Drawing up chemical formulas for compounds of two chemical elements in cases where each element has only one stoichiometric valence.

Action algorithm	Drawing up the chemical formula of aluminum oxide
Establishing (based on the name of the compound) chemical symbols of elements
Determination of the valence of atoms of elements
Specifying the numerical ratio of atoms in a compound
Drawing up a formula	Al 2 ABOUT 3

Drawing up chemical formulas for compounds that exist in aqueous solution in the form of ions.

Action algorithm	Compiling the chemical formula of aluminum sulfate
Establishment (based on the name of the compound) of the chemical formulas of ions
Determination of the number of ion charges
Calculating the least common multiple
Determining Additional Multipliers
Specifying the numerical ratio of ions
Specifying stoichiometric indices
Drawing up a formula	Al 2 (SO 4 ) 3

Writing chemical formulas

The following rules exist for indicating stoichiometric indices and charges of ions in chemical formulas.

1. If the stoichiometric index refers to a group of atoms, the chemical symbols denoting this group are placed in brackets:

C 3 H 5 (OH) 3 – the glycerol molecule contains 3 hydroxy groups;

Ca(NO 3) 2 – the formula unit of calcium nitrate contains calcium ions and nitrate ions in a ratio of 1: 2.

2. Data on the charge of a complex polyatomic ion in the chemical formula refer to the entire ion:

SO 4 2– – sulfate ion – has a double negative charge;

NH 4 + – ammonium ion – has a single positive charge.

3. The chemical formula of a complex ion is placed in square brackets, followed by its charge; it consists of:

– chemical symbol of the central atom;

– chemical formula of the ligand in parentheses;

– subscript indicating the number of ligands.

4– – hexacyanoferrate(II) ion; in an ion with four negative charges, six ligands CN – (cyanide ion) are bound to the central Fe II atom (iron cation Fe 2+).

2+ – tetraammine copper (II) ion; in an ion with two positive charges, four NH 3 ligands (ammonia molecule) are bound to the central copper atom (Cu 2+ ion).

4. The chemical formula of water in hydrates and crystalline hydrates is separated by a dot from the chemical formula of the main substance.

CuSO 4 · 5H 2 O – copper (II) sulfate pentahydrate (copper sulfate).

Classification of inorganic substances and their properties

All inorganic substances are divided into simple and complex.

Simple substances are divided into metals, non-metals and inert gases.

The most important classes of complex inorganic substances are: oxides, bases, acids, amphoteric hydroxides, salts.

Oxides - these are compounds of two elements, one of which is oxygen. General formula of oxides:

E m O n

where m – number of atoms of element E;

n is the number of oxygen atoms.

Examples of oxides: K 2 O, CaO, SO 2, P 2 O 5

Grounds - these are complex substances whose molecules consist of a metal atom and one or more hydroxide groups - OH. General formula of bases:

Me(HE) y

where y – the number of hydroxide groups equal to the valence of the metal (Me).

Examples of bases: NaOH, Ca(OH) 2, Co(OH) 3

Acids - these are complex substances containing hydrogen atoms, which can be replaced by metal atoms.

General formula of acids

N X Ace at

where Ac is an acidic residue (from English, acid – acid);

X – the number of hydrogen atoms equal to the valence of the acid residue.

Examples of acids: HC1, HNO 3, H 2 SO 4, H 3 PO 4

Amphoteric hydroxides - These are complex substances that have the properties of acids and the properties of bases. Therefore, the formulas of amphoteric hydroxides can be written in the form of bases and in the form of acids. Examples of amphoteric hydroxides:

Zn(OH) 2 = H 2 ZnO 2

Al(OH) 3 = H 3 AlO 3

form form

acid bases

Salts - these are complex substances that are the products of the replacement of hydrogen atoms in acid molecules with metal atoms or the products of the replacement of hydroxide groups in base molecules with acidic residues. For example:

The composition of normal salts is expressed by the general formula:

Meh X (Ac) at

where x - number of metal atoms; at - number of acid residues.

Examples of salts: K 3 PO 4 ; Mg SO 4; Al 2 (SO) 3 ; FeCl 3.

Oxides

For example: CO – carbon monoxide (II) – (read: “carbon monoxide two”); CO 2 – carbon monoxide (IV); Fe 2 O 3 – iron (III) oxide.

If an element has a constant valency, it is not indicated in the name of the oxide. For example: Na 2 O – sodium oxide; Al 2 O 3 – aluminum oxide.

Classification

All oxides are divided into salt-forming and non-salt-forming (or indifferent).

Non-salt-forming (indifferent) oxides- these are oxides that do not form salts when interacting with acids and bases. There are not many of them. Remember the four non-salt-forming oxides: CO, SiO, N 2 O, NO.

Salt-forming oxides are oxides that form salts when reacting with acids or bases. For example:

Na 2 O + 2HC1 = 2NaCl + H 2 O

oxide acid salt

Some oxides do not interact with water, but they correspond to hydroxides, which can be obtained indirectly. Depending on the nature of the corresponding hydroxides, all salt-forming oxides are divided into three types: basic, acidic, amphoteric.

Basic oxides are oxides whose hydrates are bases. For example:

Basic oxides		Grounds

All basic oxides are metal oxides.

Acidic oxides are oxides whose hydrates are acids. For example:

Acidic oxides

Most acid oxides are non-metal oxides. Acidic oxides are also oxides of certain metals with high valence. For example: ,

Amphoteric oxides are oxides that correspond to amphoteric hydroxides.

All amphoteric oxides are metal oxides.

Hence, nonmetals form only acid oxides; metals form all basic, All amphoteric and some acidic oxides

All oxides monovalent metals(Na 2 O, K 2 O, Cu 2 O, etc.) are the main ones. Most oxides divalent metals(CaO, BaO, FeO, etc.) are also basic. Exceptions: BeO, ZnO, PbO, SnO, which are amphoteric. Most oxides three- And tetravalent metals are amphoteric: ,,,, etc. Metal oxides with valence V, VI, VII .are acidic: ,,etc.

Metals with variable valence can form oxides of all three types.

For example: CrO is a basic oxide, Cr 2 O 3 is an amphoteric oxide, CrO 3 is an acidic oxide.

Graphic formulas

In the oxide molecule, the metal atom is directly combined with oxygen atoms.

A chemical formula reflects the composition of a substance. For example, H 2 O - two hydrogen atoms are connected to an oxygen atom. Chemical formulas also contain some information about the structure of the substance: for example, Fe(OH) 3, Al 2 (SO 4) 3 - these formulas indicate some stable groups (OH, SO 4) that are part of the substance - its molecule or formula units.

Molecular formula indicates the number of atoms of each element in a molecule. A molecular formula describes substances with a molecular structure (gases, liquids and some solids). The composition of a substance with an atomic or ionic structure can only be described by formula units.

Formula unit indicates the simplest relationship between the number of atoms of different elements in a substance. For example, the formula unit of benzene is CH, the molecular formula is C6H6.

Structural (graphic) formula indicates the order of connection of atoms in the molecule and in the formula unit and the number of bonds between atoms.

Valence

The correct writing of such formulas is based on the idea of valency(valentia - strength) as the ability of an atom of a given element to attach to itself a certain number of other atoms. In modern chemistry, three types of valency are considered: stoichiometric, electronic and structural.

Stoichiometric valence chemical element - this is the number of equivalents that a given atom can attach to itself, or the number of equivalents in an atom. Equivalents are determined by the number of attached or substituted hydrogen atoms, so the stoichiometric valency is equal to the number of hydrogen atoms with which a given atom interacts. But not all elements interact with hydrogen, but almost all elements interact with oxygen, so stoichiometric valency can be defined as twice the number of attached oxygen atoms.

For example, the stoichiometric valence of sulfur in hydrogen sulfide H 2 S is 2, in SO 2 oxide – 4, in SO 3 oxide – 6.

When determining the stoichiometric valence of an element using the formula of a binary compound, one should be guided by the rule: the total valence of all atoms of one element must be equal to the total valence of all atoms of another element.

Knowing the valency of the elements and this rule, you can create the chemical formula of the compound. When composing formulas, the following procedure should be followed.

1. Write, in order of increasing electronegativity, the chemical symbols of the elements that are part of the compound, for example:

2. Above the symbols of chemical elements their valency is indicated (usually denoted by Roman numerals):

I II III I III II

3. Using the above rule, determine the least common multiple of the numbers expressing the stoichiometric valency of both elements (2, 3 and 6, respectively).

4) By dividing the least common multiple by the valency of the corresponding element, the number of atoms in the formula of the compounds is found:

I II III I III II

K 2 O AlCl 3 Al 2 O 3

Example 15. Create a formula for chlorine oxide, knowing that chlorine in it is heptavalent and oxygen is divalent.

Solution. We find the smallest multiple of the numbers 2 and 7 - it is equal to 14. Dividing the least common multiple by the stoichiometric valency of the corresponding element, we find the number of atoms: chlorine 14 : 7 = 2, oxygen 14 : 2 =7. Thus, the oxide formula is Cl 2 O 7.

Oxidation state also characterizes the composition of the substance and is equal to the stoichiometric valence with a plus sign (for a metal or a more electropositive element in the molecule) or minus.

1. In simple substances, the oxidation state of elements is zero.

2. The oxidation state of fluorine in all compounds is -1. The remaining halogens (chlorine, bromine, iodine) with metals, hydrogen and other more electropositive elements also have an oxidation state of -1, but when combined with more electronegative elements they have positive oxidation states.

3. Oxygen in compounds has an oxidation state of -2; the exceptions are hydrogen peroxide H 2 O 2 and its derivatives (Na 2 O 2, BaO 2, etc., in which oxygen has an oxidation state of -1, as well as oxygen fluoride OF 2, in which the oxidation state of oxygen is +2.

4. Alkaline elements (Li, Na, K, etc.) and elements of the main subgroup of the second group of the Periodic Table (Be, Mg, Ca, etc.) always have an oxidation state equal to the group number, that is, +1 and +2, respectively .

