Complex compounds
Summary of the lesson-lecture
Goals. Form an idea of the composition, structure, properties and nomenclature of complex compounds; develop skills for determining the oxidation state of a complexing agent, drawing up equations for the dissociation of complex compounds.
New concepts: complex compound, complexing agent, ligand, coordination number, external and internal spheres of the complex.
Equipment and reagents. A rack with test tubes, concentrated ammonia solution, solutions of copper (II) sulfate, silver nitrate, sodium hydroxide.
DURING THE CLASSES
Laboratory experience. Add ammonia solution to the copper (II) sulfate solution. The liquid will turn intense blue.
What happened? Chemical reaction? Until now, we did not know that ammonia can react with salt. What substance was formed? What is its formula, structure, name? To what class of compounds can it be attributed? Can ammonia react with other salts? Are there any connections similar to this? We have to answer these questions today.
To better study the properties of some compounds of iron, copper, silver, aluminum, we need knowledge about complex compounds.
Let's continue our experience. Divide the resulting solution into two parts. Add alkali to one part. The precipitation of copper (II) hydroxide Cu (OH) 2 is not observed, therefore, there are no doubly charged copper ions in the solution or there are too few of them. Hence, we can conclude that copper ions interact with the added ammonia and form some new ions that do not give an insoluble compound with OH - ions.
At the same time, the ions remain unchanged. This can be verified by adding a solution of barium chloride to the ammonia solution. A white precipitate of BaSO 4 will precipitate immediately.
Studies have established that the dark blue color of the ammonia solution is due to the presence of complex 2+ ions in it, formed by the addition of four ammonia molecules to the copper ion. When water evaporates, ions 2+ bind with ions, and dark blue crystals are released from the solution, the composition of which is expressed by the formula SO 4 H 2 O.
Complex compounds are those containing complex ions and molecules capable of existing both in crystalline form and in solutions.
Formulas of molecules or ions of complex compounds are usually enclosed in square brackets. Complex compounds are obtained from conventional (non-complex) compounds.
Examples of obtaining complex compounds
The structure of complex compounds is considered on the basis of the coordination theory proposed in 1893 by the Swiss chemist Alfred Werner, a Nobel Prize winner. His scientific activities took place at the University of Zurich. The scientist synthesized many new complex compounds, systematized previously known and newly obtained complex compounds and developed experimental methods for proving their structure.
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A. Werner
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In accordance with this theory, complex compounds are distinguished complexing agent, external and inner sphere... The complexing agent is usually a cation or neutral atom. The inner sphere is made up of a certain number of ions or neutral molecules, which are firmly bound to the complexing agent. They are called ligands... The number of ligands determines coordination number(CN) complexing agent.
Complex compound example
The compound SO 4 H 2 O or CuSO 4 5H 2 O considered in the example is a crystalline hydrate of copper (II) sulfate.
Let us determine the constituent parts of other complex compounds, for example, K 4.
(Reference. The substance with the formula HCN is hydrocyanic acid. Hydrocyanic acid salts are called cyanides.)
Complexing agent - iron ion Fe 2+, ligands - cyanide ions СN -, coordination number is equal to six. Everything in square brackets is an inner sphere. Potassium ions form the outer sphere of the complex compound.
The nature of the bond between the central ion (atom) and ligands can be twofold. On the one hand, the bond is due to the forces of electrostatic attraction. On the other hand, between the central atom and ligands a bond can be formed by the donor-acceptor mechanism by analogy with the ammonium ion. In many complex compounds, the bond between the central ion (atom) and the ligands is due to both the forces of electrostatic attraction and the bond formed due to the lone electron pairs of the complexing agent and free orbitals of the ligands.
Complex compounds with an outer sphere are strong electrolytes and in aqueous solutions dissociate almost entirely into a complex ion and ions outer sphere. For example:
SO 4 2+ +.
In exchange reactions, complex ions pass from one compound to another without changing their composition:
SO 4 + BaCl 2 = Cl 2 + BaSO 4.
The inner sphere can have a positive, negative, or zero charge.
