Salts – complex substances that consist of metal atoms and acidic residues. This is the most numerous class of inorganic compounds.

Classification. By composition and properties: medium, acidic, basic, double, mixed, complex

Medium salts are products of complete replacement of hydrogen atoms of a polybasic acid with metal atoms.

Upon dissociation, only metal cations (or NH 4 +) are produced. For example:

Na 2 SO 4 ® 2Na + +SO

CaCl 2 ® Ca 2+ + 2Cl -

Acid salts are products of incomplete replacement of hydrogen atoms of a polybasic acid with metal atoms.

Upon dissociation, they give metal cations (NH 4 +), hydrogen ions and anions of the acid residue, for example:

NaHCO 3 ® Na + + HCO « H + +CO .

Basic salts are products of incomplete replacement of OH groups - the corresponding base with acidic residues.

Upon dissociation, they give metal cations, hydroxyl anions and an acid residue.

Zn(OH)Cl ® + + Cl - « Zn 2+ + OH - + Cl - .

Double salts contain two metal cations and upon dissociation give two cations and one anion.

KAl(SO 4) 2 ® K + + Al 3+ + 2SO

Complex salts contain complex cations or anions.

Br ® + + Br - « Ag + +2 NH 3 + Br -

Na ® Na + + - « Na + + Ag + + 2 CN -

Genetic relationship between different classes of compounds

EXPERIMENTAL PART

Equipment and utensils: rack with test tubes, washing machine, alcohol lamp.

Reagents and materials: red phosphorus, zinc oxide, Zn granules, slaked lime powder Ca(OH) 2, 1 mol/dm 3 solutions of NaOH, ZnSO 4, CuSO 4, AlCl 3, FeCl 3, HСl, H 2 SO 4, universal indicator paper, solution phenolphthalein, methyl orange, distilled water.

Work order

1. Pour zinc oxide into two test tubes; add an acid solution (HCl or H 2 SO 4) to one and an alkali solution (NaOH or KOH) to the other and heat slightly on an alcohol lamp.

Observations: Does zinc oxide dissolve in an acid and alkali solution?

Write equations

Conclusions: 1.What type of oxide does ZnO belong to?

2. What properties do amphoteric oxides have?

Preparation and properties of hydroxides

2.1. Dip the tip of the universal indicator strip into the alkali solution (NaOH or KOH). Compare the resulting color of the indicator strip with the standard color scale.

Observations: Record the pH value of the solution.

2.2. Take four test tubes, pour 1 ml of ZnSO 4 solution into the first, CuSO 4 into the second, AlCl 3 into the third, and FeCl 3 into the fourth. Add 1 ml of NaOH solution to each test tube. Write observations and equations for the reactions occurring.

Observations: Does precipitation occur when alkali is added to a salt solution? Indicate the color of the sediment.

Write equations occurring reactions (in molecular and ionic form).

Conclusions: How can metal hydroxides be prepared?

2.3. Transfer half of the sediments obtained in experiment 2.2 to other test tubes. Treat one part of the sediment with a solution of H 2 SO 4 and the other with a solution of NaOH.

Observations: Does precipitate dissolution occur when alkali and acid are added to precipitates?

Write equations occurring reactions (in molecular and ionic form).

Conclusions: 1. What type of hydroxides are Zn(OH) 2, Al(OH) 3, Cu(OH) 2, Fe(OH) 3?

2. What properties do amphoteric hydroxides have?

Obtaining salts.

3.1. Pour 2 ml of CuSO 4 solution into a test tube and dip a cleaned nail into this solution. (The reaction is slow, changes on the surface of the nail appear after 5-10 minutes).

Observations: Are there any changes to the surface of the nail? What is being deposited?

Write the equation for the redox reaction.

Conclusions: Taking into account the range of metal stresses, indicate the method of obtaining salts.

3.2. Place one zinc granule in a test tube and add HCl solution.

Observations: Is there any gas evolution?

Write the equation

Conclusions: Explain this method obtaining salts?

3.3. Pour some slaked lime powder Ca(OH) 2 into a test tube and add HCl solution.

Observations: Is there gas evolution?

Write the equation the reaction taking place (in molecular and ionic form).