5. All elements of the third group, except thallium, have a constant oxidation state equal to the group number, i.e. +3.

6. The highest oxidation state of an element is equal to the group number of the Periodic Table, and the lowest is the difference: group number is 8. For example, the highest oxidation state of nitrogen (it is located in the fifth group) is +5 (in nitric acid and its salts), and the lowest is equal to -3 (in ammonia and ammonium salts).

7. The oxidation states of the elements in a compound cancel each other out so that their sum for all atoms in a molecule or neutral formula unit is zero, and for an ion it is its charge.

These rules can be used to determine the unknown oxidation state of an element in a compound if the oxidation states of the others are known, and to construct formulas for multielement compounds.

Example 16. Determine the degree of oxidation of chromium in the salt K 2 CrO 4 and in the ion Cr 2 O 7 2 - .

Solution. The oxidation state of potassium is +1 (rule 4) and oxygen is -2 (rule 3). The oxidation state of chromium is denoted by X. For the formula unit K 2 CrO 4 we have:

2∙(+1) + X + 4∙(-2) = 0,

therefore, the oxidation state of chromium is X = +6.

For the Cr 2 O 7 2 - ion we have: 2∙X + 7∙(-2) = -2, X = +6.

We see that the oxidation state of chromium is the same in both cases.

Example 17. Determine the degree of oxidation of phosphorus in the compounds P 2 O 3 and PH 3.

Solution. In the compound P 2 O 3 the oxidation state of oxygen is -2. Based on the fact that the algebraic sum of the oxidation states of a molecule must be equal to zero, we find the oxidation state of phosphorus: 2∙X + 3∙(-2) = 0, hence X = +3.

In the compound PH 3, the oxidation state of hydrogen is +1, hence X + 3∙(+1) = 0, X = -3.

Example 18. Write the formulas of the oxides that can be obtained by thermal decomposition of the following hydroxides (bases and acids): Fe(OH) 3, Cu(OH) 2, H 2 SiO 3, H 3 AsO 4, H 2 WO 4.

Solution. Fe(OH) 3 - the charge of the hydroxide ion is -1, therefore, the oxidation state of iron is +3 and the formula of the corresponding oxide is Fe 2 O 3.

Cu(OH) 2 - since there are two hydroxide ions, the total charge of which is -2, the oxidation state of copper is +2 and the oxide formula is CuO.

H 2 SiO 3 . The oxidation state of hydrogen is +1, oxygen -2, silicon – X. Algebraic equation: 2∙(+1) + X + 3∙(-2) = 0. X = +4. The oxide formula is SiO 2.

H 3 AsO 4 - the oxidation state of arsenic in acid is calculated using the equation:

3. (+1) + X + 4·(-2) = 0; X = +5.

Thus, the oxide formula is As 2 O 5.

H2WO4. The oxidation state of tungsten, calculated in the same way (check!) is +6. Therefore, the formula of the corresponding oxide is WO 3.

Chemical elements are divided into elements of constant and variable valency; accordingly, the former have a constant oxidation state in any compound, and the latter have a different state, which depends on the composition of the compound/

Let's consider how, using the Periodic System D.I. Mendeleev can determine the oxidation states of elements.

For stable oxidation states of elements main subgroups The following patterns are observed.

1. Elements of groups I-III have only oxidation states - positive, and equal in value to the group numbers, except for thallium, which has oxidation states +1 and +3.

2. For elements of groups IV-VI, in addition to the maximum positive oxidation state corresponding to the group number, and the negative one, equal to the difference between the number 8 and the group number, there are also intermediate oxidation states, usually differing by 2 units. For group IV, the oxidation states are +4, +2, -4, -2; for group V +5, +3, -3, -1; for group VI - +6, +4, -2.

3. Elements of group VII have all oxidation states from +7 to -1, differing by two units, i.e. +7, +5, +3, +1 and -1. But in this group (halogens) fluorine is released, which does not have positive oxidation states and, in combination with other elements, exists only in one oxidation state -1.

Note. There are several known unstable compounds of chlorine, bromine and iodine with even oxidation states +2, +4 and +6 (ClO, ClO 2, ClO 3, etc.).

The elements side subgroups there is no simple relationship between stable oxidation states and group number. For the most common elements of elements of secondary subgroups, stable oxidation states should simply be remembered. These elements include: chromium Cr (+3 and +6), manganese Mn (+7, +6, +4 and +2), iron Fe, cobalt Co and nickel Ni (+3 and +2), copper Cu ( +2 and +1), silver Ag (+1), gold Au (+3 and +1), zinc Zn and cadmium Cd (+2), mercury Hg (+2 and +1).

To compile formulas for three- and multi-element compounds, it is necessary to know the oxidation states of all elements. In this case, the number of atoms of elements in the formula is determined from the condition that the sum of the oxidation states of all atoms is equal to zero (in a formula unit) or charge (in an ion). For example, if it is known that the formula unit contains K, Cr and O atoms with oxidation states equal to +1, +6 and -2, respectively, then the formulas K 2 CrO 4, K 2 Cr 2 O 7, K 2 Cr 3 O 10 and many others; similarly, this ion with a charge -2 containing Cr +6 and O - 2 will correspond to the formulas CrO 4 2 -, Cr 2 O 7 2 -, Cr 3 O 10 2 -, Cr 4 O 13 2 -, etc.

Electronic valency of an element is equal to the number of chemical bonds formed by an atom of this element.

In most compounds, the electronic valency of the elements is equal to the stoichiometric one. But there are many exceptions. For example, in hydrogen peroxide H 2 O 2 the stoichiometric valency of oxygen is equal to one (for each oxygen atom there is one hydrogen atom), and the electronic valence is two, which follows from the structural formula that shows the chemical bonds of atoms: H–O–O–H . The discrepancy between the stoichiometric and electronic valence values ​​in this case is explained by the fact that oxygen atoms are bonded not only to hydrogen atoms, but also to each other.

Thus, there are chemical compounds in which the stoichiometric and electronic valencies do not coincide. These include, for example, complex compounds.

Structural (coordination) valency, or the coordination number is determined by the number of neighboring atoms. For example, in the SO3 molecule of sulfur, the number of neighboring oxygen atoms is 3 and the structural valency and coordination number are 3, while the stoichiometric valency is 6.

Electronic and coordination valencies are discussed in more detail in the chapters “Chemical Bonding” and “Complex Compounds”.

Key words: Chemistry 8th grade. All formulas and definitions, symbols of physical quantities, units of measurement, prefixes for designating units of measurement, relationships between units, chemical formulas, basic definitions, briefly, tables, diagrams.

1. Symbols, names and units of measurement
some physical quantities used in chemistry

Physical quantity	Designation	Unit
Time	t	With
Pressure	p	Pa, kPa
Quantity of substance	ν	mole
Mass of substance	m	kg, g
Mass fraction	ω	Dimensionless
Molar mass	M	kg/mol, g/mol
Molar volume	Vn	m 3 /mol, l/mol
Volume of substance	V	m 3, l
Volume fraction		Dimensionless
Relative atomic mass	A r	Dimensionless
	Mr	Dimensionless
Relative density of gas A to gas B	D B (A)	Dimensionless
Density of matter	R	kg/m 3, g/cm 3, g/ml
Avogadro's constant	N A	1/mol
Absolute temperature	T	K (Kelvin)
Temperature in Celsius	t	°C (degrees Celsius)
Thermal effect of a chemical reaction	Q	kJ/mol

2. Relationships between units of physical quantities

3. Chemical formulas in 8th grade

4. Basic definitions in 8th grade

• Atom- the smallest chemically indivisible particle of a substance.
• Chemical element- a certain type of atom.
• Molecule- the smallest particle of a substance that retains its composition and chemical properties and consists of atoms.
• Simple substances- substances whose molecules consist of atoms of the same type.
• Complex substances- substances whose molecules consist of atoms of different types.
• Qualitative composition of the substance shows which atoms of elements it consists of.
• Quantitative composition of the substance shows the number of atoms of each element in its composition.
• Chemical formula- conventional recording of the qualitative and quantitative composition of a substance using chemical symbols and indices.
• Atomic mass unit(amu) - a unit of measurement of atomic mass, equal to the mass of 1/12 of a carbon atom 12 C.
• Mole- the amount of a substance that contains a number of particles equal to the number of atoms in 0.012 kg of carbon 12 C.
• Avogadro's constant (Na = 6*10 23 mol -1) - the number of particles contained in one mole.
• Molar mass of a substance (M ) is the mass of a substance taken in an amount of 1 mole.
• Relative atomic mass element A r - the ratio of the mass of an atom of a given element m 0 to 1/12 of the mass of a carbon atom 12 C.
• Relative molecular weight substances M r - the ratio of the mass of a molecule of a given substance to 1/12 of the mass of a carbon atom 12 C. The relative molecular mass is equal to the sum of the relative atomic masses of the chemical elements forming the compound, taking into account the number of atoms of a given element.
• Mass fraction chemical element ω(X) shows what part of the relative molecular mass of substance X is accounted for by a given element.

ATOMIC-MOLECULAR TEACHING
1. There are substances with molecular and non-molecular structure.
2. There are gaps between the molecules, the sizes of which depend on the state of aggregation of the substance and temperature.
3. Molecules are in continuous motion.
4. Molecules are made up of atoms.
6. Atoms are characterized by a certain mass and size.
During physical phenomena, molecules are preserved; during chemical phenomena, as a rule, they are destroyed. Atoms rearrange during chemical phenomena, forming molecules of new substances.

LAW OF CONSTANT COMPOSITION OF MATTER
Each chemically pure substance of molecular structure, regardless of the method of preparation, has a constant qualitative and quantitative composition.

VALENCE
Valence is the property of an atom of a chemical element to attach or replace a certain number of atoms of another element.

CHEMICAL REACTION
A chemical reaction is a phenomenon as a result of which other substances are formed from one substance. Reactants are substances that enter into a chemical reaction. Reaction products are substances formed as a result of a reaction.
Signs of chemical reactions:
1. Release of heat (light).
2. Change in color.
3. Odor appears.
4. Formation of sediment.
5. Gas release.

• Chemical equation- recording a chemical reaction using chemical formulas. Shows which substances and in what quantities react and are obtained as a result of the reaction.

LAW OF CONSERVATION OF MASS
The mass of substances that entered into a chemical reaction is equal to the mass of substances formed as a result of the reaction. As a result of chemical reactions, atoms do not disappear or appear, but they are rearranged.

The most important classes of inorganic substances

Lesson summary “Chemistry 8th grade. All formulas and definitions."

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Chemistry– the science of the composition, structure, properties and transformations of substances.

Atomic-molecular science. Substances consist of chemical particles (molecules, atoms, ions), which have a complex structure and consist of elementary particles (protons, neutrons, electrons).

Atom– a neutral particle consisting of a positive nucleus and electrons.

Molecule– a stable group of atoms connected by chemical bonds.

Chemical element– a type of atoms with the same nuclear charge. Element denote

where X is the symbol of the element, Z– serial number of the element in the Periodic Table of Elements D.I. Mendeleev, A– mass number. Serial number Z equal to the charge of the atomic nucleus, the number of protons in the atomic nucleus and the number of electrons in the atom. Mass number A equal to the sum of the numbers of protons and neutrons in an atom. The number of neutrons is equal to the difference A–Z.

Isotopes– atoms of the same element having different mass numbers.

Relative atomic mass(A r) is the ratio of the average mass of an atom of an element of natural isotopic composition to 1/12 of the mass of an atom of the carbon isotope 12 C.

Relative molecular weight(M r) is the ratio of the average mass of a molecule of a substance of natural isotopic composition to 1/12 of the mass of an atom of the 12 C carbon isotope.

Atomic mass unit(a.u.m) – 1/12 of the mass of an atom of the carbon isotope 12 C. 1 a.u. m = 1.66? 10 -24 years

Mole– the amount of a substance containing as many structural units (atoms, molecules, ions) as there are atoms in 0.012 kg of the carbon isotope 12 C. Mole– the amount of a substance containing 6.02 10 23 structural units (atoms, molecules, ions).

n = N/N A, Where n– amount of substance (mol), N– number of particles, a N A– Avogadro’s constant. The amount of a substance can also be denoted by the symbol v.

Avogadro's constant N A = 6.02 10 23 particles/mol.

Molar massM(g/mol) – ratio of the mass of the substance m(d) to the amount of substance n(mol):

M = m/n, where: m = M n And n = m/M.

Molar volume of gasV M(l/mol) – gas volume ratio V(l) to the amount of substance of this gas n(mol). Under normal conditions V M = 22.4 l/mol.

Normal conditions: temperature t = 0°C, or T = 273 K, pressure p = 1 atm = 760 mm. rt. Art. = 101,325 Pa = 101.325 kPa.

V M = V/n, where: V = V Mn And n = V/V M .

The result is a general formula:

n = m/M = V/V M = N/N A .

Equivalent- a real or fictitious particle that interacts with one hydrogen atom, or replaces it, or is equivalent to it in some other way.

Molar mass equivalents M e– the ratio of the mass of a substance to the number of equivalents of this substance: M e = m/n (eq) .

In charge exchange reactions, the molar mass of substance equivalents is

with molar mass M equal to: M e = M/(n ? m).

In redox reactions, the molar mass of equivalents of a substance with molar mass M equal to: M e = M/n(e), Where n(e)– number of transferred electrons.

Law of equivalents– the masses of reactants 1 and 2 are proportional to the molar masses of their equivalents. m 1 / m 2= M E1/M E2, or m 1 /M E1 = m 2 /M E2, or n 1 = n 2, Where m 1 And m 2– masses of two substances, M E1 And M E2– molar masses of equivalents, n 1 And n 2– the number of equivalents of these substances.

For solutions, the law of equivalents can be written as follows:

c E1 V 1 = c E2 V 2, Where with E1, with E2, V 1 And V 2– molar concentrations of equivalents and volumes of solutions of these two substances.

United gas law: pV = nRT, Where p– pressure (Pa, kPa), V– volume (m 3, l), n– amount of gas substance (mol), T – temperature (K), T(K) = t(°C) + 273, R– constant, R= 8.314 J/(K? mol), with J = Pa m 3 = kPa l.

2. Atomic structure and Periodic Law

Wave-particle duality matter - the idea that every object can have both wave and corpuscular properties. Louis de Broglie proposed a formula connecting the wave and corpuscular properties of objects: ? = h/(mV), Where h– Planck’s constant, ? – wavelength that corresponds to each body with mass m and speed V. Although wave properties exist for all objects, they can be observed only for micro-objects with masses on the order of the mass of an atom and an electron.

Heisenberg Uncertainty Principle: ?(mV x) ?х > h/2n or ?V x ?x > h/(2?m), Where m– particle mass, x– its coordinate, Vx– speed in direction x, ?– uncertainty, error of determination. The uncertainty principle means that it is impossible to simultaneously indicate the position (coordinate) x) and speed (V x) particles.

Particles with small masses (atoms, nuclei, electrons, molecules) are not particles in the sense of Newtonian mechanics and cannot be studied by classical physics. They are studied by quantum physics.

Principal quantum numbern takes values ​​1, 2, 3, 4, 5, 6 and 7, corresponding to the electronic levels (layers) K, L, M, N, O, P and Q.

Level– the space where electrons with the same number are located n. Electrons of different levels are spatially and energetically separated from each other, since the number n determines electron energy E(the more n, the more E) and distance R between electrons and nucleus (the more n, the more R).

Orbital (side, azimuthal) quantum numberl takes values ​​depending on the number n:l= 0, 1,…(n- 1). For example, if n= 2, then l = 0, 1; If n= 3, then l = 0, 1, 2. Number l characterizes the sublevel (sublayer).

Sublevel– the space where electrons with certain n And l. Sublevels of a given level are designated depending on the number l:s- If l = 0, p- If l = 1, d- If l = 2, f- If l = 3. The sublevels of a given atom are designated depending on the numbers n And l, for example: 2s (n = 2, l = 0), 3d(n= 3, l = 2), etc. Sublevels of a given level have different energies (the more l, the more E): E s< E < Е А < … and the different shapes of the orbitals that make up these sublevels: the s-orbital has the shape of a ball, p-the orbital is shaped like a dumbbell, etc.

Magnetic quantum numberm 1 characterizes the orientation of the orbital magnetic moment, equal to l, in space relative to the external magnetic field and takes the following values: – l,…-1, 0, 1,…l, i.e. total (2l + 1) value. For example, if l = 2, then m 1 =-2, -1, 0, 1, 2.

Orbital(part of a sublevel) – the space where electrons (no more than two) are located with certain n, l, m 1. Sublevel contains 2l+1 orbital. For example, d– the sublevel contains five d-orbitals. Orbitals of the same sublevel having different numbers m 1, have the same energy.

Magnetic spin numberm s characterizes the orientation of the electron’s own magnetic moment s, equal to?, relative to the external magnetic field and takes two values: +? And _ ?.

Electrons in an atom occupy levels, sublevels and orbitals according to the following rules.

Pauli's rule: In one atom, two electrons cannot have four identical quantum numbers. They must differ in at least one quantum number.

From the Pauli rule it follows that an orbital can contain no more than two electrons, a sublevel can contain no more than 2(2l + 1) electrons, a level can contain no more 2n 2 electrons.

Klechkovsky's rule: electronic sublevels are filled in in order of increasing amount (n + l), and in case of the same amount (n+l)– in ascending order of number n.

Graphic form of Klechkovsky's rule.

According to Klechkovsky’s rule, sublevels are filled in in the following order: 1s, 2s, 2р, 3s, Зр, 4s, 3d, 4р, 5s, 4d, 5р, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s,…

Although the filling of sublevels occurs according to the Klechkovsky rule, in the electronic formula the sublevels are written sequentially by level: 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f etc. Thus, the electronic formula of the bromine atom is written as follows: Br(35e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 5 .

The electronic configurations of a number of atoms differ from those predicted by Klechkovsky's rule. So, for Cr and Cu:

Сr(24e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 and Cu(29e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1.

Rule of Hunda (Gunda): The filling of the orbitals of a given sublevel is carried out so that the total spin is maximum. The orbitals of a given sublevel are filled first with one electron at a time.

Electronic configurations of atoms can be written by levels, sublevels, orbitals. For example, the electronic formula P(15e) can be written:

a) by levels)2)8)5;

b) by sublevels 1s 2 2s 2 2p 6 3s 2 3p 3;

c) by orbital

Examples of electronic formulas of some atoms and ions:

V(23e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 3 4s 2;

V 3+ (20e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 0.

3. Chemical bond

3.1. Valence bond method

According to the valence bond method, a bond between atoms A and B is formed by sharing a pair of electrons.

Covalent bond. Donor-acceptor connection.

Valence characterizes the ability of atoms to form chemical bonds and is equal to the number of chemical bonds formed by an atom. According to the valence bond method, valence is equal to the number of shared pairs of electrons, and in the case of a covalent bond, valence is equal to the number of unpaired electrons in the outer level of an atom in its ground or excited states.

Valence of atoms

For example, for carbon and sulfur:

Saturability covalent bond: atoms form a limited number of bonds equal to their valency.

Hybridization of atomic orbitals– mixing of atomic orbitals (AO) of different sublevels of the atom, the electrons of which participate in the formation of equivalent?-bonds. Hybrid orbital (HO) equivalence explains the equivalence of the chemical bonds formed. For example, in the case of a tetravalent carbon atom there is one 2s– and three 2p-electron. To explain the equivalence of the four?-bonds formed by carbon in the molecules CH 4, CF 4, etc., atomic one s- and three R- orbitals are replaced by four equivalent hybrid ones sp 3-orbitals:

Focus A covalent bond is that it is formed in the direction of maximum overlap of the orbitals that form a common pair of electrons.

Depending on the type of hybridization, hybrid orbitals have a specific location in space:

sp– linear, the angle between the axes of the orbitals is 180°;

sp 2– triangular, angles between the axes of the orbitals are 120°;

sp 3– tetrahedral, angles between the axes of the orbitals are 109°;

sp 3 d 1– trigonal-bipyramidal, angles 90° and 120°;

sp 2 d 1– square, angles between the axes of the orbitals are 90°;

sp 3 d 2– octahedral, the angles between the axes of the orbitals are 90°.

3.2. Molecular orbital theory

According to the theory of molecular orbitals, a molecule consists of nuclei and electrons. In molecules, electrons are located in molecular orbitals (MO). The MOs of the outer electrons have a complex structure and are considered as a linear combination of the outer orbitals of the atoms that make up the molecule. The number of formed MOs is equal to the number of AOs involved in their formation. The energies of MOs can be lower (bonding MOs), equal (non-bonding MOs) or higher (antibonding MOs), than the energies of the AOs that form them.

Terms of interaction of JSC

1. AO interact if they have similar energies.

2. AOs interact if they overlap.

3. AO interact if they have the appropriate symmetry.

For a diatomic molecule AB (or any linear molecule), the symmetry of MO can be:

If a given MO has an axis of symmetry,

If a given MO has a plane of symmetry,

If the MO has two perpendicular planes of symmetry.

The presence of electrons on the bonding MOs stabilizes the system, as it reduces the energy of the molecule compared to the energy of the atoms. The stability of the molecule is characterized bond order n, equal to: n = (n light – n size)/2, Where n light and n size - number of electrons in bonding and antibonding orbitals.

The filling of MOs with electrons occurs according to the same rules as the filling of AOs in an atom, namely: Pauli’s rule (there cannot be more than two electrons on a MO), Hund’s rule (the total spin must be maximum), etc.

The interaction of 1s-AO atoms of the first period (H and He) leads to the formation of bonding?-MO and antibonding?*-MO:

Electronic formulas of molecules, bond orders n, experimental bond energies E and intermolecular distances R for diatomic molecules from atoms of the first period are given in the following table:

Other atoms of the second period contain, in addition to 2s-AO, also 2p x -, 2p y – and 2p z -AO, which upon interaction can form?– and?-MO. For O, F, and Ne atoms, the energies of the 2s– and 2p-AOs differ significantly, and the interaction between the 2s-AO of one atom and the 2p-AO of another atom can be neglected, considering the interaction between the 2s-AO of two atoms separately from the interaction of their 2p-AO. The MO scheme for molecules O 2, F 2, Ne 2 has the following form:

For atoms B, C, N, the energies of 2s– and 2p-AO are close in their energies, and the 2s-AO of one atom interacts with the 2p z-AO of another atom. Therefore, the order of MOs in molecules B 2, C 2 and N 2 differs from the order of MOs in molecules O 2, F 2 and Ne 2. Below is the MO scheme for molecules B 2, C 2 and N 2:

Based on the given MO schemes, it is possible, for example, to write down the electronic formulas of the molecules O 2 , O 2 + and O 2 ?:

O 2 + (11e)? s2? s *2 ? z 2 (? x 2 ? y 2)(? x *1 ? y *0)

n = 2 R = 0.121 nm;

O 2 (12e)? s2? s *2 ? z 2 (? x 2 ? y 2)(? x *1 ? y *1)

n = 2.5 R = 0.112 nm;

O 2 ?(13e)? s2? s *2 ? z 2 (? x 2 ? y 2)(? x *2 ? y *1)

n = 1.5 R = 0.126 nm.

In the case of the O 2 molecule, MO theory allows us to foresee greater strength of this molecule, since n = 2, the nature of changes in binding energies and internuclear distances in the series O 2 + – O 2 – O 2 ?, as well as the paramagnetism of the O 2 molecule, the upper MOs of which have two unpaired electrons.

3.3. Some types of connections

Ionic bond– electrostatic bond between ions of opposite charges. An ionic bond can be considered an extreme case of a polar covalent bond. An ionic bond is formed if the difference in electronegativity of the atoms? X is greater than 1.5–2.0.

An ionic bond is non-directional non-saturable communication In a NaCl crystal, the Na+ ion is attracted by all the Cl ions? and is repelled by all other Na + ions, regardless of the direction of interaction and the number of ions. This determines the greater stability of ionic crystals compared to ionic molecules.

Hydrogen bond– a bond between a hydrogen atom of one molecule and an electronegative atom (F, CI, N) of another molecule.

The existence of a hydrogen bond explains the anomalous properties of water: the boiling point of water is much higher than that of its chemical analogues: t kip (H 2 O) = 100 °C, and t kip (H 2 S) = -61 ° C. No hydrogen bonds are formed between H 2 S molecules.

4. Patterns of chemical processes

4.1. Thermochemistry

Energy(E)- ability to produce work. Mechanical work (A) is performed, for example, by gas during its expansion: A = p?V.

Reactions that occur with the absorption of energy are: endothermic.

Reactions that involve the release of energy are: exothermic.

Types of energy: heat, light, electrical, chemical, nuclear energy, etc.

Energy types: kinetic and potential.

Kinetic energy– the energy of a moving body, this is the work that a body can do before it reaches rest.

Heat (Q)– a type of kinetic energy – associated with the movement of atoms and molecules. When communicating to a body of mass (m) and specific heat capacity (c) of heat? Q its temperature increases by? t: ?Q = m with ?t, where? t = ?Q/(c t).

Potential energy- energy acquired by a body as a result of a change in position in space by it or its component parts. The energy of chemical bonds is a type of potential energy.

First law of thermodynamics: energy can pass from one type to another, but cannot disappear or arise.

Internal energy (U) – the sum of the kinetic and potential energies of the particles that make up the body. The heat absorbed in the reaction is equal to the difference in the internal energy of the reaction products and reagents (Q = ?U = U 2 – U 1), provided that the system has not done any work on the environment. If the reaction occurs at constant pressure, then the released gases do work against external pressure forces, and the heat absorbed during the reaction is equal to the sum of the changes in internal energy ?U and work A = p?V. This heat absorbed at constant pressure is called the enthalpy change: ? Н = ?U + p?V, defining enthalpy How H = U + pV. Reactions of liquid and solid substances occur without significant changes in volume (?V = 0), so what about these reactions? N close to ?U (?Н = ?U). For reactions with a change in volume we have ?Н > ?U, if expansion is in progress, and ?N< ?U , if there is compression.

The change in enthalpy is usually referred to the standard state of a substance: that is, for a pure substance in a certain state (solid, liquid or gaseous), at a pressure of 1 atm = 101,325 Pa, a temperature of 298 K and a concentration of substances of 1 mol/l.

Standard enthalpy of formation?– heat released or absorbed during the formation of 1 mole of a substance from the simple substances that constitute it, under standard conditions. For example, ?N arr.(NaCl) = -411 kJ/mol. This means that in the reaction Na(s) + ?Cl 2 (g) = NaCl(s) when 1 mole of NaCl is formed, 411 kJ of energy is released.

Standard enthalpy of reaction?H– change in enthalpy during a chemical reaction, determined by the formula: ?N = ?N arr.(products) – ?N arr.(reagents).

So for the reaction NH 3 (g) + HCl (g) = NH 4 Cl (tv), knowing? H o 6 p (NH 3) = -46 kJ/mol, ? H o 6 p (HCl) = -92 kJ /mol and?H o 6 p (NH 4 Cl) = -315 kJ/mol we have:

H = ?H o 6 p (NH 4 Cl) – ?H o 6 p (NH 3) – ?H o 6 p (HCl) = -315 – (-46) – (-92) = -177 kJ.

If? N< 0, then the reaction is exothermic. If? N> 0, then the reaction is endothermic.

Law Hess: The standard enthalpy of a reaction depends on the standard enthalpies of the reactants and products and does not depend on the path of the reaction.

Spontaneous processes can be not only exothermic, i.e. processes with a decrease in energy (?N< 0), but can also be endothermic processes, i.e. processes with increasing energy (?N> 0). In all these processes, the “disorder” of the system increases.

EntropyS – a physical quantity characterizing the degree of disorder of the system. S – standard entropy, ?S – change in standard entropy. If?S > 0, disorder increases if AS< 0, то беспорядок системы уменьшается. Для процессов в которых растет число частиц, ?S >0. For processes in which the number of particles decreases, ?S< 0. Например, энтропия меняется в ходе реакций:

CaO(solid) + H 2 O(l) = Ca(OH) 2 (solid), ?S< 0;

CaCO 3 (tv) = CaO (tv) + CO 2 (g), ?S > 0.

Processes occur spontaneously with the release of energy, i.e. for which? N< 0, and with increasing entropy, i.e. for which?S > 0. Taking both factors into account leads to the expression for Gibbs energy: G = H – TS or? G = ?H – T?S. Reactions in which the Gibbs energy decreases, i.e. ?G< 0, могут идти самопроизвольно. Реакции, в ходе которых энергия Гиббса увеличивается, т. е. ?G >0, do not go spontaneously. The condition?G = 0 means that equilibrium has been established between the products and reactants.

At low temperatures, when the value T is close to zero, only exothermic reactions occur, since T?S– little and?G = ? N< 0. At high temperatures the values T?S great, and, neglecting the size? N, we have?G = – T?S, i.e., processes with increasing entropy will occur spontaneously, for which?S > 0, a?G< 0. При этом чем больше по абсолютной величине значение?G, тем более полно проходит данный процесс.

The value of AG for a particular reaction can be determined by the formula:

G = ?С arr (products) – ?G o b p (reagents).

In this case, the values ​​of ?G o br, as well as? N arr. and?S o br for a large number of substances are given in special tables.

4.2. Chemical kinetics

Chemical reaction rate(v) is determined by the change in the molar concentration of reactants per unit time:

Where v– reaction rate, s – molar concentration of the reagent, t- time.

The rate of a chemical reaction depends on the nature of the reactants and the reaction conditions (temperature, concentration, presence of a catalyst, etc.)

Effect of concentration. IN In the case of simple reactions, the reaction rate is proportional to the product of the concentrations of the reacting substances, taken in powers equal to their stoichiometric coefficients.

For reaction

where 1 and 2 are the directions of the forward and reverse reactions, respectively:

v 1 = k 1 ? [A] m ? [B]n and

v 2 = k 2 ? [C]p ? [D]q

Where v- speed reaction, k– rate constant, [A] – molar concentration of substance A.

Molecularity of the reaction– the number of molecules participating in the elementary act of the reaction. For simple reactions, for example: mA + nB> рС + qD, molecularity is equal to the sum of the coefficients (m + n). Reactions can be single-molecule, double-molecule, and rarely triple-molecule. Reactions of higher molecular weight do not occur.

Reaction order is equal to the sum of the exponents of the degrees of concentration in the experimental expression of the rate of a chemical reaction. So, for a complex reaction

mA + nB > рС + qD the experimental expression for the reaction rate is

v 1 = k 1 ? [A] ? ? [IN] ? and the reaction order is (? + ?). Wherein? And? are found experimentally and may not coincide with m And n accordingly, since the equation of a complex reaction is the result of several simple reactions.

Effect of temperature. The rate of a reaction depends on the number of effective collisions between molecules. An increase in temperature increases the number of active molecules, giving them the necessary energy for the reaction to occur. activation energy E act and increases the rate of a chemical reaction.

Van't Hoff's rule. When the temperature increases by 10°, the reaction rate increases by 2–4 times. Mathematically this is written as:

v 2 = v 1 ? ?(t 2 – t 1)/10

where v 1 and v 2 are the reaction rates at the initial (t 1) and final (t 2) temperatures, ? – temperature coefficient of reaction rate, which shows how many times the reaction rate increases with an increase in temperature by 10°.

More precisely, the dependence of the reaction rate on temperature is expressed Arrhenius equation:

k = A? e - E/(RT)

Where k– rate constant, A– constant independent of temperature, e = 2.71828, E– activation energy, R= 8.314 J/(K? mol) – gas constant; T– temperature (K). It can be seen that the rate constant increases with increasing temperature and decreasing activation energy.

4.3. Chemical equilibrium

A system is in equilibrium if its state does not change over time. Equality of the rates of forward and reverse reactions is a condition for maintaining the equilibrium of the system.

An example of a reversible reaction is the reaction

N 2 + 3H 2 - 2NH 3 .

Law of mass action: the ratio of the product of concentrations of reaction products to the product of concentrations of starting substances (all concentrations are indicated in powers equal to their stoichiometric coefficients) is a constant called equilibrium constant.

The equilibrium constant is a measure of the progress of a forward reaction.

K = O – direct reaction does not occur;

K =? – the direct reaction goes to completion;

K > 1 – balance shifted to the right;

TO< 1 – balance is shifted to the left.

Reaction equilibrium constant TO is related to the magnitude of the change in the standard Gibbs energy?G for the same reaction:

G= – RT ln K, or?G = -2.3RT lg K, or K= 10 -0.435?G/RT

If K > 1, then lg K> 0 and?G< 0, т. е. если равновесие сдвинуто вправо, то реакция – переход от исходного состояния к равновесному – идет самопроизвольно.

If TO< 1, then lg K < 0 и?G >0, i.e. if the equilibrium is shifted to the left, then the reaction does not spontaneously go to the right.

Law of equilibrium shift: If an external influence is exerted on a system in equilibrium, a process arises in the system that counteracts the external influence.

5. Redox reactions

Redox reactions– reactions that occur with a change in the oxidation states of elements.

Oxidation– process of electron donation.

Recovery– the process of adding electrons.

Oxidizer– an atom, molecule, or ion that accepts electrons.

Reducing agent– an atom, molecule, or ion that donates electrons.

Oxidizing agents, accepting electrons, go into a reduced form:

F 2 [approx. ] + 2e > 2F? [restored].

Reductants, giving up electrons, go into the oxidized form:

Na 0 [recovery ] – 1e > Na + [approx.].

The equilibrium between the oxidized and reduced forms is characterized by Nernst equations for redox potential:

Where E 0– standard value of redox potential; n– number of transferred electrons; [restored ] and [approx. ] are the molar concentrations of the compound in reduced and oxidized forms, respectively.

Values ​​of standard electrode potentials E 0 are given in tables and characterize the oxidative and reduction properties of compounds: the more positive the value E 0, the stronger the oxidizing properties, and the more negative the value E 0, the stronger the restorative properties.

For example, for F 2 + 2e - 2F? E 0 = 2.87 volts, and for Na + + 1e - Na 0 E 0 =-2.71 volts (the process is always recorded for reduction reactions).

A redox reaction is a combination of two half-reactions, oxidation and reduction, and is characterized by an electromotive force (emf) ? E 0:?E 0= ?E 0 ok – ?E 0 restore, Where E 0 ok And? E 0 restore– standard potentials of the oxidizing agent and reducing agent for this reaction.

E.m.f. reactions? E 0 is related to the change in the Gibbs free energy?G and the equilibrium constant of the reaction TO:

?G = –nF?E 0 or? E = (RT/nF) ln K.

E.m.f. reactions at non-standard concentrations? E equal to: ? E =?E 0 – (RT/nF) ? Ig K or? E =?E 0 –(0,059/n)lg K.

In the case of equilibrium?G = 0 and?E = 0, where does it come from? E =(0.059/n)lg K And K = 10 n?E/0.059 .

For the reaction to proceed spontaneously, the following relations must be satisfied: ?G< 0 или K >> 1, to which the condition corresponds? E 0> 0. Therefore, to determine the possibility of a given redox reaction, it is necessary to calculate the value? E 0. If? E 0 > 0, the reaction is in progress. If? E 0< 0, no response.

Chemical current sources

Galvanic cells– devices that convert the energy of a chemical reaction into electrical energy.

Daniel's galvanic cell consists of zinc and copper electrodes immersed in solutions of ZnSO 4 and CuSO 4, respectively. Electrolyte solutions communicate through a porous partition. In this case, oxidation occurs on the zinc electrode: Zn > Zn 2+ + 2e, and reduction occurs on the copper electrode: Cu 2+ + 2e > Cu. In general, the reaction goes: Zn + CuSO 4 = ZnSO 4 + Cu.

Anode– electrode on which oxidation occurs. Cathode– the electrode on which the reduction takes place. In galvanic cells, the anode is negatively charged and the cathode is positively charged. On element diagrams, metal and mortar are separated by a vertical line, and two mortars are separated by a double vertical line.

So, for the reaction Zn + CuSO 4 = ZnSO 4 + Cu, the circuit diagram of the galvanic cell is written: (-)Zn | ZnSO 4 || CuSO 4 | Cu(+).

The electromotive force (emf) of the reaction is? E 0 = E 0 ok – E 0 restore= E 0(Cu 2+ /Cu) – E 0(Zn 2+ /Zn) = 0.34 – (-0.76) = 1.10 V. Due to losses, the voltage created by the element will be slightly less than? E 0. If the concentrations of solutions differ from the standard ones, equal to 1 mol/l, then E 0 ok And E 0 restore are calculated using the Nernst equation, and then the emf is calculated. corresponding galvanic cell.

Dry element consists of a zinc body, NH 4 Cl paste with starch or flour, a mixture of MnO 2 with graphite and a graphite electrode. During its operation, the following reaction occurs: Zn + 2NH 4 Cl + 2MnO 2 = Cl + 2MnOOH.

Element diagram: (-)Zn | NH4Cl | MnO 2 , C(+). E.m.f. element - 1.5 V.

Batteries. A lead battery consists of two lead plates immersed in a 30% sulfuric acid solution and coated with a layer of insoluble PbSO 4 . When charging a battery, the following processes occur on the electrodes:

PbSO 4 (tv) + 2e > Pb (tv) + SO 4 2-

PbSO 4 (tv) + 2H 2 O > PbO 2 (tv) + 4H + + SO 4 2- + 2e

When the battery is discharged, the following processes occur on the electrodes:

Pb(tv) + SO 4 2- > PbSO 4 (tv) + 2e

PbO 2 (tv) + 4H + + SO 4 2- + 2e > PbSO 4 (tv) + 2H 2 O

The total reaction can be written as:

To operate, the battery requires regular charging and monitoring of the concentration of sulfuric acid, which may decrease slightly during battery operation.

6. Solutions

6.1. Concentration of solutions

Mass fraction of substance in solution w equal to the ratio of the mass of the solute to the mass of the solution: w = m water / m solution or w = m in-va /(V ? ?), because m solution = V p-pa ? ?r-ra.

Molar concentration With equal to the ratio of the number of moles of solute to the volume of solution: c = n(mol)/ V(l) or c = m/(M? V( l )).

Molar concentration of equivalents (normal or equivalent concentration) with e is equal to the ratio of the number of equivalents of a dissolved substance to the volume of solution: with e = n(mol eq.)/ V(l) or with e = m/(M e? V(l)).

6.2. Electrolytic dissociation

Electrolytic dissociation– decomposition of the electrolyte into cations and anions under the influence of polar solvent molecules.

Degree of dissociation?– ratio of the concentration of dissociated molecules (with diss) to the total concentration of dissolved molecules (with vol): ? = with diss / with ob.

Electrolytes can be divided into strong(? ~ 1) and weak.

Strong electrolytes(for them? ~ 1) – salts and bases soluble in water, as well as some acids: HNO 3, HCl, H 2 SO 4, HI, HBr, HClO 4 and others.

Weak electrolytes(for them?<< 1) – Н 2 O, NH 4 OH, малорастворимые основания и соли и многие кислоты: HF, H 2 SO 3 , H 2 CO 3 , H 2 S, CH 3 COOH и другие.

Ionic reaction equations. IN In ionic equations of reactions, strong electrolytes are written in the form of ions, and weak electrolytes, poorly soluble substances and gases are written in the form of molecules. For example:

CaCO 3 v + 2HCl = CaCl 2 + H 2 O + CO 2 ^

CaCO 3 v + 2H + + 2Cl? = Ca 2+ + 2Cl? + H 2 O + CO 2 ^

CaCO 3 v + 2H + = Ca 2+ + H 2 O + CO 2 ^

Reactions between ions go towards the formation of a substance that produces fewer ions, i.e. towards a weaker electrolyte or a less soluble substance.

6.3. Dissociation of weak electrolytes

Let us apply the law of mass action to the equilibrium between ions and molecules in a solution of a weak electrolyte, for example acetic acid:

CH 3 COOH - CH 3 COO? +H+

The equilibrium constants for dissociation reactions are called dissociation constants. Dissociation constants characterize the dissociation of weak electrolytes: the lower the constant, the less the weak electrolyte dissociates, the weaker it is.

Polybasic acids dissociate stepwise:

H 3 PO 4 - H + + H 2 PO 4 ?

The equilibrium constant of the total dissociation reaction is equal to the product of the constants of the individual stages of dissociation:

N 3 PO 4 - ZN + + PO 4 3-

Ostwald's dilution law: the degree of dissociation of a weak electrolyte (a) increases with decreasing its concentration, i.e., with dilution:

Effect of a common ion on the dissociation of a weak electrolyte: the addition of a common ion reduces the dissociation of the weak electrolyte. So, when adding CH 3 COOH to a solution of a weak electrolyte

CH 3 COOH - CH 3 COO? +H+ ?<< 1

a strong electrolyte containing an ion common to CH 3 COOH, i.e. an acetate ion, for example CH 3 COONa

CH 3 COOna - CH 3 COO? + Na + ? = 1

the concentration of acetate ion increases, and the CH 3 COOH dissociation equilibrium shifts to the left, i.e., acid dissociation decreases.

6.4. Dissociation of strong electrolytes

Ion activity A – concentration of an ion, manifested in its properties.

Activity factorf– ion activity ratio A to concentration with: f= a/c or A = fc.

If f = 1, then the ions are free and do not interact with each other. This occurs in very dilute solutions, in solutions of weak electrolytes, etc.

If f< 1, то ионы взаимодействуют между собой. Чем меньше f, тем больше взаимодействие между ионами.

The activity coefficient depends on the ionic strength of solution I: the higher the ionic strength, the lower the activity coefficient.

Ionic strength of solution I depends on charges z and concentrations from ions:

I = 0.52?s z2.

The activity coefficient depends on the charge of the ion: the greater the charge of the ion, the lower the activity coefficient. Mathematically, the dependence of the activity coefficient f on ionic strength I and ion charge z written using the Debye-Hückel formula:

Ion activity coefficients can be determined using the following table:

6.5 Ionic product of water. pH value

Water, a weak electrolyte, dissociates, forming H+ and OH? ions. These ions are hydrated, that is, connected to several water molecules, but for simplicity they are written in non-hydrated form

H 2 O - H + + OH?.

Based on the law of mass action, for this equilibrium:

The concentration of water molecules [H 2 O], i.e. the number of moles in 1 liter of water, can be considered constant and equal to [H 2 O] = 1000 g/l: 18 g/mol = 55.6 mol/l. From here:

TO[H 2 O] = TO(H 2 O ) = [H + ] = 10 -14 (22°C).

Ionic product of water– the product of concentrations [H + ] and – is a constant value at a constant temperature and equal to 10 -14 at 22°C.

The ionic product of water increases with increasing temperature.

pH value– negative logarithm of the concentration of hydrogen ions: pH = – log. Similarly: pOH = – log.

Taking the logarithm of the ionic product of water gives: pH + pHOH = 14.

The pH value characterizes the reaction of the medium.

If pH = 7, then [H + ] = is a neutral medium.

If pH< 7, то [Н + ] >– acidic environment.

If pH > 7, then [H + ]< – щелочная среда.

6.6. Buffer solutions

Buffer solutions are solutions that have a certain concentration of hydrogen ions. The pH of these solutions does not change when diluted and changes little when small amounts of acids and alkalis are added.

I. A solution of the weak acid HA, concentration – from the acid, and its salt with the strong base BA, concentration – from the salt. For example, an acetate buffer is a solution of acetic acid and sodium acetate: CH 3 COOH + CHgCOONa.

pH = pK acidic + log(salt/s sour).

II. A solution of the weak base BOH, concentration - from basic, and its salt with a strong acid BA, concentration - from salt. For example, an ammonia buffer is a solution of ammonium hydroxide and ammonium chloride NH 4 OH + NH 4 Cl.

pH = 14 – рК basic – log(with salt/with basic).

6.7. Hydrolysis of salts

Hydrolysis of salts– interaction of salt ions with water to form a weak electrolyte.

Examples of hydrolysis reaction equations.

I. A salt is formed by a strong base and a weak acid:

Na 2 CO 3 + H 2 O - NaHCO 3 + NaOH

2Na + + CO 3 2- + H 2 O - 2Na + + HCO 3 ? +OH?

CO 3 2- + H 2 O - HCO 3 ? + OH?, pH > 7, alkaline environment.

In the second stage, hydrolysis practically does not occur.

II. A salt is formed by a weak base and a strong acid:

AlCl 3 + H 2 O - (AlOH)Cl 2 + HCl

Al 3+ + 3Cl? + H 2 O - AlOH 2+ + 2Cl? + H + + Cl?

Al 3+ + H 2 O - AlOH 2+ + H +, pH< 7.

In the second stage, hydrolysis occurs less, and in the third stage there is practically no hydrolysis.

III. A salt is formed by a strong base and a strong acid:

K + + NO 3 ? + H 2 O ? no hydrolysis, pH? 7.

IV. A salt is formed by a weak base and a weak acid:

CH 3 COONH 4 + H 2 O - CH 3 COOH + NH 4 OH

CH 3 COO? + NH 4 + + H 2 O - CH 3 COOH + NH 4 OH, pH = 7.

In some cases, when the salt is formed by very weak bases and acids, complete hydrolysis occurs. In the solubility table for such salts the symbol is “decomposed by water”:

Al 2 S 3 + 6H 2 O = 2Al(OH) 3 v + 3H 2 S^

The possibility of complete hydrolysis should be taken into account in exchange reactions:

Al 2 (SO 4) 3 + 3Na 2 CO 3 + 3H 2 O = 2Al(OH) 3 v + 3Na 2 SO 4 + 3CO 2 ^

Degree of hydrolysish – the ratio of the concentration of hydrolyzed molecules to the total concentration of dissolved molecules.

For salts formed by a strong base and a weak acid:

= chрOH = – log, рН = 14 – рOH.

From the expression it follows that the degree of hydrolysis h(i.e. hydrolysis) increases:

a) with increasing temperature, as K(H 2 O) increases;

b) with a decrease in the dissociation of the acid forming the salt: the weaker the acid, the greater the hydrolysis;

c) with dilution: the smaller the c, the greater the hydrolysis.

For salts formed by a weak base and a strong acid

[H + ] = ch pH = – log.

For salts formed by a weak base and a weak acid

6.8. Protolytic theory of acids and bases

Protolysis– proton transfer process.

Protoliths– acids and bases that donate and accept protons.

Acid– a molecule or ion capable of donating a proton. Each acid has a corresponding conjugate base. The strength of acids is characterized by the acid constant K k.

H 2 CO 3 + H 2 O - H 3 O + + HCO 3 ?

K k = 4 ? 10 -7

3+ + H 2 O - 2+ + H 3 O +

K k = 9 ? 10 -6

Base– a molecule or ion that can accept a proton. Each base has a corresponding conjugate acid. The strength of bases is characterized by the base constant K 0.

NH3? H 2 O (H 2 O) - NH 4 + + OH?

K 0 = 1,8 ?10 -5

Ampholytes– protoliths capable of releasing and acquiring a proton.

HCO3? + H 2 O - H 3 O + + CO 3 2-

HCO3? – acid.

HCO3? + H 2 O - H 2 CO 3 + OH?

HCO3? – foundation.

For water: H 2 O+ H 2 O - H 3 O + + OH?

K(H 2 O) = [H 3 O + ] = 10 -14 and pH = – log.

Constants K k And K 0 for conjugate acids and bases are linked.

HA + H 2 O - H 3 O + + A?,

A? + H 2 O - HA + OH?,

7. Solubility constant. Solubility

In a system consisting of a solution and a precipitate, two processes take place - dissolution of the precipitate and precipitation. The equality of the rates of these two processes is a condition of equilibrium.

Saturated solution– a solution that is in equilibrium with the precipitate.

The law of mass action applied to the equilibrium between precipitate and solution gives:

Since = const,

TO = K s (AgCl) = .

In general we have:

A m B n(TV) - m A +n+n B -m

K s ( A m B n)= [A +n ] m[IN -m ] n .

Solubility constantK s(or solubility product PR) - the product of ion concentrations in a saturated solution of a slightly soluble electrolyte - is a constant value and depends only on temperature.

Solubility of a sparingly soluble substance s can be expressed in moles per liter. Depending on the size s substances can be divided into poorly soluble – s< 10 -4 моль/л, среднерастворимые – 10 -4 моль/л? s? 10 -2 mol/l and highly soluble s>10 -2 mol/l.

The solubility of compounds is related to their solubility product.

Condition for precipitation and dissolution of sediment

In the case of AgCl: AgCl - Ag + + Cl?

K s= :

a) condition of equilibrium between precipitate and solution: = Ks.

b) deposition condition: > Ks; during precipitation, ion concentrations decrease until equilibrium is established;

c) the condition for the dissolution of the precipitate or the existence of a saturated solution:< Ks; As the precipitate dissolves, the ion concentration increases until equilibrium is established.

8. Coordination compounds

Coordination (complex) compounds are compounds with a donor-acceptor bond.

For K 3:

ions of the outer sphere – 3K +,

inner sphere ion – 3-,

complexing agent – ​​Fe 3+,

ligands – 6CN?, their dentation – 1,

coordination number – 6.

Examples of complexing agents: Ag +, Cu 2+, Hg 2+, Zn 2+, Ni 2+, Fe 3+, Pt 4+, etc.

Examples of ligands: polar molecules H 2 O, NH 3, CO and anions CN?, Cl?, OH? and etc.

Coordination numbers: usually 4 or 6, less often 2, 3, etc.

Nomenclature. The anion is named first (in the nominative case), then the cation (in the genitive case). Names of some ligands: NH 3 - ammin, H 2 O - aquo, CN? – cyano, Cl? – chloro, OH? – hydroxo. Names of coordination numbers: 2 – di, 3 – three, 4 – tetra, 5 – penta, 6 – hexa. The oxidation state of the complexing agent is indicated:

Cl—diamminesilver(I) chloride;

SO 4 – tetrammine copper(II) sulfate;

K 3 – potassium hexacyanoferrate(III).

Chemical connection.

Valence bond theory assumes hybridization of the orbitals of the central atom. The location of the resulting hybrid orbitals determines the geometry of the complexes.

Diamagnetic complex ion Fe(CN) 6 4-.

Cyanide ion – donor

The iron ion Fe 2+ – acceptor – has the formula 3d 6 4s 0 4p 0. Taking into account the diamagnetic nature of the complex (all electrons are paired) and the coordination number (6 free orbitals are needed), we have d 2 sp 3-hybridization:

The complex is diamagnetic, low-spin, intraorbital, stable (no external electrons are used), octahedral ( d 2 sp 3-hybridization).

Paramagnetic complex ion FeF 6 3-.

Fluoride ion is a donor.

The iron ion Fe 3+ – acceptor – has the formula 3d 5 4s 0 4p 0 . Taking into account the paramagnetism of the complex (electrons are coupled) and the coordination number (6 free orbitals are needed), we have sp 3 d 2-hybridization:

The complex is paramagnetic, high-spin, outer-orbital, unstable (outer 4d orbitals are used), octahedral ( sp 3 d 2-hybridization).

Dissociation of coordination compounds.

Coordination compounds in solution completely dissociate into ions of the inner and outer spheres.

NO 3 > Ag(NH 3) 2 + + NO 3 ?, ? = 1.

The ions of the inner sphere, i.e., complex ions, dissociate into metal ions and ligands, like weak electrolytes, in stages.

Where K 1 , TO 2 , TO 1 _ 2 are called instability constants and characterize the dissociation of complexes: the lower the instability constant, the less the complex dissociates, the more stable it is.

Well, to complete our acquaintance with alcohols, I will also give the formula of another well-known substance - cholesterol. Not everyone knows that it is a monohydric alcohol!

|`/`\\`|<`|w>`\`/|<`/w$color(red)HO$color()>\/`|0/`|/\<`|w>|_q_q_q<-dH>:a_q|0<|dH>`/<`|wH>`\|dH; #a_(A-72)<_(A-120,d+)>-/-/<->`\

I marked the hydroxyl group in it in red.

Carboxylic acids

Any winemaker knows that wine should be stored without access to air. Otherwise it will turn sour. But chemists know the reason - if you add another oxygen atom to an alcohol, you get an acid.
Let's look at the formulas of acids that are obtained from alcohols already familiar to us:

Substance			Skeletal formula	Gross formula
Methane acid (formic acid)	H/C`|O|\OH	HCOOH	O//\OH
Ethanoic acid (acetic acid)	H-C-C\O-H; H|#C|H	CH3-COOH	/`|O|\OH
Propanic acid (methylacetic acid)	H-C-C-C\O-H; H|#2|H; H|#3|H	CH3-CH2-COOH	\/`|O|\OH
Butanoic acid (butyric acid)	H-C-C-C-C\O-H; H|#2|H; H|#3|H; H|#4|H	CH3-CH2-CH2-COOH	/\/`|O|\OH
Generalized formula	(R)-C\O-H	(R)-COOH or (R)-CO2H	(R)/`|O|\OH

A distinctive feature of organic acids is the presence of a carboxyl group (COOH), which gives such substances acidic properties.

Anyone who has tried vinegar knows that it is very sour. The reason for this is the presence of acetic acid in it. Typically table vinegar contains between 3 and 15% acetic acid, with the rest (mostly) water. Consumption of acetic acid in undiluted form poses a danger to life.

Carboxylic acids can have multiple carboxyl groups. In this case they are called: dibasic, tribasic etc...

Food products contain many other organic acids. Here are just a few of them:

The name of these acids corresponds to the food products in which they are contained. By the way, please note that here there are acids that also have a hydroxyl group, characteristic of alcohols. Such substances are called hydroxycarboxylic acids(or hydroxy acids).
Below, under each of the acids, there is a sign specifying the name of the group of organic substances to which it belongs.

Radicals

Radicals are another concept that has influenced chemical formulas. The word itself is probably known to everyone, but in chemistry radicals have nothing in common with politicians, rebels and other citizens with an active position.
Here these are just fragments of molecules. And now we will figure out what makes them special and get acquainted with a new way of writing chemical formulas.

Generalized formulas have already been mentioned several times in the text: alcohols - (R)-OH and carboxylic acids - (R)-COOH. Let me remind you that -OH and -COOH are functional groups. But R is a radical. It’s not for nothing that he is depicted as the letter R.

To be more specific, a monovalent radical is a part of a molecule lacking one hydrogen atom. Well, if you subtract two hydrogen atoms, you get a divalent radical.

Radicals in chemistry received their own names. Some of them even received Latin designations similar to the designations of the elements. And besides, sometimes in formulas radicals can be indicated in abbreviated form, more reminiscent of gross formulas.
All this is demonstrated in the following table.

Name	Structural formula	Designation	Brief formula	Example of alcohol
Methyl	CH3-()	Me	CH3	(Me)-OH	CH3OH
Ethyl	CH3-CH2-()	Et	C2H5	(Et)-OH	C2H5OH
I cut through	CH3-CH2-CH2-()	Pr	C3H7	(Pr)-OH	C3H7OH
Isopropyl	H3C\CH(*`/H3C*)-()	i-Pr	C3H7	(i-Pr)-OH	(CH3)2CHOH
Phenyl	`/`=`\//-\\-{}	Ph	C6H5	(Ph)-OH	C6H5OH

I think everything is clear here. I just want to draw your attention to the column where examples of alcohols are given. Some radicals are written in a form that resembles the gross formula, but the functional group is written separately. For example, CH3-CH2-OH turns into C2H5OH.
And for branched chains like isopropyl, structures with brackets are used.

There is also such a phenomenon as free radicals. These are radicals that, for some reason, have separated from functional groups. In this case, one of the rules with which we began studying the formulas is violated: the number of chemical bonds no longer corresponds to the valence of one of the atoms. Well, or we can say that one of the connections becomes open at one end. Free radicals usually live for a short time as the molecules tend to return to a stable state.

Introduction to nitrogen. Amines

I propose to get acquainted with another element that is part of many organic compounds. This nitrogen.
It is denoted by the Latin letter N and has a valency of three.

Let's see what substances are obtained if nitrogen is added to the familiar hydrocarbons:

Substance	Expanded structural formula	Simplified structural formula	Skeletal formula	Gross formula
Aminomethane (methylamine)	H-C-N\H;H|#C|H	CH3-NH2	\NH2
Aminoethane (ethylamine)	H-C-C-N\H;H|#C|H;H|#3|H	CH3-CH2-NH2	/\NH2
Dimethylamine	H-C-N<`|H>-C-H; H|#-3|H; H|#2|H	$L(1.3)H/N<_(A80,w+)CH3>\dCH3	/N<_(y-.5)H>\
Aminobenzene (Aniline)	H\N|C\\C|C<\H>`//C<|H>`\C<`/H>`||C<`\H>/	NH2|C\\CH|CH`//C<_(y.5)H>`\HC`||HC/	NH2|\|`/`\`|/_o
Triethylamine	$slope(45)H-C-C/N\C-C-H;H|#2|H; H|#3|H; H|#5|H;H|#6|H; #N`|C<`-H><-H>`|C<`-H><-H>`|H	CH3-CH2-N<`|CH2-CH3>-CH2-CH3	\/N<`|/>\|

As you probably already guessed from the names, all these substances are united under the general name amines. The functional group ()-NH2 is called amino group. Here are some general formulas of amines:

In general, there are no special innovations here. If these formulas are clear to you, then you can safely engage in further study of organic chemistry using a textbook or the Internet.
But I would also like to talk about formulas in inorganic chemistry. You will see how easy it is to understand them after studying the structure of organic molecules.

Rational formulas

It should not be concluded that inorganic chemistry is easier than organic chemistry. Of course, inorganic molecules tend to look much simpler because they don't tend to form complex structures like hydrocarbons. But then we have to study more than a hundred elements that make up the periodic table. And these elements tend to combine according to their chemical properties, but with numerous exceptions.

So, I won’t tell you any of this. The topic of my article is chemical formulas. And with them everything is relatively simple.
Most often used in inorganic chemistry rational formulas. And now we’ll figure out how they differ from those already familiar to us.

First, let's get acquainted with another element - calcium. This is also a very common element.
It is designated Ca and has a valency of two. Let's see what compounds it forms with the carbon, oxygen and hydrogen we know.

Substance	Structural formula	Rational formula	Gross formula
Calcium oxide	Ca=O	CaO
Calcium hydroxide	H-O-Ca-O-H	Ca(OH)2
Calcium carbonate	$slope(45)Ca`/O\C|O`|/O`\#1	CaCO3
Calcium bicarbonate	HO/`|O|\O/Ca\O/`|O|\OH	Ca(HCO3)2
Carbonic acid	H|O\C|O`|/O`|H	H2CO3

At first glance, you can see that the rational formula is something between a structural and a gross formula. But it is not yet very clear how they are obtained. To understand the meaning of these formulas, you need to consider the chemical reactions in which substances participate.

