Periodic table Dmitri Ivanovich Mendeleev is very convenient and versatile in its use. It can be used to determine some of the characteristics of the elements, and most surprisingly, to predict some of the properties of yet undiscovered, not discovered by scientists, chemical elements(for example, we know some of the properties of the proposed unbihexium, although it has not yet been discovered and synthesized).
What are metallic and non-metallic properties
These properties depend on the element's ability donate or attract electrons. It is important to remember one rule, metals - give electrons, and non-metals - accept. Accordingly, metallic properties are the ability of a certain chemical element to donate its electrons (from an external electron cloud) to another chemical element. For non-metals, the opposite is true. The easier a non-metal accepts electrons, the higher its non-metal properties.
Metals will never accept electrons from another chemical element. This is typical for the following elements;
- sodium;
- potassium;
- lithium;
- France and so on.
The situation is similar with non-metals. Fluorine exhibits its properties more than all other non-metals, it can only attract particles of another element to itself, but under no circumstances will it give up its own. It has the most non-metallic properties. Oxygen (according to its characteristics) comes immediately after fluorine. Oxygen can form a compound with fluorine, giving up its electrons, but it takes negative particles from other elements.
List of non-metals with the most pronounced characteristics:
- fluorine;
- oxygen;
- nitrogen;
- chlorine;
- bromine.
Non-metallic and metallic properties are explained by the fact that all chemical substances tend to complete their energy level. To do this, the last electron level must have 8 electrons. The fluorine atom has 7 electrons on the last electron shell, trying to complete it, it attracts one more electron. The sodium atom has one electron on the outer shell, in order to get 8, it is easier for it to give 1, and at the last level there will be 8 negatively charged particles.
Noble gases do not interact with other substances precisely because they have a completed energy level, they do not need to either attract or give away electrons.
How metallic properties change in the periodic system
The periodic table of Mendeleev consists of groups and periods. The periods are arranged horizontally in such a way that the first period includes: lithium, beryllium, boron, carbon, nitrogen, oxygen, and so on. Chemical elements are arranged strictly in order of increasing serial number.
The groups are arranged vertically in such a way that the first group includes: lithium, sodium, potassium, copper, rubidium, silver, and so on. The group number indicates the number of negative particles at the outer level of a certain chemical element. While the period number indicates the number of electron clouds.
The metallic properties are enhanced in the series from right to left or, in other words, weaken in the period. That is, magnesium has greater metallic properties than aluminum, but less than sodium. This is because in the period the number of electrons in the outer shell increases, therefore, it is more difficult for a chemical element to give up its electrons.
In the group, the opposite is true, the metallic properties are enhanced in a row from top to bottom. For example, potassium is stronger than copper, but weaker than sodium. The explanation for this is very simple, in the group the number of electron shells, and the farther the electron is from the nucleus, the easier element give it away. The force of attraction between the nucleus of an atom and an electron in the first shell is greater than between the nucleus and an electron in the 4th shell.
Let's compare two elements - calcium and barium. Barium ranks lower in the periodic table than calcium. And this means that the electrons from the outer shell of calcium are located closer to the nucleus, therefore, they are better attracted than those of barium.
It is more difficult to compare elements that are in different groups and periods. Take, for example, calcium and rubidium. Rubidium will give off negative particles better than calcium. Because it is below and to the left. But using only the periodic table, it is impossible to unequivocally answer this question by comparing magnesium and scandium (since one element is lower and to the right, and the other is higher and to the left). To compare these elements, special tables will be needed (for example, the electrochemical series of metal voltages).
How non-metallic properties change in the periodic system
Non-metallic properties in the periodic system of Mendeleev change exactly the opposite than metal ones. In fact, these two signs are antagonists.
Strengthen in the period (in a row from right to left). For example, sulfur is less able to attract electrons than chlorine, but more than phosphorus. The explanation for this phenomenon is the same. The number of negatively charged particles on the outer layer increases and therefore it is easier for the element to complete its energy level.
Non-metallic properties decrease in a row from top to bottom (in a group). For example, phosphorus is able to give off negatively charged particles more than nitrogen, but at the same time it is able to attract better than arsenic. Phosphorus particles are attracted to the core better than arsenic particles, which gives it the advantage of an oxidizing agent in reactions to lower and increase the oxidation state (redox reactions).
Compare, for example, sulfur and arsenic. Sulfur is higher and to the right, which means that it is easier for her to complete her energy level. Like metals, non-metals are difficult to compare if they are in different groups and periods. For example, chlorine and oxygen. One of these elements is above and to the left, and the other is below and to the right. To answer, we will have to refer to the table of electronegativity of non-metals, from which we see that oxygen more easily attracts negative particles to itself than chlorine.
Periodic table of Mendeleev helps to find out not only the number of protons in an atom, atomic mass and ordinal, but also helps to determine the properties of the elements.
Video
The video will help you understand the regularities of the properties of chemical elements and their compounds by periods and groups.
Lecture: Patterns of changes in the properties of elements and their compounds by periods and groups
Law D.I. Mendeleev
The Russian scientist D. I. Mendeleev worked successfully in many fields of science. However, he was most famous for the unique discovery of the periodic law of chemical elements in 1869. Initially, it sounded like this: “The properties of all elements, and as a result of the quality of the simple ones formed by them, as well as complex substances, stand in a periodic relationship with their atomic weight.
Currently, the wording of the law is different. The fact is that at the time of the discovery of the law, scientists had no idea about the structure of the atom, and the weight of a chemical element was taken as the atomic weight. Subsequently, an active study of the atom and obtaining new information about its structure, a law was derived that is relevant today: "Properties of atoms chem. elements and the simple substances formed by them in a periodic dependence on the charges of the nuclei of their atoms.
The law is also expressed graphically. The table shows it clearly:
Periodic table of D.I. Mendeleev
In this lesson, we will learn how to extract from it important and necessary information for understanding science. In it you see lines. This periods. There are seven in total. Recall from the previous lesson that the number of each period indicates the number of energy levels in which the electrons of an atom of a chemical element are located. For example, sodium (Na) and magnesium (Mg) are in the third period, which means that their electrons are located on three energy levels. All periods, with the exception of the 1st, originate from alkali metal, and culminate in a noble gas.
Electronic configuration:
alkali metal - ns 1,
noble gas - ns 2 p 6, with the exception of helium (He) - 1s2.
Where n - is the period number.
We also see vertical columns in the table - these are groups. In some tables you can see 18 groups numbered with Arabic numerals. This form of the table is called long, it appeared after the discovery of differences between d-elements and s- and p-elements. But the traditional one created by Mendeleev is short form, where the elements are grouped into 8 groups, numbered with Roman numerals:
In the future, we will use the short table that is already familiar and familiar to you.So what information do the group numbers give us? From the number, we find out the number of electrons that form chemical bonds. They're called valence. 8 groups are divided into two subgroups: main and side.
- What other information can we extract from the table? You can see that each element is assigned a serial number. It's not accidental either. Judging by the number of an element, we can judge the number of electrons in an atom of a given element. For example, calcium (Ca) is number 20, which means there are 20 electrons in its atom.
The electrons of the s- and p-sublevels enter into the main one. These are subgroups IA, IIA, IIIA, IVA, VA, VIA, VIIA and VIIIA. For example, aluminum (Al) - an element of the main subgroup of group III has ... 3s 2 3p 1 valence electrons.
Elements located in side subgroups contain electrons of the d - sublevel. Side are groups IB, IIB, IIIB, IVB, VB, VIB, VIIB and VIIIB. For example, manganese (Mn), an element of the main subgroup of group VII, has …3d 5 4s 2 valence electrons.
In the short table, s-elements are in red, p-elements in yellow, d-elements in blue, and f-elements in white.
The following conclusion can also be drawn from the table: the higher the element's serial number, the smaller the radius of the atom. Why? The fact is that with an increase in the total number of electrons, there is a decrease in the radius of the atom. The more electrons, the higher the energy of their binding with the nucleus. For example, the nucleus of a phosphorus (P) atom holds the electrons of its outer level much more strongly than the nucleus of a sodium (Na) atom, which has one electron in the outer level. And if the atoms of phosphorus and sodium react, phosphorus will take this electron from sodium, because phosphorus is more electronegative. This process is called electronegativity. Remember, when moving to the right along one row of elements of the table, their electronegativity increases, and within one subgroup it decreases. We will talk about this property of elements in more detail in the next lessons.
Remember:
1. In periods with an increase in the serial number, we can observe:- increase in nuclear charge and decrease in atomic radius;
- increase in the number of external electrons;
- increase in ionization and electronegativity;
- an increase in non-metallic oxidizing properties and a decrease in metal reducing properties;
- an increase in acidity and a decrease in the basicity of hydroxides and oxides.
- increase in nuclear charge and increase in atomic radius;
- decrease in ionization and electronegativity;
- a decrease in non-metallic oxidizing properties and an increase in metal reducing properties;
- an increase in basicity and a decrease in acidity of hydroxides and oxides.
Ionization is the process of converting atoms into ions (positively charged cations or negatively charged anions) during a chemical reaction.
Electronegativity is the ability of an atom to attracting an electron from another atom during a chemical reaction.
Oxidation- the process of transferring an electron from a reducing agent atom (electron donor) to an oxidizing atom (electron acceptor) and increasing the degree of oxidation of a substance atom.
There are three values for the degree of oxidation:
- with a high electronegativity of an element, it attracts electrons to itself more strongly and its atoms acquire a negative oxidation state (for example, fluorine always has an oxidation state of - 1);
- at low electronegativity, the element gives up electrons and acquires a positive oxidation state (all metals have a + degree, for example, potassium +1, calcium +2, aluminum +3);
- atoms of simple substances consisting of one element have atoms with high and free atoms have a zero degree.
(Z) is periodic. Within the same period with increasing Z there is a tendency to a decrease in the size of atoms. For example, in the second period, atomic radii have the following values:
|
r , nm |
0,155 |
0,113 |
0,091 |
0,077 |
0,071 |
0,066 |
0,064 |
This is explained by an increase in the attraction of the electrons of the outer layer to the nucleus as the charge of the nucleus increases. In subgroups, from top to bottom, atomic radii increase, because the number of electron layers increases:
|
r , nm |
r , nm |
||
|
0,155 |
0,071 |
||
|
0,189 |
0,130 |
||
|
0,236 |
0,148 |
||
|
0,248 |
0,161 |
||
|
0,268 |
0,182 |
The loss of electrons by an atom leads to a decrease in its effective size, and the addition of excess electrons leads to an increase. Therefore, the radius of a positive ion (cation) is always less, and the radius of a negative ion (anion) is always greater than the radius of the corresponding electrically neutral atom. For example:
|
r , nm |
r , nm |
||
|
0,236 |
Cl 0 |
0,099 |
|
|
0,133 |
Cl - |
0,181 |
The radius of the ion differs the more from the radius of the atom, the greater the charge of the ion:
|
cr 0 |
Cr2+ |
Cr3+ |
|
|
r , nm |
0,127 |
0,083 |
0,064 |
Within one subgroup, the radii of ions of the same charge increase with increasing nuclear charge:
|
r , nm |
r , nm |
||
|
Li + |
0,068 |
0,133 |
|
|
Na+ |
0,098 |
Cl - |
0,181 |
|
0,133 |
Br - |
0,196 |
|
|
Rb+ |
0,149 |
0,220 |
This regularity is explained by the increase in the number of electron layers and the growing distance of the outer electrons from the nucleus.
b) Ionization energy and electron affinity. In chemical reactions, the nuclei of atoms do not undergo changes, while the electron shell is rebuilt, and the atoms are able to turn into positively and negatively charged ions. This ability can be quantified by the ionization energy of an atom and its electron affinity.
Ionization energy (ionization potential) I is the amount of energy required to detach an electron from an unexcited atom to form a cation:
X- e→ X +
Energy ionization is measured in kJ/mol or in electronvolts 1 eV = 1.602. 10 -19 J or 96.485 kJ / mol.(eV). The detachment of the second electron is more difficult than the first, because the second electron is detached not from a neutral atom, but from a positive ion:
X+- e→ X 2+
Therefore, the second ionization potential I 2 more than the first ( I 2 >I one). Obviously, the removal of each next electron will require more energy than the removal of the previous one. To characterize the properties of elements, the energy of detachment of the first electron is usually taken into account.
In groups, the ionization potential decreases with increasing atomic number of the element:
|
I, eV |
6,39 |
5,14 |
4,34 |
4,18 |
3,89 |
This is due to the greater distance of valence electrons from the nucleus and, consequently, their easier detachment as the number of electron layers increases. The value of the ionization potential can serve as a measure of the “metallicity” of an element: the lower the ionization potential, the easier it is to remove an electron from an atom, the more pronounced the metallic properties.
In periods from left to right, the charge of the nucleus increases, and the radius of the atom decreases. Therefore, the ionization potential gradually increases, and the metallic properties weaken:
|
I, eV |
5,39 |
9,32 |
8,30 |
11,26 |
14,53 |
13,61 |
17,42 |
21,56 |
Breaking the upward trend I observed for atoms with a completely filled external energy sublevel, or for atoms in which the external energy sublevel is exactly half filled:
This indicates an increased energy stability of electronic configurations with completely or exactly half-occupied sublevels.
The degree of attraction of an electron to the nucleus and, consequently, the ionization potential depends on a number of factors, and above all on nuclear charge The charge of the nucleus is equal to the ordinal number of the element in the periodic table., on the distance between the electron and the nucleus, on the screening effect of other electrons. So, for all atoms, except for the elements of the first period, the influence of the nucleus on the electrons of the outer layer is screened by the electrons of the inner layers.
The field of the nucleus of an atom, which holds the electrons, also attracts a free electron if it is near the atom. True, this electron experiences repulsion from the electrons of the atom. For many atoms, the energy of attraction of an additional electron to the nucleus exceeds the energy of its repulsion from the electron shells. These atoms can add an electron, forming a stable singly charged anion. The energy of detachment of an electron from a negative singly charged ion in the process X - - e → X 0 is called the affinity of an atom for an electron ( A), measured in kJ/mol or eV. When two or more electrons are attached to an atom, repulsion prevails over attraction - the affinity of an atom for two or more electrons is always negative. Therefore, monatomic multiply charged negative ions (O 2-, S 2-, N 3-, etc.) cannot exist in the free state.
Electron affinity is not known for all atoms. Halogen atoms have the highest electron affinity.
B) electronegativity. This value characterizes the ability of an atom in a molecule to attract binding electrons to itself. Electronegativity should not be confused with electron affinity: the first concept refers to an atom in a molecule, and the second to an isolated atom. Absolute electronegativity(kJ/mol or eV 1 electronvolt = 1.602. 10 -19 J or 96.485 kJ / mol.) is equal to the sum of the ionization energy and electron affinity :AEO= I+A. In practice, the relative value is often used electronegativity, equal to the ratio of the AEO of this element to the AEO of lithium (535 kJ/mol):
A.I. Khlebnikov, I.N. Arzhanova, O.A. Napilkova
So, when oxides react with water, depending on the properties of the main element, either a base or hydroxide, or an acid is obtained.
For example:
SO3+H2O=H2SO4 - manifestation;
CaO+H2O=Ca(OH)2 - manifestation of basic properties;
Periodic table of Mendeleev, as an indicator of acid-base properties
The periodic table can help in determining the acid-base properties of the elements. If you look at the periodic table, you can see such a pattern that non-metallic or acid properties. Accordingly, metals are closer to the left edge, amphoteric elements are in the center, and non-metals are to the right. If you look at the electrons and their attraction to the nucleus, it is noticeable that on the left side the elements have a weak nucleus charge, and the electrons are at the s-level. As a result, it is easier for such elements to donate an electron than for elements located on the right side. Non-metals have a fairly high nuclear charge. This complicates the return of free electrons. It is easier for such elements to attach electrons to themselves, showing acidic properties.Three theories for defining properties
There are three approaches that define a compound: the Bronsted-Lowry proton theory, the Lewis aproton electron theory, and the Arrhenius theory.According to the proton theory, compounds that can donate their protons have acidic properties. Such compounds were called donors. And the main properties are manifested by the ability to accept or attach a proton.
The aprotic approach implies that the acceptance and donation of protons to determine the acid-base properties is not necessary. Acid properties, according to this theory, are manifested by the ability to accept an electron pair, and basic, on the contrary, give this pair.
The Arrhenius theory is the most relevant for determining acid-base properties. In the course of the study, it was proved that acidic properties appear when, during dissociation aqueous solutions chemical compound is divided into anions and hydrogen ions, and the main properties are divided into cations and hydroxide ions.
A strong base is an inorganic chemical compound formed by a hydroxyl group -OH and an alkaline (elements of group I periodic system: Li, K, Na, RB, Cs) or alkaline earth metal (group II elements Ba, Ca). They are written as formulas LiOH, KOH, NaOH, RbOH, CsOH, Ca(OH) ₂, Ba(OH) ₂.
You will need
- evaporating cup
- burner
- indicators
- metal rod
- H₃RO₄
Instruction
Strong bases exhibit, characteristic of all. The presence in the solution is determined by the change in color of the indicator. Add phenolphthalein to the sample with the test solution or omit litmus paper. Methyl orange gives a yellow color, phenolphthalein gives a purple color, and litmus paper turns a blue color. The stronger the base, the more intense the color of the indicator.
If you need to find out which alkalis are presented to you, then spend qualitative analysis solutions. The most common strong bases are lithium, potassium, sodium, barium, and calcium. Bases react with acids (neutralization reactions) to form salt and water. In this case, Ca(OH) ₂, Ba(OH) ₂ and LiOH can be distinguished. When interacting with orthophosphoric acid, insoluble precipitates are formed. The remaining hydroxides will not give precipitation, tk. all K and Na salts are soluble.
3 Ca(OH) ₂ + 2 H₃RO₄ --→ Ca₃(PO₄)₂↓+ 6 H₂O
3 Va(OH) ₂ +2 H₃RO₄ --→ Va₃(PO₄)₂↓+ 6 H₂О
3 LiOH + Н₃РО₄ --→ Li₃РО₄↓ + 3 H₂О
Strain them and dry them. Inject the dried sediments into the flame of the burner. Lithium, calcium and barium ions can be qualitatively determined by changing the color of the flame. Accordingly, you will determine where which hydroxide is. Lithium salts color the burner flame carmine red. Barium salts - in green, and calcium salts - in raspberry.
The remaining alkalis form soluble orthophosphates.
3 NaOH + Н₃РО₄--→ Na₃РО₄ + 3 H₂О
3 KOH + H₃PO₄--→ K₃PO₄ + 3 H₂О
Evaporate the water to a dry residue. Evaporated salts on a metal rod alternately bring into the burner flame. Where sodium salt is located, the flame will turn bright yellow, and potassium orthophosphate will turn pink-violet. Thus, having a minimum set of equipment and reagents, you have determined all the strong reasons given to you.
a) Patterns associated with the metallic and non-metallic properties of elements.
1. When moving along the period FROM RIGHT TO LEFT metallic element properties GAIN. In the opposite direction, non-metallic ones increase.
From left to right in the period, the charge of the nucleus also increases. Consequently, the attraction to the nucleus of valence electrons increases and their return becomes more difficult.
2. Moving TOP DOWN along groups REINFORCED METAL element properties. This is due to the fact that below in the groups there are elements that already have quite a lot of filled electron shells. Their outer shells are further from the core.
b) Patterns associated with redox properties. Changes in the electronegativity of elements.
1. FROM LEFT TO RIGHT OXIDATIVE properties, and when moving TOP DOWN - RECOVERY element properties.
2. ELECTRICITY INCREASES too FROM LEFT TO RIGHT, reaching a maximum for halogens.
3. Moving TOP DOWN by groups ELECTRICITY DECREASES. This is due to an increase in the number of electron shells, on the last of which the electrons are attracted to the nucleus more and more weakly.
c) Regularities related to the size of atoms.
1. Dimensions of atoms (ATOMIC RADIUS) when moving FROM LEFT TO RIGHT along the period DECREASE.
2. Moving FROM TOP DOWN ATOMIC RADIUS elements GROW, because more electron shells are filled.
Question 3.
The structure of matter. hybridization of orbitals. Types of chemical bonds. Ionization potential and electronegativity.
The structure of matter
All bodies are made up of individual particles - molecules and atoms. Molecules are the smallest particles of matter. Molecules are made up of atoms.
Basic information about the composition of the substance:
1) All bodies consist of individual particles (molecules and atoms), between which there are gaps.
2) Molecules are constantly and randomly moving.
3) Molecules interact with each other (attract and repel).
Molecule properties:
1) Molecules of the same substance are the same.
2) When heated, the gaps between molecules increase, and when cooled, they decrease.
3) With increasing temperature, the speed of movement of molecules increases.
According to the type of structure, all substances are divided into molecular And non-molecular. Among organic matter molecular substances predominate, among inorganic - non-molecular.
According to the type of chemical bond, substances are divided into substances with covalent bonds, substances with ionic bonds(ionic substances) and substances with metallic bonds (metals).
Substances with covalent bonds can be molecular or non-molecular. This significantly affects their physical properties.
molecular substances consist of molecules interconnected by weak intermolecular bonds, these include: H 2, O 2, N 2, Cl 2, Br 2, S 8, P 4 and others simple substances; CO 2 , SO 2 , N 2 O 5 , H 2 O, HCl, HF, NH 3 , CH 4 , C 2 H 5 OH, organic polymers and many other substances. These substances do not have high strength, have low temperatures melting and boiling, do not conduct electricity, some of them are soluble in water or other solvents.
Non-molecular substances with covalent bonds or atomic substances (diamond, graphite, Si, SiO 2 , SiC and others) form very strong crystals (the exception is layered graphite), they are insoluble in water and other solvents, have high melting and boiling points, most of them do not conducts electric current (except for graphite, which has electrical conductivity, and semiconductors - silicon, germanium, etc.)
All ionic substances are naturally non-molecular. These are solid refractory substances whose solutions and melts conduct electric current. Many of them are soluble in water.
Hybridization of orbitals
Hybridization of orbitals- this is a change in the shape of some orbitals during the formation covalent bond to achieve more efficient orbital overlap.
sp 3 - Hybridization. One s-orbital and three p-orbitals turn into four identical "hybrid" orbitals, the angle between the axes of which is 109 ° 28 ". Molecules in which sp 3 hybridization occurs have a tetrahedral geometry (CH 4, NH 3).
sp 2 - Hybridization. One s-orbital and two p-orbitals turn into three identical "hybrid" orbitals, the angle between the axes of which is 120°.
Molecules in which sp 2 hybridization is carried out have a flat geometry.
sp- hybridization. One s-orbital and one p-orbital are transformed into two identical "hybrid" orbitals, the angle between the axes of which is 180°. Molecules in which sp-hybridization occurs have a linear geometry.
Types of chemical bonds.
1) Ionic(metal + non-metal)
2) covalent(non-metal + non-metal using shared electron pairs)
Types: * polar (various non-metals)
* non-polar (identical non-metals)
Species: * formed by the exchange mechanism
* formed by a donor-acceptor mechanism
exchange mechanism- one-electron atomic orbitals participate in the formation of a bond, i.e. Each of the atoms provides for the general use of one electron:
Donor-acceptor mechanism(coordination link) - chemical bond between two atoms or a group of atoms, carried out due to the lone pair of electrons of one atom (donor) and the free orbital of another atom (acceptor).
3) metal(between metal atoms, between metal ions and common free electrons)
4) Hydrogen(between the hydrogen of one molecule and another more electronegative element (O, S, N, F) and with another molecule)
Ionization potential of an atom- minimum potential difference U, which an electron must pass in an accelerating electric field in order to acquire the kinetic energy sufficient to ionize the atom.
Electronegativity (EO) is the relative ability of atoms to attract electrons when they bond with other atoms. Electronegativity characterizes the ability of an atom to polarize chemical bonds.