Conclusion: 1. What type of reaction is the interaction between a hydroxide and an acid?

2.What substances are the products of this reaction?

3.5. Pour 1 ml of salt solutions into two test tubes: into the first - copper sulfate, into the second - cobalt chloride. Add to both test tubes drop by drop sodium hydroxide solution until precipitation forms. Then add excess alkali to both test tubes.

Observations: Indicate the changes in the color of precipitation in the reactions.

Write the equation the reaction taking place (in molecular and ionic form).

Conclusion: 1. As a result of what reactions are basic salts formed?

2. How can you convert basic salts to medium salts?

Test tasks:

1. From the listed substances, write down the formulas of salts, bases, acids: Ca(OH) 2, Ca(NO 3) 2, FeCl 3, HCl, H 2 O, ZnS, H 2 SO 4, CuSO 4, KOH
Zn(OH) 2, NH 3, Na 2 CO 3, K 3 PO 4.

2. Indicate the formulas of the oxides corresponding to the listed substances H 2 SO 4, H 3 AsO 3, Bi(OH) 3, H 2 MnO 4, Sn(OH) 2, KOH, H 3 PO 4, H 2 SiO 3, Ge( OH) 4 .

3. Which hydroxides are amphoteric? Write down reaction equations characterizing the amphotericity of aluminum hydroxide and zinc hydroxide.

4. Which of the following compounds will interact in pairs: P 2 O 5 , NaOH, ZnO, AgNO 3 , Na 2 CO 3 , Cr(OH) 3 , H 2 SO 4 . Write down equations for possible reactions.

Laboratory work No. 2 (4 hours)

Subject: Qualitative analysis of cations and anions

Target: master the technique of conducting qualitative and group reactions on cations and anions.

THEORETICAL PART

The main task of qualitative analysis is to establish chemical composition substances found in various objects (biological materials, medicines, food, environmental objects). This paper examines a qualitative analysis inorganic substances, which are electrolytes, i.e., essentially a qualitative analysis of ions. From the entire set of occurring ions, the most important in medical and biological terms were selected: (Fe 3+, Fe 2+, Zn 2+, Ca 2+, Na +, K +, Mg 2+, Cl -, PO, CO, etc. ). Many of these ions are part of various medicines and food products.

In qualitative analysis, not all possible reactions are used, but only those that are accompanied by a clear analytical effect. The most common analytical effects: the appearance of a new color, the release of gas, the formation of a precipitate.

There are two fundamentally different approaches to qualitative analysis: fractional and systematic . In systematic analysis, group reagents are necessarily used to separate the ions present into separate groups, and in some cases into subgroups. To do this, some of the ions are transferred into insoluble compounds, and some of the ions are left in solution. After separating the precipitate from the solution, they are analyzed separately.

For example, the solution contains A1 3+, Fe 3+ and Ni 2+ ions. If this solution is exposed to excess alkali, a precipitate of Fe(OH) 3 and Ni(OH) 2 precipitates, and [A1(OH) 4 ] - ions remain in the solution. The precipitate containing iron and nickel hydroxides will partially dissolve when treated with ammonia due to the transition to 2+ solution. Thus, using two reagents - alkali and ammonia, two solutions were obtained: one contained [A1(OH) 4 ] - ions, the other contained 2+ ions and a Fe(OH) 3 precipitate. Using characteristic reactions, the presence of certain ions is then proven in solutions and in the precipitate, which must first be dissolved.

Systematic analysis is used mainly for the detection of ions in complex multicomponent mixtures. It is very labor-intensive, but its advantage lies in the easy formalization of all actions that fit into a clear scheme (methodology).

To carry out fractional analysis, only characteristic reactions are used. Obviously, the presence of other ions can significantly distort the results of the reaction (overlapping colors, unwanted precipitation, etc.). To avoid this, fractional analysis mainly uses highly specific reactions that give an analytical effect with a small number of ions. For successful reactions, it is very important to maintain certain conditions, in particular pH. Very often in fractional analysis it is necessary to resort to masking, that is, to convert ions into compounds that are not capable of producing an analytical effect with the selected reagent. For example, dimethylglyoxime is used to detect nickel ion. The Fe 2+ ion gives a similar analytical effect to this reagent. To detect Ni 2+, the Fe 2+ ion is transferred to a stable fluoride complex 4- or oxidized to Fe 3+, for example, with hydrogen peroxide.

Fractional analysis is used to detect ions in simpler mixtures. Analysis time is significantly reduced, but at the same time the experimenter is required to have a deeper knowledge of the flow patterns chemical reactions, since it is quite difficult to take into account in one specific technique all possible cases of mutual influence of ions on the nature of the observed analytical effects.

In analytical practice, the so-called fractional-systematic method. With this approach, a minimum number of group reagents is used, which makes it possible to outline analysis tactics in general outline, which is then carried out using the fractional method.

According to the technique of conducting analytical reactions, reactions are distinguished: sedimentary; microcrystalscopic; accompanied by the release of gaseous products; conducted on paper; extraction; colored in solutions; flame coloring.

When carrying out sedimentary reactions, the color and nature of the precipitate (crystalline, amorphous) must be noted; if necessary, additional tests are carried out: the precipitate is checked for solubility in strong and weak acids, alkalis and ammonia, excess reagent. When carrying out reactions accompanied by the release of gas, its color and smell are noted. In some cases, additional tests are carried out.

For example, if the gas released is suspected to be carbon monoxide (IV), it is passed through an excess of lime water.

In fractional and systematic analyses, reactions during which a new color appears are widely used, most often these are complexation reactions or redox reactions.

In some cases, it is convenient to carry out such reactions on paper (droplet reactions). Reagents that do not decompose under normal conditions are applied to paper in advance. Thus, to detect hydrogen sulfide or sulfide ions, paper impregnated with lead nitrate is used [blackening occurs due to the formation of lead(II) sulfide]. Many oxidizing agents are detected using iodine starch paper, i.e. paper soaked in solutions of potassium iodide and starch. In most cases, the necessary reagents are applied to paper during the reaction, for example, alizarin for the A1 3+ ion, cupron for the Cu 2+ ion, etc. To enhance the color, extraction into an organic solvent is sometimes used. For preliminary tests, flame color reactions are used.

Let me briefly remind you of specific examples, how should salts be called correctly?

Example 1. The salt K 2 SO 4 is formed by a sulfuric acid residue (SO 4) and metal K. Salts of sulfuric acid are called sulfates. K 2 SO 4 - potassium sulfate.

Example 2. FeCl 3 - the salt contains iron and a hydrochloric acid residue (Cl). Name of salt: iron (III) chloride. Please note: in in this case we must not only name the metal, but also indicate its valency (III). In the previous example, this was not necessary, since the valency of sodium is constant.

Important: the name of the salt should indicate the valency of the metal only if the metal has a variable valence!

Example 3. Ba(ClO) 2 - the salt contains barium and the remainder of hypochlorous acid (ClO). Salt name: barium hypochlorite. The valency of the metal Ba in all its compounds is two; it does not need to be indicated.

Example 4. (NH 4) 2 Cr 2 O 7. The NH 4 group is called ammonium, the valence of this group is constant. Name of salt: ammonium dichromate (dichromate).

In the above examples we only encountered the so-called. medium or normal salts. Acidic, basic, double and complex salts, salts of organic acids will not be discussed here.

If you are interested not only in the nomenclature of salts, but also in the methods of their preparation and Chemical properties, I recommend turning to the relevant sections of the chemistry reference book: "

Salts – substances that can be considered as the product of the replacement of hydrogen atoms in an acid with metal atoms or a group of atoms. There are 5 types of salts: medium (normal), acidic, basic, double, complex, differing in the nature of the ions formed during dissociation.

1.Medium salts are products of complete replacement of hydrogen atoms in the molecule acids. Salt composition: cation - metal ion, anion - acid residue ion. Na 2 CO 3 - sodium carbonate

Na 3 PO 4 - sodium phosphate

Na 3 PO 4 = 3Na + + PO 4 3-

cation anion

2.Sour salts – products of incomplete replacement of hydrogen atoms in an acid molecule. The anion contains hydrogen atoms.

NaH 2 PO 4 =Na + + H 2 PO 4 -

Dihydrogen phosphate cation anion

Acidic salts produce only polybasic acids when the amount of base taken is insufficient.

H 2 SO 4 +NaOH=NaHSO 4 +H 2 O

hydrogen sulfate

By adding excess alkali, the acidic salt can be converted to medium

NaHSO 4 +NaOH=Na 2 SO 4 +H 2 O

3.Basic salts – products of incomplete substitution of hydroxide ions based on acid residue. The cation contains a hydroxo group.

CuOHCl=CuOH + +Cl -

hydroxochloride cation anion

Basic salts can only be formed by polyacid bases

(bases containing several hydroxyl groups), when they interact with acids.

Cu(OH) 2 +HCl=CuOHCl+H2O

You can convert a basic salt to a medium salt by treating it with an acid:

CuOHCl+HCl=CuCl 2 +H 2 O

4.Double salts – they contain cations of several metals and anions of one acid

KAl(SO 4) 2 = K + + Al 3+ + 2SO 4 2-

potassium aluminum sulfate

Characteristic properties All types of salts considered are: exchange reactions with acids, alkalis and with each other.

For naming salts use Russian and international nomenclature.

The Russian name of the salt is made up of the name of the acid and the name of the metal: CaCO 3 - calcium carbonate.

For acidic salts, the “sour” additive is introduced: Ca(HCO 3) 2 - acidic calcium carbonate. To name the main salts, add “basic”: (СuOH) 2 SO 4 – basic copper sulfate.

The most widespread is the international nomenclature. The name of the salt according to this nomenclature consists of the name of the anion and the name of the cation: KNO 3 - potassium nitrate. If the metal has a different valence in the compound, then it is indicated in brackets: FeSO 4 - iron sulfate (III).

For salts of oxygen-containing acids, the suffix “at” is added to the name if the acid-forming element has a higher valency: KNO 3 – potassium nitrate; suffix “it” if the acid-forming element has a lower valency: KNO 2 - potassium nitrite. In cases where an acid-forming element forms acids in more than two valence states, the suffix “at” is always used. Moreover, if it exhibits a higher valency, the prefix “per” is added. For example: KClO 4 – potassium perchlorate. If the acid-forming element forms a lower valency, the suffix “it” is used, with the addition of the prefix “hypo”. For example: KClO – potassium hypochlorite. For salts formed by acids containing different amounts of water, the prefixes “meta” and “ortho” are added. For example: NaPO 3 - sodium metaphosphate (salt of metaphosphoric acid), Na 3 PO 4 - sodium orthophosphate (salt of orthophosphoric acid). The prefix “hydro” is introduced into the name of the acidic salt. For example: Na 2 HPO 4 – sodium hydrogen phosphate (if the anion has one hydrogen atom) and the prefix “hydro” with the Greek numeral (if there is more than one hydrogen atom) – NaH 2 PO 4 – sodium dihydrogen phosphate. The prefix “hydroxo” is introduced into the names of the main salts. For example: FeOHCl – iron hydroxychloride (I).

5. Complex salts – compounds that form complex ions (charged complexes) upon dissociation. When writing complex ions, it is customary to enclose them in square brackets. For example:

Ag(NH 3) 2  Cl = Ag(NH 3) 2  + + Cl -

K 2 PtCl 6  = 2K + + PtCl 6  2-

According to the ideas proposed by A. Werner, in a complex connection there are internal and external spheres. So, for example, in the complex compounds considered, the inner sphere is composed of complex ions Ag(NH 3) 2  + and PtCl 6  2-, and the outer sphere is Cl - and K +, respectively. The central atom or ion of the inner sphere is called a complexing agent. In the proposed compounds these are Ag +1 and Pt +4. Molecules or ions of opposite sign coordinated around a complexing agent are ligands. In the compounds under consideration, these are 2NH 3 0 and 6Cl -. The number of ligands of a complex ion determines its coordination number. In the proposed compounds it is equal to 2 and 6, respectively.

Complexes are distinguished by the sign of the electric charge

1.Cationic (coordination around the positive ion of neutral molecules):

Zn +2 (NH 3 0) 4 Cl 2 -1 ; Al +3 (H 2 O 0) 6  Cl 3 -1

2.Anionic (coordination around a complexing agent in a positive oxidation state of a ligand having negative degree oxidation):

K 2 +1 Be +2 F 4 -1 ; K 3 +1 Fe +3 (CN -1) 6 

3. Neutral complexes – complex compounds without an outer spherePt + (NH 3 0) 2 Cl 2 -  0. Unlike compounds with anionic and cationic complexes, neutral complexes are not electrolytes.

Dissociation of complex compounds into inner and outer spheres is called primary . It proceeds almost entirely like strong electrolytes.

Zn (NH 3) 4 Cl 2 → Zn (NH 3) 4  +2 + 2Cl ─

K 3 Fe(CN) 6 → 3 K + +Fe(CN) 6  3 ─

Complex ion (charged complex) in a complex compound forms the inner coordination sphere, the remaining ions form the outer sphere.

In a complex compound K 3, the complex ion 3-, consisting of a complexing agent - the Fe 3+ ion and ligands - CN ─ ions, is the inner sphere of the compound, and the K + ions form the outer sphere.

The ligands located in the inner sphere of the complex are bound by the complexing agent much more tightly and their elimination during dissociation occurs only to a small extent. The reversible dissociation of the inner sphere of a complex compound is called secondary .

Fe(CN) 6  3 ─ Fe 3+ + 6CN ─

Secondary dissociation of the complex proceeds according to the type of weak electrolytes. The algebraic sum of the charges of particles formed during the dissociation of a complex ion is equal to the charge of the complex.

The names of complex compounds, as well as the names of ordinary substances, are formed from the Russian names of cations and the Latin names of anions; just as in ordinary substances, in complex compounds the first is called the anion. If the anion is complex, its name is formed from the name of the ligands with the ending “o” (Cl - - chloro, OH - - hydroxo, etc.) and the Latin name of the complexing agent with the suffix “at”; the number of ligands is, as usual, indicated by the corresponding numeral. If the complexing agent is an element capable of exhibiting a variable oxidation state, the numerical value of the oxidation state, as in the names of ordinary compounds, is indicated by a Roman numeral in parentheses

Example: Names of complex compounds with a complex anion.

K 3 – potassium hexacyanoferrate (III)

Complex cations in the vast majority of cases contain neutral water molecules H 2 O, called “aqua,” or ammonia NH 3, called “ammine,” as ligands. In the first case, complex cations are called aqua complexes, in the second - ammonia. The name of the complex cation consists of the name of the ligands indicating their number and the Russian name of the complexing agent with the indicated value of its oxidation state, if necessary.

Example: Names of complex compounds with a complex cation.

Cl 2 – tetrammine zinc chloride

Complexes, despite their stability, can be destroyed in reactions in which ligands are bound into even more stable weakly dissociating compounds.

Example: Destruction of a hydroxo complex by an acid due to the formation of weakly dissociating H 2 O molecules.

K 2 + 2H 2 SO 4 = K 2 SO 4 + ZnSO 4 + 2H 2 O.

Name of the complex compound they begin by indicating the composition of the inner sphere, then name the central atom and its oxidation state.

In the inner sphere, anions are first named, adding the ending “o” to the Latin name.

F -1 – fluoro Cl - - chloroCN - - cyanoSO 2 -2 –sulfito

OH - - hydroxoNO 2 - - nitrito, etc.

Then the neutral ligands are called:

NH 3 – ammin H 2 O – aqua

The number of ligands is marked with Greek numerals:

I – mono (usually not indicated), 2 – di, 3 – three, 4 – tetra, 5 – penta, 6 – hexa. Next we move on to the name of the centralatom (complexing agent). The following are taken into account:

If the complexing agent is part of the cation, then the Russian name of the element is used and the degree of its oxidation is indicated in parentheses in Roman numerals;

If the complexing agent is part of an anion, then the Latin name of the element is used, its oxidation state is indicated before it, and the ending “at” is added at the end.

After the designation of the inner sphere, cations or anions located in the outer sphere are indicated.

When forming the name of a complex compound, one must remember that the ligands included in its composition can be mixed: electrically neutral molecules and charged ions; or charged ions of different types.