If the charge of the ligands compensates for the charge of the complexing agent, then such complex compounds are called neutral or non-electrolyte complexes: they consist only of the complexing agent and ligands of the inner sphere.
Such a neutral complex is, for example,.
The most typical complexing agents are cations d-elements.
Ligands can be:
a) polar molecules - NH 3, H 2 O, CO, NO;
b) simple ions - F -, Cl -, Br -, I -, H -, H +;
c) complex ions - CN -, SCN -, NO 2 -, OH -.
Consider the table showing the coordination numbers of some complexing agents.
Nomenclature of complex compounds. In a compound, the anion is first named and then the cation. When specifying the composition of the inner sphere, anions are first of all called, adding the suffix - O-, for example: Cl - - chloro, CN - - cyano, OH - - hydroxo, etc. Hereinafter, neutral ligands are called and primarily ammonia and its derivatives. In this case, they use the terms: for coordinated ammonia - ammin, for water - aqua... The number of ligands is indicated in Greek words: 1 - mono, 2 - di, 3 - three, 4 - tetra, 5 - penta, 6 - hexa. Then they move on to the name of the central atom. If the central atom is part of the cations, then the Russian name of the corresponding element is used and its oxidation state (in Roman numerals) is indicated in brackets. If the central atom is contained in the anion, then use the Latin name of the element, and at the end add the ending - at... In the case of non-electrolytes, the oxidation state of the central atom is not given, since it is unambiguously determined from the condition that the complex is electrically neutral.
Examples. To name the complex Сl 2, the oxidation state is determined
(S.O.)
NS complexing agent - Cu ion NS+ :
1 x + 2 (–1) = 0,x = +2, C.O. (Cu) = +2.
The oxidation state of the cobalt ion is found in a similar way:
y + 2 (–1) + (–1) = 0,y = +3, S.O. (Co) = +3.
What is the coordination number of cobalt in this compound? How many molecules and ions are there around the central ion? The coordination number of cobalt is six.
The name of the complex ion is written in one word. The oxidation state of the central atom is indicated by a Roman numeral in parentheses. For example:
Cl 2 - tetraammine copper (II) chloride,
NO 3 –
dichloroaquatriamminecobalt (III) nitrate,
K 3 - hexacyanoferrate (III)
potassium,
K 2 - tetrachloroplatinate (II)
potassium,
- dichlorotetraamminezinc,
H 2 - hexachloro tin acid.
Using the example of several complex compounds, we will determine the structure of molecules (ion-complexing agent, its SO, coordination number, ligands, inner and outer spheres), give a name to the complex, write down the equations of electrolytic dissociation.
K 4 - potassium hexacyanoferrate (II),
K 4 4K + + 4–.
H - tetrachloroauric acid (formed by dissolving gold in aqua regia),
H H + + -.
OH - diammine silver (I) hydroxide (this substance participates in the "silver mirror" reaction),
OH + + OH -.
Na - tetrahydroxoaluminate sodium,
Na Na + + -.
Complex compounds also include many organic substances, in particular, the products of the interaction of amines with water and acids known to you. For example, methyl ammonium chloride salts and phenylammonium chloride are complex compounds. According to the coordination theory, they have the following structure:
Here, the nitrogen atom is a complexing agent, hydrogen atoms at nitrogen, methyl and phenyl radicals are ligands. Together they form the inner sphere. The outer sphere contains chloride ions.
Many organic substances, which are of great importance in the life of organisms, are complex compounds. These include hemoglobin, chlorophyll, enzymes and dr.
Complex compounds are widely used:
1) in analytical chemistry for the determination of many ions;
2) to separate some metals and obtain metals of high purity;
3) as dyes;
4) to eliminate water hardness;
5) as catalysts for important biochemical processes.
Complex metal compounds
Metals in living systems, as a rule, exist in the composition of various complex compounds with bioligands. Therefore, this most important property of metals - their ability to form a variety of complex structures - will be considered primarily on individual examples.
