sublevels and levels

(period number)

n

The number of sublevels is equal to n

Shape (type) of orbitals

Number of orbitals

Maximum number of electrons

in the sublevel

in level equal n 2

at sublevels

at levels

TO (n=1)

1 s

Practical work

Purpose of the work:

6) Conclusion

Task No. 1

6 . Chargeatomic nucleus, Z

7. Mass number, A

8. Number of neutrons,N n 0 = A -N r +

a) by group

b) by period

Task No. 2

1) the electronic formula of an element’s atom, according to the number of electrons at the outer level, metallic and non-metallic in nature (if there are 1-3 electrons at the outer level, then the element is a metal, if there are more than 3, then the element is a non-metal;

2) electronic structural formula of the valence shell of an element’s atom, normal and excited states of the atom, negative and positive oxidation states for p - elements (non-metals), highest and lowest positive oxidation states for metals ( s - And d - families);

3) formula of a hydrogen compound (for s -element hydride with N - , For p - element gaseous hydrogen compound with H + ), name;

4) name the formulas of oxides in which positive oxidation states appear, indicate the nature;

5) name the formulas of bases and acids corresponding to the oxides; Formulas of salts, name.

Characteristic p - element S - sulfur, located in III period of the main subgroup VI groups

1) 16 S 1 s 2 2 s 2 2 p 6 3 s 2 3 p 4 - a non-metal, since at the outer level the atom has more than three electrons - six

2) S 3 s 2 3 p 4 p - element

↓ the normal state of an atom is 2 unpaired electrons, therefore,Ssulfur

S ↓ 3p 4 exhibits a negative oxidation state (-2):

3 s 2 S 0 + 2 ē →S -2

S * the first excited state is 4 unpaired electrons, therefore,S

3 d 1 exhibits a positive oxidation state (+4):

↓ 3 p 3 S 0 - 4 ē →S +4

3 s 2

the second excited state is 6 unpaired electrons, therefore,

3 d 2 sulfur exhibits a positive oxidation state (+6):

S ** 3 p 3 S 0 - 6 ē →S +6

3 s 1

3) S -2 → H 2 S- hydrogen sulfide, an aqueous solution of which is hydrosulfide acid.

SaltsH 2 Scalled sulfides; (name) K 2 S- potassium sulfide.

4) S +4 → SO 2 (sulfur oxideIV) → acidH 2 SO 3 → salt:

TO 2 SO 3 and KNSO 3

5) S +6 → SO 3 (sulfur oxideVI) → acidH 2 SO 4 → salt: K 2 SO 4 and KNSO 4

Characteristic s - element Ca - calcium, located in the fourth period of the main subgroup of the second group

1) 20 Ca 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 Kcalcium is a metal, since at the outer level the atom has less than three electrons - 2 electrons

2) Ca 4 s 2 s- element; Ca 4s 2 - normal state of the atom - no unpaired electrons

Ca * the excited state of the atom is two unpaired electrons, therefore,

Ca 0 - 2 ē → Sa +2

4p 1 Ca - exhibits a positive oxidation state (+2);negative degree

4 s 1 Metals do not oxidize

3) Ca +2 N 2 - - hydrogen connection; SaN 2 (calcium hydride)

4) Ca +2 → CaO oxide → Ca(OH) base 2 → salts: 1) CaC.I. 2 and CaOHC.I. 2) CaSO 3 AndCa(HSO 3 ) 2

Task No. 3

element

Element

Valence

shell

Lowest oxidation state

Hydrogen connection

Highest oxidation state

Supreme Oxide Formula

Hydroxide formula

Salt formula

s-element

r- element

Conclusion:

Practical work

Option 1

Drawing up electronic formulas of atoms of elements and graphic diagrams, filling them with electrons

Work progress

Task No. 1

Fill out the table: