Chemical properties of bases table. Grounds

Metal in the voltage series after H 2

Metal sulfate in max s.o.

+
+ +

Metal (others)

+
+ +

HNO 3 concentrated

Metal nitrate in max d.o.

Metal (others; Al, Cr, Fe, Co, Ni when heated)

Nitratemetal

la in max s.o.

+
+

Metal (the rest in the yard of stresses up to N 2)

NO/N 2 O/N 2 /NH 3 (NH 4 NO 3)

depending on conditions

+

Metal (the rest in the series of stresses after H 2)

Fig.2. INTERACTION OF ACIDS WITH METALS

SALT

Salts – These are complex substances that dissociate in solutions to form positively charged ions (cations - basic residues), with the exception of hydrogen ions, and negatively charged ions (anions - acidic residues), other than hydroxide ions.

Bases, amphoteric hydroxides

Bases are complex substances consisting of metal atoms and one or more hydroxyl groups (-OH). The general formula is Me +y (OH) y, where y is the number of hydroxo groups equal to the oxidation state of the metal Me. The table shows the classification of bases.

Properties of alkalis, hydroxides of alkali and alkaline earth metals

1. Aqueous solutions of alkalis are soapy to the touch and change the color of indicators: litmus - blue, phenolphthalein - crimson.

2. Aqueous solutions dissociate:

3. Interact with acids, entering into an exchange reaction:

Polyacid bases can give medium and basic salts:

4. React with acidic oxides, forming medium and acidic salts depending on the basicity of the acid corresponding to this oxide:

5. Interact with amphoteric oxides and hydroxides:

a) fusion:

b) in solutions:

6. Interact with water-soluble salts if a precipitate or gas is formed:

Insoluble bases (Cr(OH) 2, Mn(OH) 2, etc.) interact with acids and decompose when heated:

Amphoteric hydroxides

Amphoteric compounds are compounds that, depending on conditions, can be both donors of hydrogen cations and exhibit acidic properties, and their acceptors, i.e., exhibit basic properties.

Chemical properties of amphoteric compounds

1. Interacting with strong acids, they exhibit basic properties:

Zn(OH) 2 + 2HCl = ZnCl 2 + 2H 2 O

2. Interacting with alkalis - strong bases, they exhibit acidic properties:

Zn(OH) 2 + 2NaOH = Na 2 ( complex salt)

Al(OH) 3 + NaOH = Na ( complex salt)

Complex compounds are those in which at least one covalent bond is formed by a donor-acceptor mechanism.

The general method for preparing bases is based on exchange reactions, with the help of which both insoluble and soluble bases can be obtained.

CuSO 4 + 2KOH = Cu(OH) 2 ↓ + K 2 SO 4

K 2 CO 3 + Ba(OH) 2 = 2 KOH + BaCO 3 ↓

When soluble bases are obtained by this method, an insoluble salt precipitates.

When preparing water-insoluble bases with amphoteric properties, excess alkali should be avoided, since dissolution of the amphoteric base may occur, for example:

AlCl 3 + 4KOH = K[Al(OH) 4 ] + 3KCl

In such cases, ammonium hydroxide is used to obtain hydroxides, in which amphoteric hydroxides do not dissolve:

AlCl 3 + 3NH 3 + ZH 2 O = Al(OH) 3 ↓ + 3NH 4 Cl

Silver and mercury hydroxides decompose so easily that when trying to obtain them by exchange reaction, instead of hydroxides, oxides precipitate:

2AgNO 3 + 2KOH = Ag 2 O↓ + H 2 O + 2KNO 3

In industry, alkalis are usually obtained by electrolysis of aqueous solutions of chlorides.

2NaCl + 2H 2 O → ϟ → 2NaOH + H 2 + Cl 2

Alkalis can also be obtained by reacting alkali and alkaline earth metals or their oxides with water.

2Li + 2H 2 O = 2LiOH + H 2

SrO + H 2 O = Sr(OH) 2

Acids

Acids are complex substances whose molecules consist of hydrogen atoms that can be replaced by metal atoms and acidic residues. Under normal conditions, acids can be solid (phosphoric H 3 PO 4; silicon H 2 SiO 3) and liquid (in its pure form, sulfuric acid H 2 SO 4 will be a liquid).

Gases such as hydrogen chloride HCl, hydrogen bromide HBr, hydrogen sulfide H 2 S form the corresponding acids in aqueous solutions. The number of hydrogen ions formed by each acid molecule during dissociation determines the charge of the acid residue (anion) and the basicity of the acid.

According to protolytic theory of acids and bases, proposed simultaneously by the Danish chemist Brønsted and the English chemist Lowry, an acid is a substance splitting off with this reaction protons, A basis- a substance that can accept protons.

acid → base + H +

Based on such ideas, it is clear basic properties of ammonia, which, due to the presence of a lone electron pair at the nitrogen atom, effectively accepts a proton when interacting with acids, forming an ammonium ion through a donor-acceptor bond.

HNO 3 + NH 3 ⇆ NH 4 + + NO 3 —

acid base acid base

More general definition of acids and bases proposed by the American chemist G. Lewis. He suggested that acid-base interactions are completely do not necessarily occur with the transfer of protones. In the Lewis determination of acids and bases, the main role in chemical reactions is played by electron pairs

Cations, anions, or neutral molecules that can accept one or more pairs of electrons are called Lewis acids.

For example, aluminum fluoride AlF 3 is an acid, since it is able to accept an electron pair when interacting with ammonia.

AlF 3 + :NH 3 ⇆ :

Cations, anions, or neutral molecules capable of donating electron pairs are called Lewis bases (ammonia is a base).

Lewis's definition covers all acid-base processes that were considered by previously proposed theories. The table compares the definitions of acids and bases currently used.

Nomenclature of acids

Since there are different definitions of acids, their classification and nomenclature are rather arbitrary.

According to the number of hydrogen atoms capable of elimination in an aqueous solution, acids are divided into monobasic(e.g. HF, HNO 2), dibasic(H 2 CO 3, H 2 SO 4) and tribasic(H 3 PO 4).

According to the composition of the acid, they are divided into oxygen-free(HCl, H 2 S) and oxygen-containing(HClO 4, HNO 3).

Usually names of oxygen-containing acids are derived from the name of the non-metal with the addition of the endings -kai, -vaya, if the oxidation state of the non-metal is equal to the group number. As the oxidation state decreases, the suffixes change (in order of decreasing oxidation state of the metal): -opaque, rusty, -ovish:

If we consider the polarity of the hydrogen-nonmetal bond within a period, we can easily relate the polarity of this bond to the position of the element in the Periodic Table. From metal atoms, which easily lose valence electrons, hydrogen atoms accept these electrons, forming a stable two-electron shell like the shell of a helium atom, and give ionic metal hydrides.

In hydrogen compounds of elements of groups III-IV of the Periodic Table, boron, aluminum, carbon, and silicon form covalent, weakly polar bonds with hydrogen atoms that are not prone to dissociation. For elements of groups V-VII of the Periodic Table, within a period, the polarity of the nonmetal-hydrogen bond increases with the charge of the atom, but the distribution of charges in the resulting dipole is different than in hydrogen compounds of elements that tend to donate electrons. Non-metal atoms, which require several electrons to complete the electron shell, attract (polarize) a pair of bonding electrons the more strongly, the greater the nuclear charge. Therefore, in the series CH 4 - NH 3 - H 2 O - HF or SiH 4 - PH 3 - H 2 S - HCl, bonds with hydrogen atoms, while remaining covalent, become more polar in nature, and the hydrogen atom in the element-hydrogen bond dipole becomes more electropositive. If polar molecules find themselves in a polar solvent, a process of electrolytic dissociation can occur.

Let us discuss the behavior of oxygen-containing acids in aqueous solutions. These acids have an H-O-E bond and, naturally, the polarity of the H-O bond is influenced by the O-E bond. Therefore, these acids, as a rule, dissociate more easily than water.

H 2 SO 3 + H 2 O ⇆ H 3 O + + HSO 3

HNO 3 + H 2 O ⇆ H 3 O + + NO 3

Let's look at a few examples properties of oxygen-containing acids, formed by elements that are capable of exhibiting different degrees of oxidation. It is known that hypochlorous acid HClO very weak chlorous acid HClO 2 also weak, but stronger than hypochlorous, hypochlorous acid HClO 3 strong. Perchloric acid HClO 4 is one of the strongest inorganic acids.

For acidic dissociation (with the elimination of the H ion), the cleavage of the O-H bond is necessary. How can we explain the decrease in the strength of this bond in the series HClO - HClO 2 - HClO 3 - HClO 4? In this series, the number of oxygen atoms associated with the central chlorine atom increases. Each time a new oxygen-chlorine bond is formed, electron density is drawn from the chlorine atom, and therefore from the O-Cl single bond. As a result, the electron density partially leaves the O-H bond, which is weakened as a result.

This pattern - strengthening of acidic properties with increasing degree of oxidation of the central atom - characteristic not only of chlorine, but also of other elements. For example, nitric acid HNO 3, in which the oxidation state of nitrogen is +5, is stronger than nitrous acid HNO 2 (the oxidation state of nitrogen is +3); sulfuric acid H 2 SO 4 (S +6) is stronger than sulfurous acid H 2 SO 3 (S +4).

Obtaining acids

1. Oxygen-free acids can be obtained by direct combination of non-metals with hydrogen.

H 2 + Cl 2 → 2HCl,

H 2 + S ⇆ H 2 S

2. Some oxygen-containing acids can be obtained interaction of acid oxides with water.

3. Both oxygen-free and oxygen-containing acids can be obtained by metabolic reactions between salts and other acids.

BaBr 2 + H 2 SO 4 = BaSO 4 ↓ + 2НВr

CuSO 4 + H 2 S = H 2 SO 4 + CuS↓

FeS + H 2 SO 4 (pa zb) = H 2 S + FeSO 4

NaCl (T) + H 2 SO 4 (conc) = HCl + NaHSO 4

AgNO 3 + HCl = AgCl↓ + HNO 3

CaCO 3 + 2HBr = CaBr 2 + CO 2 + H 2 O

4. Some acids can be obtained using redox reactions.

H 2 O 2 + SO 2 = H 2 SO 4

3P + 5HNO 3 + 2H 2 O = ZN 3 PO 4 + 5NO 2

Sour taste, effect on indicators, electrical conductivity, interaction with metals, basic and amphoteric oxides, bases and salts, formation of esters with alcohols - these properties are common to inorganic and organic acids.

can be divided into two types of reactions:

1) are common For acids reactions are associated with the formation of hydronium ion H 3 O + in aqueous solutions;

2) specific(i.e. characteristic) reactions specific acids.

The hydrogen ion can enter into redox reaction, reducing to hydrogen, as well as in a compound reaction with negatively charged or neutral particles having lone pairs of electrons, i.e. acid-base reactions.

The general properties of acids include reactions of acids with metals in the voltage series up to hydrogen, for example:

Zn + 2Н + = Zn 2+ + Н 2

Acid-base reactions include reactions with basic oxides and bases, as well as with intermediate, basic, and sometimes acidic salts.

2 CO 3 + 4HBr = 2CuBr 2 + CO 2 + 3H 2 O

Mg(HCO 3) 2 + 2HCl = MgCl 2 + 2CO 2 + 2H 2 O

2KHSO 3 + H 2 SO 4 = K 2 SO 4 + 2SO 2 + 2H 2 O

Note that polybasic acids dissociate stepwise, and at each subsequent step the dissociation is more difficult, therefore, with an excess of acid, acidic salts are most often formed, rather than average ones.

Ca 3 (PO 4) 2 + 4H 3 PO 4 = 3Ca (H 2 PO 4) 2

Na 2 S + H 3 PO 4 = Na 2 HPO 4 + H 2 S

NaOH + H 3 PO 4 = NaH 2 PO 4 + H 2 O

KOH + H 2 S = KHS + H 2 O

At first glance, the formation of acid salts may seem surprising monobasic hydrofluoric acid. However, this fact can be explained. Unlike all other hydrohalic acids, hydrofluoric acid in solutions is partially polymerized (due to the formation of hydrogen bonds) and various particles (HF) X may be present in it, namely H 2 F 2, H 3 F 3, etc.

A special case of acid-base equilibrium - reactions of acids and bases with indicators that change their color depending on the acidity of the solution. Indicators are used in qualitative analysis to detect acids and bases in solutions.

The most commonly used indicators are litmus(V neutral environment purple, V sour - red, V alkaline - blue), methyl orange(V sour environment red, V neutral - orange, V alkaline - yellow), phenolphthalein(V highly alkaline environment raspberry red, V neutral and acidic - colorless).

Specific properties different acids can be of two types: firstly, reactions leading to the formation insoluble salts, and secondly, redox transformations. If the reactions associated with the presence of the H + ion are common to all acids (qualitative reactions for detecting acids), specific reactions are used as qualitative reactions for individual acids:

Ag + + Cl - = AgCl (white precipitate)

Ba 2+ + SO 4 2- = BaSO 4 (white precipitate)

3Ag + + PO 4 3 - = Ag 3 PO 4 (yellow precipitate)

Some specific reactions of acids are due to their redox properties.

Anoxic acids in an aqueous solution can only be oxidized.

2KMnO 4 + 16HCl = 5Сl 2 + 2КСl + 2МnСl 2 + 8Н 2 O

H 2 S + Br 2 = S + 2НВг

Oxygen-containing acids can be oxidized only if the central atom in them is in a lower or intermediate oxidation state, as, for example, in sulfurous acid:

H 2 SO 3 + Cl 2 + H 2 O = H 2 SO 4 + 2HCl

Many oxygen-containing acids, in which the central atom has the maximum oxidation state (S +6, N +5, Cr +6), exhibit the properties of strong oxidizing agents. Concentrated H 2 SO 4 is a strong oxidizing agent.

Cu + 2H 2 SO 4 (conc) = CuSO 4 + SO 2 + 2H 2 O

Pb + 4HNO 3 = Pb(NO 3) 2 + 2NO 2 + 2H 2 O

C + 2H 2 SO 4 (conc) = CO 2 + 2SO 2 + 2H 2 O

It should be remembered that:

• Acid solutions react with metals that are to the left of hydrogen in the electrochemical voltage series, subject to a number of conditions, the most important of which is the formation of a soluble salt as a result of the reaction. The interaction of HNO 3 and H 2 SO 4 (conc.) with metals proceeds differently.

Concentrated sulfuric acid in the cold passivates aluminum, iron, and chromium.

• In water, acids dissociate into hydrogen cations and anions of acid residues, for example:

• Inorganic and organic acids react with basic and amphoteric oxides, provided that a soluble salt is formed:

• Both acids react with bases. Polybasic acids can form both intermediate and acid salts (these are neutralization reactions):

• The reaction between acids and salts occurs only if a precipitate or gas is formed:

The interaction of H 3 PO 4 with limestone will stop due to the formation of the last insoluble precipitate of Ca 3 (PO 4) 2 on the surface.

The peculiarities of the properties of nitric HNO 3 and concentrated sulfuric H 2 SO 4 (conc.) acids are due to the fact that when they interact with simple substances (metals and non-metals), the oxidizing agents will not be H + cations, but nitrate and sulfate ions. It is logical to expect that as a result of such reactions, not hydrogen H2 is formed, but other substances are obtained: necessarily salt and water, as well as one of the products of the reduction of nitrate or sulfate ions, depending on the concentration of acids, the position of the metal in the voltage series and reaction conditions (temperature, degree of metal grinding, etc.).

These features of the chemical behavior of HNO 3 and H 2 SO 4 (conc.) clearly illustrate the thesis of the theory of chemical structure about the mutual influence of atoms in the molecules of substances.

The concepts of volatility and stability (stability) are often confused. Volatile acids are acids whose molecules easily pass into a gaseous state, that is, evaporate. For example, hydrochloric acid is a volatile but stable acid. It is impossible to judge the volatility of unstable acids. For example, non-volatile, insoluble silicic acid decomposes into water and SiO 2. Aqueous solutions of hydrochloric, nitric, sulfuric, phosphoric and a number of other acids are colorless. An aqueous solution of chromic acid H 2 CrO 4 is yellow in color, and manganese acid HMnO 4 is crimson.

Reference material for taking the test:

Mendeleev table

Solubility table

After reading the article, you will be able to separate substances into salts, acids and bases. The article describes what the pH of a solution is and what general properties acids and bases have.

Like metals and nonmetals, acids and bases are the division of substances based on similar properties. The first theory of acids and bases belonged to the Swedish scientist Arrhenius. According to Arrhenius, an acid is a class of substances that, when reacting with water, dissociate (decay), forming the hydrogen cation H +. Arrhenius bases in aqueous solution form OH - anions. The next theory was proposed in 1923 by scientists Bronsted and Lowry. The Brønsted-Lowry theory defines acids as substances capable of donating a proton in a reaction (a hydrogen cation is called a proton in reactions). Bases, accordingly, are substances that can accept a proton in a reaction. The currently relevant theory is the Lewis theory. Lewis theory defines acids as molecules or ions capable of accepting electron pairs, thereby forming Lewis adducts (an adduct is a compound formed by combining two reactants without forming by-products).

In inorganic chemistry, as a rule, an acid means a Bronsted-Lowry acid, that is, substances capable of donating a proton. If they mean the definition of a Lewis acid, then in the text such an acid is called a Lewis acid. These rules apply to acids and bases.

Dissociation

Dissociation is the process of decomposition of a substance into ions in solutions or melts. For example, the dissociation of hydrochloric acid is the decomposition of HCl into H + and Cl -.

Properties of acids and bases

Bases tend to feel soapy to the touch, while acids generally taste sour.

When a base reacts with many cations, a precipitate is formed. When an acid reacts with anions, a gas is usually released.

Commonly used acids:
H 2 O, H 3 O +, CH 3 CO 2 H, H 2 SO 4, HSO 4 −, HCl, CH 3 OH, NH 3

Commonly used bases:
OH − , H 2 O , CH 3 CO 2 − , HSO 4 − , SO 4 2 − , Cl −

Strong and weak acids and bases

Strong acids

Such acids that completely dissociate in water, producing hydrogen cations H + and anions. An example of a strong acid is hydrochloric acid HCl:

HCl (solution) + H 2 O (l) → H 3 O + (solution) + Cl - (solution)

Examples of strong acids: HCl, HBr, HF, HNO 3, H 2 SO 4, HClO 4

List of strong acids

• HCl - hydrochloric acid
• HBr - hydrogen bromide
• HI - hydrogen iodide
• HNO 3 - nitric acid
• HClO 4 - perchloric acid
• H 2 SO 4 - sulfuric acid

Weak acids

Only partially dissolved in water, for example, HF:

HF (solution) + H2O (l) → H3O + (solution) + F - (solution) - in such a reaction more than 90% of the acid does not dissociate:
= < 0,01M для вещества 0,1М

Strong and weak acids can be distinguished by measuring the conductivity of solutions: conductivity depends on the number of ions, the stronger the acid, the more dissociated it is, therefore, the stronger the acid, the higher the conductivity.

List of weak acids

• HF hydrogen fluoride
• H 3 PO 4 phosphoric
• H 2 SO 3 sulfurous
• H 2 S hydrogen sulfide
• H 2 CO 3 coal
• H 2 SiO 3 silicon

Strong grounds

Strong bases completely dissociate in water:

NaOH (solution) + H 2 O ↔ NH 4

Strong bases include metal hydroxides of the first (alkalines, alkali metals) and second (alkalinotherrenes, alkaline earth metals) groups.

List of strong bases

• NaOH sodium hydroxide (caustic soda)
• KOH potassium hydroxide (caustic potash)
• LiOH lithium hydroxide
• Ba(OH) 2 barium hydroxide
• Ca(OH) 2 calcium hydroxide (slaked lime)

Weak foundations

In a reversible reaction in the presence of water, it forms OH - ions:

NH 3 (solution) + H 2 O ↔ NH + 4 (solution) + OH - (solution)

Most weak bases are anions:

F - (solution) + H 2 O ↔ HF (solution) + OH - (solution)

List of weak bases

• Mg(OH) 2 magnesium hydroxide
• Fe(OH) 2 iron(II) hydroxide
• Zn(OH) 2 zinc hydroxide
• NH 4 OH ammonium hydroxide
• Fe(OH) 3 iron(III) hydroxide

Reactions of acids and bases

Strong acid and strong base

This reaction is called neutralization: when the amount of reagents is sufficient to completely dissociate the acid and base, the resulting solution will be neutral.

Example:
H 3 O + + OH - ↔ 2H 2 O

Weak base and weak acid

General type of reaction:
Weak base (solution) + H 2 O ↔ Weak acid (solution) + OH - (solution)

Strong base and weak acid

The base dissociates completely, the acid dissociates partially, the resulting solution has weak properties of a base:

HX (solution) + OH - (solution) ↔ H 2 O + X - (solution)

Strong acid and weak base

The acid dissociates completely, the base does not dissociate completely:

Dissociation of water

Dissociation is the breakdown of a substance into its component molecules. The properties of an acid or base depend on the equilibrium that is present in water:

H 2 O + H 2 O ↔ H 3 O + (solution) + OH - (solution)
K c = / 2
The equilibrium constant of water at t=25°: K c = 1.83⋅10 -6, the following equality also holds: = 10 -14, which is called the dissociation constant of water. For pure water = = 10 -7, hence -lg = 7.0.

This value (-lg) is called pH - hydrogen potential. If pH< 7, то вещество имеет кислотные свойства, если pH >7, then the substance has basic properties.

Methods for determining pH

Instrumental method

A special device, a pH meter, is a device that transforms the concentration of protons in a solution into an electrical signal.

Indicators

A substance that changes color in a certain pH range depending on the acidity of the solution; using several indicators you can achieve a fairly accurate result.

Salt

A salt is an ionic compound formed by a cation other than H+ and an anion other than O2-. In a weak aqueous solution, the salts completely dissociate.

To determine the acid-base properties of a salt solution, it is necessary to determine which ions are present in the solution and consider their properties: neutral ions formed from strong acids and bases do not affect pH: they do not release either H + or OH - ions in water. For example, Cl -, NO - 3, SO 2- 4, Li +, Na +, K +.

Anions formed from weak acids exhibit alkaline properties (F -, CH 3 COO -, CO 2- 3); cations with alkaline properties do not exist.

All cations except metals of the first and second groups have acidic properties.

Buffer solution

Solutions that maintain their pH level when a small amount of a strong acid or a strong base is added are mainly composed of:

• A mixture of a weak acid, its corresponding salt and a weak base
• Weak base, corresponding salt and strong acid

To prepare a buffer solution of a certain acidity, it is necessary to mix a weak acid or base with the appropriate salt, taking into account:

• pH range in which the buffer solution will be effective
• Solution capacity - the amount of strong acid or strong base that can be added without affecting the pH of the solution
• There should be no unwanted reactions that could change the composition of the solution

Test:

a) obtaining grounds.

1) The general method for preparing bases is an exchange reaction, with the help of which both insoluble and soluble bases can be obtained:

CuSO 4 + 2 KOH = Cu(OH) 2  + K 2 SO 4,

K 2 CO 3 + Ba(OH) 2 = 2KOH + BaCO 3 .

When soluble bases are obtained by this method, an insoluble salt precipitates.

2) Alkalis can also be obtained by reacting alkali and alkaline earth metals or their oxides with water:

2Li + 2H 2 O = 2LiOH + H 2,

SrO + H 2 O = Sr(OH) 2.

3) Alkalis in technology are usually obtained by electrolysis of aqueous solutions of chlorides:

b)chemicalproperties of bases.

1) The most characteristic reaction of bases is their interaction with acids - the neutralization reaction. Both alkalis and insoluble bases enter into it:

NaOH + HNO 3 = NaNO 3 + H 2 O,

Cu(OH) 2 + H 2 SO 4 = CuSO 4 + 2 H 2 O.

2) It was shown above how alkalis interact with acidic and amphoteric oxides.

3) When alkalis interact with soluble salts, a new salt and a new base are formed. Such a reaction proceeds to completion only when at least one of the resulting substances precipitates.

FeCl 3 + 3 KOH = Fe(OH) 3  + 3 KCl

4) When heated, most bases, with the exception of alkali metal hydroxides, decompose into the corresponding oxide and water:

2 Fe(OH) 3 = Fe 2 O 3 + 3 H 2 O,

Ca(OH) 2 = CaO + H 2 O.

ACIDS – complex substances whose molecules consist of one or more hydrogen atoms and an acid residue. The composition of acids can be expressed by the general formula H x A, where A is the acid residue. Hydrogen atoms in acids can be replaced or exchanged with metal atoms, resulting in the formation of salts.

If an acid contains one such hydrogen atom, then it is a monobasic acid (HCl - hydrochloric, HNO 3 - nitric, HСlO - hypochlorous, CH 3 COOH - acetic); two hydrogen atoms - dibasic acids: H 2 SO 4 - sulfuric, H 2 S - hydrogen sulfide; three hydrogen atoms are tribasic: H 3 PO 4 – orthophosphoric, H 3 AsO 4 – orthoarsenic.

Depending on the composition of the acid residue, acids are divided into oxygen-free (H 2 S, HBr, HI) and oxygen-containing (H 3 PO 4, H 2 SO 3, H 2 CrO 4). In molecules of oxygen-containing acids, hydrogen atoms are connected through oxygen to the central atom: H – O – E. The names of oxygen-free acids are formed from the root of the Russian name for a non-metal, the connecting vowel - O- and the words “hydrogen” (H 2 S – hydrogen sulfide). The names of oxygen-containing acids are given as follows: if the non-metal (less often a metal) included in the acid residue is in the highest degree of oxidation, then suffixes are added to the root of the Russian name of the element -n-, -ev-, or - ov- and then the ending -and I-(H 2 SO 4 - sulfur, H 2 CrO 4 - chrome). If the oxidation state of the central atom is lower, then the suffix is ​​used -ist-(H 2 SO 3 – sulfurous). If a non-metal forms a number of acids, other suffixes are used (HClO - chlorine ovatist aya, HClO 2 – chlorine ist aya, HClO 3 – chlorine ovat aya, HClO 4 – chlorine n and I).

WITH
From the point of view of the theory of electrolytic dissociation, acids are electrolytes that dissociate in an aqueous solution to form only hydrogen ions as cations:

N x A xN + +A x-

The presence of H + ions causes the color change of indicators in acid solutions: litmus (red), methyl orange (pink).

Preparation and properties of acids

A) production of acids.

1) Oxygen-free acids can be obtained by directly combining non-metals with hydrogen and then dissolving the corresponding gases in water:

2) Oxygen-containing acids can often be obtained by reacting acid oxides with water.

3) Both oxygen-free and oxygen-containing acids can be obtained by exchange reactions between salts and other acids:

BaBr 2 + H 2 SO 4 = BaSO 4 + 2 HBr,

CuSO 4 + H 2 S = H 2 SO 4 + CuS ,

FeS+ H 2 SO 4 (dissolved) = H 2 S  + FeSO 4,

NaCl (solid) + H 2 SO 4 (conc.) = HCl  + NaHSO 4,

AgNO 3 + HCl = AgCl + HNO 3,

4) In some cases, redox reactions can be used to produce acids:

3P + 5HNO 3 + 2H 2 O = 3H 3 PO 4 + 5NO 

b ) chemical properties of acids.

1) Acids interact with bases and amphoteric hydroxides. In this case, practically insoluble acids (H 2 SiO 3, H 3 BO 3) can only react with soluble alkalis.

H 2 SiO 3 +2NaOH=Na 2 SiO 3 +2H 2 O

2) The interaction of acids with basic and amphoteric oxides is discussed above.

3) The interaction of acids with salts is an exchange reaction with the formation of salt and water. This reaction proceeds to completion if the reaction product is an insoluble or volatile substance, or a weak electrolyte.

Ni 2 SiO 3 +2HCl=2NaCl+H 2 SiO 3

Na 2 CO 3 +H 2 SO 4 =Na 2 SO 4 +H 2 O+CO 2 

4) The interaction of acids with metals is an oxidation-reduction process. Reductant - metal, oxidizing agent - hydrogen ions (non-oxidizing acids: HCl, HBr, HI, H 2 SO 4 (diluted), H 3 PO 4) or an anion of the acid residue (oxidizing acids: H 2 SO 4 (conc), HNO 3(end and break)). The reaction products of the interaction of non-oxidizing acids with metals in the voltage series up to hydrogen are salt and hydrogen gas:

Zn+H 2 SO 4(dil) =ZnSO 4 +H 2 

Zn+2HCl=ZnCl 2 +H 2 

Oxidizing acids interact with almost all metals, including low-active ones (Cu, Hg, Ag), and the products of reduction of the acid anion, salt and water are formed:

Cu + 2H 2 SO 4 (conc.) = CuSO 4 + SO 2  + 2 H 2 O,

Pb + 4HNO 3(conc) = Pb(NO 3) 2 +2NO 2 + 2H 2 O

AMPHOTERIC HYDROXIDES exhibit acid-base duality: they react with acids as bases:

2Cr(OH) 3 + 3H 2 SO 4 = Cr 2 (SO 4) 3 + 6H 2 O,

and with bases - like acids:

Cr(OH) 3 + NaOH = Na (the reaction takes place in an alkali solution);

Cr(OH) 3 + NaOH = NaCrO 2 + 2H 2 O (the reaction occurs between solid substances during fusion).

Amphoteric hydroxides form salts with strong acids and bases.

Like other insoluble hydroxides, amphoteric hydroxides decompose when heated into oxide and water:

Be(OH) 2 = BeO+H 2 O.

SALT– ionic compounds consisting of metal cations (or ammonium) and anions of acid residues. Any salt can be considered as a product of the reaction of neutralization of a base with an acid. Depending on the ratio of acid and base, salts are obtained: average(ZnSO 4, MgCl 2) – the product of complete neutralization of the base with acid, sour(NaHCO 3, KH 2 PO 4) - with excess acid, basic(CuOHCl, AlOHSO 4) – with an excess of base.

The names of salts according to the international nomenclature are formed from two words: the name of the acid anion in the nominative case and the metal cation in the genitive, indicating the degree of its oxidation, if it is variable, with a Roman numeral in parentheses. For example: Cr 2 (SO 4) 3 – chromium (III) sulfate, AlCl 3 – aluminum chloride. The names of acid salts are formed by adding the word hydro- or dihydro-(depending on the number of hydrogen atoms in the hydroanion): Ca(HCO 3) 2 - calcium bicarbonate, NaH 2 PO 4 - sodium dihydrogen phosphate. The names of the main salts are formed by adding the words hydroxo- or dihydroxo-: (AlOH)Cl 2 – aluminum hydroxychloride, 2 SO 4 – chromium(III) dihydroxosulfate.

Preparation and properties of salts

A ) chemical properties of salts.

1) The interaction of salts with metals is an oxidation-reduction process. In this case, the metal located to the left in the electrochemical series of voltages displaces the subsequent ones from solutions of their salts:

Zn+CuSO 4 =ZnSO 4 +Cu

Alkali and alkaline earth metals do not use for the reduction of other metals from aqueous solutions of their salts, since they interact with water, displacing hydrogen:

2Na+2H 2 O=H 2 +2NaOH.

2) The interaction of salts with acids and alkalis was discussed above.

3) The interaction of salts with each other in solution occurs irreversibly only if one of the products is a slightly soluble substance:

BaCl 2 + Na 2 SO 4 = BaSO 4  + 2NaCl.

4) Hydrolysis of salts - exchange decomposition of some salts with water. The hydrolysis of salts will be discussed in detail in the topic “electrolytic dissociation”.

b) methods of obtaining salts.

In laboratory practice, the following methods for obtaining salts are usually used, based on the chemical properties of various classes of compounds and simple substances:

1) Interaction of metals with non-metals:

Cu+Cl 2 = CuCl 2,

2) Interaction of metals with salt solutions:

Fe+CuCl 2 =FeCl 2 +Cu.

3) Interaction of metals with acids:

Fe+2HCl=FeCl 2 +H 2 .

4) Interaction of acids with bases and amphoteric hydroxides:

3HCl+Al(OH) 3 =AlCl 3 +3H 2 O.

5) Interaction of acids with basic and amphoteric oxides:

2HNO 3 +CuO=Cu(NO 3) 2 +2H 2 O.

6) Interaction of acids with salts:

HCl+AgNO 3 =AgCl+HNO 3.

7) Interaction of alkalis with salts in solution:

3KOH+FeCl 3 =Fe(OH) 3 +3KCl.

8) Interaction of two salts in solution:

NaCl + AgNO 3 = NaNO 3 + AgCl.

9) Interaction of alkalis with acidic and amphoteric oxides:

Ca(OH) 2 +CO 2 =CaCO 3 +H 2 O.

10) Interaction of oxides of various types with each other:

CaO+CO 2 = CaCO 3.

Salts are found in nature in the form of minerals and rocks, in a dissolved state in the water of oceans and seas.

Bases (hydroxides)– complex substances whose molecules contain one or more hydroxy OH groups. Most often, bases consist of a metal atom and an OH group. For example, NaOH is sodium hydroxide, Ca(OH) 2 is calcium hydroxide, etc.

There is a base - ammonium hydroxide, in which the hydroxy group is attached not to the metal, but to the NH 4 + ion (ammonium cation). Ammonium hydroxide is formed when ammonia is dissolved in water (the reaction of adding water to ammonia):

NH 3 + H 2 O = NH 4 OH (ammonium hydroxide).

The valency of the hydroxy group is 1. The number of hydroxyl groups in the base molecule depends on the valence of the metal and is equal to it. For example, NaOH, LiOH, Al (OH) 3, Ca(OH) 2, Fe(OH) 3, etc.

All reasons - solids that have different colors. Some bases are highly soluble in water (NaOH, KOH, etc.). However, most of them are not soluble in water.

Bases soluble in water are called alkalis. Alkali solutions are “soapy”, slippery to the touch and quite caustic. Alkalies include hydroxides of alkali and alkaline earth metals (KOH, LiOH, RbOH, NaOH, CsOH, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2, etc.). The rest are insoluble.

Insoluble bases- these are amphoteric hydroxides, which act as bases when interacting with acids, and behave like acids with alkalis.

Different bases have different abilities to remove hydroxy groups, so they are divided into strong and weak bases.

Strong bases in aqueous solutions easily give up their hydroxy groups, but weak bases do not.

Chemical properties of bases

The chemical properties of bases are characterized by their relationship to acids, acid anhydrides and salts.

1. Act on indicators. Indicators change color depending on interaction with different chemicals. In neutral solutions they have one color, in acid solutions they have another color. When interacting with bases, they change their color: the methyl orange indicator turns yellow, the litmus indicator turns blue, and phenolphthalein becomes fuchsia.

2. Interact with acid oxides with formation of salt and water:

2NaOH + SiO 2 → Na 2 SiO 3 + H 2 O.

3. React with acids, forming salt and water. The reaction of a base with an acid is called a neutralization reaction, since after its completion the medium becomes neutral:

2KOH + H 2 SO 4 → K 2 SO 4 + 2H 2 O.

4. Reacts with salts forming a new salt and base:

2NaOH + CuSO 4 → Cu(OH) 2 + Na 2 SO 4.

5. When heated, they can decompose into water and the main oxide:

Cu(OH) 2 = CuO + H 2 O.

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