Ag +1 NH 3  2 Cl– diamine silver (I) chloride

K 3 Fe +3 CN 6 - hexacyano (III) potassium ferrate

NH 4  2 Pt +4 OH 2 Cl 4 – dihydroxotetrachloro(IV) ammonium platinate

Pt +2 NH 3  2 Cl 2 -1  o - diammine dichloride-platinum x)

X) in neutral complexes the name of the complexing agent is given in the nominative case

These are substances that dissociate in solutions to form hydrogen ions.

Acids are classified by their strength, by their basicity, and by the presence or absence of oxygen in the acid.

By strengthacids are divided into strong and weak. The most important strong acids- nitrogen HNO 3, sulfuric H2SO4, and hydrochloric HCl.

According to the presence of oxygen distinguish between oxygen-containing acids ( HNO3, H3PO4 etc.) and oxygen-free acids ( HCl, H 2 S, HCN, etc.).

By basicity, i.e. According to the number of hydrogen atoms in an acid molecule that can be replaced by metal atoms to form a salt, acids are divided into monobasic (for example, HNO 3, HCl), dibasic (H 2 S, H 2 SO 4), tribasic (H 3 PO 4), etc.

The names of oxygen-free acids are derived from the name of the non-metal with the addition of the ending -hydrogen: HCl - hydrochloric acid, H2S e - hydroselenic acid, HCN - hydrocyanic acid.

The names of oxygen-containing acids are also formed from the Russian name of the corresponding element with the addition of the word “acid”. In this case, the name of the acid in which the element is in the highest oxidation state ends in “naya” or “ova”, for example, H2SO4 - sulfuric acid, HClO4 - perchloric acid, H3AsO4 - arsenic acid. With a decrease in the oxidation degree of the acid-forming element, the endings change in the following sequence: “ovate” ( HClO3 - perchloric acid), “solid” ( HClO2 - chlorous acid), “ovate” ( H O Cl - hypochlorous acid). If an element forms acids while being in only two oxidation states, then the name of the acid corresponding to the lowest oxidation state of the element receives the ending “iste” ( HNO3 - Nitric acid, HNO2 - nitrous acid).

Table - The most important acids and their salts

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. Oxygen-containing acids can often be obtained by directly combining acid oxides with water:

SO 3 + H 2 O = H 2 SO 4,

CO 2 + H 2 O = H 2 CO 3,

P 2 O 5 + H 2 O = 2 HPO 3.

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 + 2HBr,

CuSO 4 + H 2 S = H 2 SO 4 + CuS,

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

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

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

3P + 5HNO3 + 2H2O = 3H3PO4 + 5NO.

Chemical properties of acids

1. The most characteristic chemical property of acids is their ability to react with bases (as well as with basic and amphoteric oxides) to form salts, for example:

H 2 SO 4 + 2NaOH = Na 2 SO 4 + 2H 2 O,

2HNO 3 + FeO = Fe(NO 3) 2 + H 2 O,

2 HCl + ZnO = ZnCl 2 + H 2 O.

2. The ability to interact with some metals in the voltage series up to hydrogen, with the release of hydrogen:

Zn + 2HCl = ZnCl 2 + H 2,

2Al + 6HCl = 2AlCl3 + 3H2.

3. With salts, if a slightly soluble salt or volatile substance is formed:

H 2 SO 4 + BaCl 2 = BaSO 4 ↓ + 2HCl,

2HCl + Na 2 CO 3 = 2NaCl + H 2 O + CO 2,

2KHCO 3 + H 2 SO 4 = K 2 SO 4 +2SO 2+ 2H 2 O.

Note that polybasic acids dissociate stepwise, and the ease of dissociation at each step decreases; therefore, for polybasic acids, instead of medium salts, acidic salts are often formed (in the case of an excess of the reacting acid):

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.

4. A special case of acid-base interaction is the reaction of acids with indicators, leading to a change in color, which has long been used for the qualitative detection of acids in solutions. So, litmus changes color in an acidic environment to red.

5. When heated, oxygen-containing acids decompose into oxide and water (preferably in the presence of a water-removing agent P2O5):

H 2 SO 4 = H 2 O + SO 3,

H 2 SiO 3 = H 2 O + SiO 2.

M.V. Andryukhova, L.N. Borodina