Calcium in its pure form is a soft white metal. It does not occur in nature. But it is quite possible to buy it at a chemical store. It is usually stored in special jars without access to air. Because in air it reacts with oxygen. Actually, that’s why it doesn’t occur in nature.
So, the reaction of calcium with oxygen:

2Ca + O2 -> 2CaO

The number 2 before the formula of a substance means that 2 molecules are involved in the reaction.
Calcium and oxygen produce calcium oxide. This substance also does not occur in nature because it reacts with water:

CaO + H2O -> Ca(OH2)

The result is calcium hydroxide. If you look closely at its structural formula (in the previous table), you can see that it is formed by one calcium atom and two hydroxyl groups, with which we are already familiar.
These are the laws of chemistry: if a hydroxyl group is added to an organic substance, an alcohol is obtained, and if it is added to a metal, a hydroxide is obtained.

But calcium hydroxide does not occur in nature due to the presence of carbon dioxide in the air. I think everyone has heard about this gas. It is formed during the respiration of people and animals, the combustion of coal and petroleum products, during fires and volcanic eruptions. Therefore, it is always present in the air. But it also dissolves quite well in water, forming carbonic acid:

CO2 + H2O<=>H2CO3

Sign<=>indicates that the reaction can proceed in both directions under the same conditions.

Thus, calcium hydroxide, dissolved in water, reacts with carbonic acid and turns into slightly soluble calcium carbonate:

Ca(OH)2 + H2CO3 -> CaCO3"|v" + 2H2O

A down arrow means that as a result of the reaction the substance precipitates.
With further contact of calcium carbonate with carbon dioxide in the presence of water, a reversible reaction occurs to form an acidic salt - calcium bicarbonate, which is highly soluble in water

CaCO3 + CO2 + H2O<=>Ca(HCO3)2

This process affects the hardness of the water. When the temperature rises, bicarbonate turns back into carbonate. Therefore, in regions with hard water, scale forms in kettles.

Chalk, limestone, marble, tuff and many other minerals are largely composed of calcium carbonate. It is also found in corals, mollusk shells, animal bones, etc...
But if calcium carbonate is heated over very high heat, it will turn into calcium oxide and carbon dioxide.

This short story about the calcium cycle in nature should explain why rational formulas are needed. So, rational formulas are written so that the functional groups are visible. In our case it is:

In addition, individual elements - Ca, H, O (in oxides) - are also independent groups.

Ions

I think it's time to get acquainted with ions. This word is probably familiar to everyone. And after studying the functional groups, it doesn’t cost us anything to figure out what these ions are.

In general, the nature of chemical bonds is usually that some elements give up electrons while others gain them. Electrons are particles with a negative charge. An element with a full complement of electrons has zero charge. If he gave away an electron, then its charge becomes positive, and if he accepted it, then it becomes negative. For example, hydrogen has only one electron, which it gives up quite easily, turning into a positive ion. There is a special entry for this in chemical formulas:

H2O<=>H^+ + OH^-

Here we see that as a result electrolytic dissociation water breaks down into a positively charged hydrogen ion and a negatively charged OH group. The OH^- ion is called hydroxide ion. It should not be confused with the hydroxyl group, which is not an ion, but part of some kind of molecule. The + or - sign in the upper right corner shows the charge of the ion.
But carbonic acid never exists as an independent substance. In fact, it is a mixture of hydrogen ions and carbonate ions (or bicarbonate ions):

H2CO3 = H^+ + HCO3^-<=>2H^+ + CO3^2-

The carbonate ion has a charge of 2-. This means that two electrons have been added to it.

Negatively charged ions are called anions. Typically these include acidic residues.
Positively charged ions - cations. Most often these are hydrogen and metals.

And here you can probably fully understand the meaning of rational formulas. The cation is written in them first, followed by the anion. Even if the formula does not contain any charges.

You probably already guess that ions can be described not only by rational formulas. Here is the skeletal formula of the bicarbonate anion:

Here the charge is indicated directly next to the oxygen atom, which received an extra electron and therefore lost one line. Simply put, each extra electron reduces the number of chemical bonds depicted in the structural formula. On the other hand, if some node of the structural formula has a + sign, then it has an additional stick. As always, this fact needs to be demonstrated with an example. But among the substances familiar to us there is not a single cation that consists of several atoms.
And such a substance is ammonia. Its aqueous solution is often called ammonia and is included in any first aid kit. Ammonia is a compound of hydrogen and nitrogen and has the rational formula NH3. Consider the chemical reaction that occurs when ammonia is dissolved in water:

NH3 + H2O<=>NH4^+ + OH^-

The same thing, but using structural formulas:

H|N<`/H>\H + H-O-H<=>H|N^+<_(A75,w+)H><_(A15,d+)H>`/H + O`^-# -H

On the right side we see two ions. They were formed as a result of one hydrogen atom moving from a water molecule to an ammonia molecule. But this atom moved without its electron. The anion is already familiar to us - it is a hydroxide ion. And the cation is called ammonium. It exhibits properties similar to metals. For example, it may combine with an acidic residue. The substance formed by combining ammonium with a carbonate anion is called ammonium carbonate: (NH4)2CO3.
Here is the reaction equation for the interaction of ammonium with a carbonate anion, written in the form of structural formulas:

2H|N^+<`/H><_(A75,w+)H>_(A15,d+)H + O^-\C|O`|/O^-<=>H|N^+<`/H><_(A75,w+)H>_(A15,d+)H`|0O^-\C|O`|/O^-|0H_(A-15,d-)N^+<_(A105,w+)H><\H>`|H

But in this form the reaction equation is given for demonstration purposes. Typically equations use rational formulas:

2NH4^+ + CO3^2-<=>(NH4)2CO3

Hill system

So, we can assume that we have already studied structural and rational formulas. But there is another issue that is worth considering in more detail. How do gross formulas differ from rational ones?
We know why the rational formula of carbonic acid is written H2CO3, and not some other way. (The two hydrogen cations come first, followed by the carbonate anion.) But why is the gross formula written CH2O3?

In principle, the rational formula of carbonic acid may well be considered a true formula, because it has no repeating elements. Unlike NH4OH or Ca(OH)2.
But an additional rule is very often applied to gross formulas, which determines the order of elements. The rule is quite simple: carbon is placed first, then hydrogen, and then the remaining elements in alphabetical order.
So CH2O3 comes out - carbon, hydrogen, oxygen. This is called the Hill system. It is used in almost all chemical reference books. And in this article too.

A little about the easyChem system

Instead of a conclusion, I would like to talk about the easyChem system. It is designed so that all the formulas that we discussed here can be easily inserted into the text. Actually, all the formulas in this article are drawn using easyChem.

Why do we even need some kind of system for deriving formulas? The thing is that the standard way to display information in Internet browsers is hypertext markup language (HTML). It is focused on processing text information.

Rational and gross formulas can be depicted using text. Even some simplified structural formulas can also be written in text, for example alcohol CH3-CH2-OH. Although for this you would have to use the following entry in HTML: CH ₃-CH ₂-OH.
This of course creates some difficulties, but you can live with them. But how to depict the structural formula? In principle, you can use a monospace font:

H H | | H-C-C-O-H | | H H Of course it doesn’t look very nice, but it’s also doable.

The real problem comes when trying to draw benzene rings and when using skeletal formulas. There is no other way left except connecting a raster image. Rasters are stored in separate files. Browsers can include images in gif, png or jpeg format.
To create such files, a graphic editor is required. For example, Photoshop. But I have been familiar with Photoshop for more than 10 years and I can say for sure that it is very poorly suited for depicting chemical formulas.
Molecular editors cope with this task much better. But with a large number of formulas, each of which is stored in a separate file, it is quite easy to get confused in them.
For example, the number of formulas in this article is . They are displayed in the form of graphic images (the rest using HTML tools).

The easyChem system allows you to store all formulas directly in an HTML document in text form. In my opinion, this is very convenient.
In addition, the gross formulas in this article are calculated automatically. Because easyChem works in two stages: first the text description is converted into an information structure (graph), and then various actions can be performed on this structure. Among them, the following functions can be noted: calculation of molecular weight, conversion to a gross formula, checking for the possibility of output as text, graphic and text rendering.

Thus, to prepare this article, I only used a text editor. Moreover, I didn’t have to think about which of the formulas would be graphic and which would be text.

Here are a few examples that reveal the secret of preparing the text of an article: Descriptions from the left column are automatically turned into formulas in the second column.
In the first line, the description of the rational formula is very similar to the displayed result. The only difference is that the numerical coefficients are displayed interlinearly.
In the second line, the expanded formula is given in the form of three separate chains separated by a symbol; I think it is easy to see that the textual description is in many ways reminiscent of the actions that would be required to depict the formula with a pencil on paper.
The third line demonstrates the use of slanted lines using the \ and / symbols. The ` (backtick) sign means the line is drawn from right to left (or bottom to top).

There is much more detailed documentation on using the easyChem system here.

Let me finish this article and wish you good luck in studying chemistry.

A brief explanatory dictionary of terms used in the article

Hydrocarbons Substances consisting of carbon and hydrogen. They differ from each other in the structure of their molecules. Structural formulas are schematic images of molecules, where atoms are denoted by Latin letters and chemical bonds by dashes. Structural formulas are expanded, simplified and skeletal. Expanded structural formulas are structural formulas where each atom is represented as a separate node. Simplified structural formulas are those structural formulas where hydrogen atoms are written next to the element with which they are associated. And if more than one hydrogen is attached to one atom, then the amount is written as a number. We can also say that groups act as nodes in simplified formulas. Skeletal formulas are structural formulas where carbon atoms are depicted as empty nodes. The number of hydrogen atoms bonded to each carbon atom is equal to 4 minus the number of bonds that converge at the site. For knots formed not by carbon, the rules of simplified formulas apply. Gross formula (aka true formula) - a list of all chemical elements that make up the molecule, indicating the number of atoms in the form of a number (if there is one atom, then the unit is not written) Hill system - a rule that determines the order of atoms in the gross formula formula: carbon is placed first, then hydrogen, and then the remaining elements in alphabetical order. This is a system that is used very often. And all the gross formulas in this article are written according to the Hill system. Functional groups Stable combinations of atoms that are conserved during chemical reactions. Often functional groups have their own names and affect the chemical properties and scientific name of the substance