1. Aqua complexes
In aqueous solutions, d-metal cations in free form (including in the body) exist in the form of n + aquocomplexes, which are usually denoted as Me n + or Me n + hydr. "Aquocomplexes of some metals, in particular copper (II) , manganese (II), silver (1), are quite stable, so the salts of these metals do not undergo hydrolysis.
2. Ammonia
Ammonia complexes are good models for understanding the structures associated with the formation of biological compounds containing amino groups.
A. Formation of copper (II) ammonia.
2+ (blue) + 4NH 3 2+ (blue) + 4H 2 0
In molecular form, this process can be represented as follows:
SO 4 + 4NH 3 S0 4 + 4H 2 O
And simplified, without being reflected in the record of the formation of the aqua complex, the equation will take the form:
In what follows, when writing reactions in ionic or molecular form, we will write down metal ions in a simplified manner Me n +, meaning hydrated ions.
CuSO 4 + 4NH 3 SO 4
An important aspect of the behavior of "biocomplexes", i.e. complexes in living systems is their stability. Therefore, it is important to know the factors affecting the stability of complex systems and the possible ways of their destruction.
The reason for the destruction of the complex may be the removal of the complexing agent (Cu 2+) from the inner sphere of the complex and its binding in the form of a poorly soluble compound (CuS in the first reaction) or the removal of ligands (NH3) and their binding to a more stable compound (NH 4 + ion in the second reaction ).
B. Dissolution of silver chloride in a solution of excess ammonia to form silver ammonia.
AgCl + 2NH 3 (excess) -> Cl (colorless)
This complex can also be destroyed in several ways.
B. The interaction of zinc and cadmium salts with ammonia also leads to the formation of ammonia complexes.
D. The reaction of mercury (II) chloride (mercuric chloride) with ammonia ends with the formation of a white precipitate - aminorturate chloride (white precipitate - antiseptic), which is not a complex compound.
HgCl 2 + 2NH 3 -> Cl-Hg-NH 2 + NH 4 C1
3. Chelated complexes with amino acids
Many metalloenzymes, in which metal ions bind to a protein through the oxygen of the carboxyl groups and the nitrogen of the amino groups, are bioclusters (protein complexes) - stable chelate compounds.
The process of interaction of aquocomplexes of biometals with amino acids, leading to such structures, is accompanied by a sharp increase in the entropy of the system (AS> 0) due to a significant increase in the number of particles (entropy effect). For example, in the case of copper (II) ions and glycine 1:
Chelated (entropic) effect - an increase in entropy and the formation of five- and six-membered rings are the reason for the relatively higher stability of chelate compounds in comparison with analogous metal complexes with monodentate ligands or with chelating reagents, but with a smaller number of chelate cycles.
Note that the toxicity of copper compounds is due not only to the binding of thiol (see above), but also amino groups of proteins, which leads to disruption of enzymatic activity, and, consequently, normal life.
4. Chelate complexes with ethylenediaminetetraacetic acid (EDTA). trilon B (Na 2 EDTA). pentacin- complexones used in the widespread method of chelation therapy.
In this diagram, Trilon B is shown as a tetradentate ligand, but it should be borne in mind that this complex is capable of forming six bonds with the complexing agent, and it is more correct to write the final product in a different form.
5. Macrocyclic complexes
Many bioactive compounds are based on complexes based on macroheterocycles. Examples of such structures are discussed below. A. Porphyrin cycle.
Chlorophylls (a, b): Me = Mg 2+, X and Y are absent.
Heme proteins(hemoglobin, myoglobin, cytochromes, enzymes - catalase, peroxidase): Me = Fe 2+ (Fe 3+); X is H 2 O, O 2, CO, CN -; Y - organic residue.
B. Corrine cycle(similar to porphyrin, differing in several details).
Vitamin B 12 (growth factor, hematopoietic stimulator): Me = Co 3+, X = CN, Y - organic residue.
V. Membrane / active complexes.
Among natural complex compounds, a special place is occupied by macrocomplexes based on cyclic polypeptides containing internal cavities of certain sizes, in which there are several oxygen-containing groups capable of binding the cations of those metals, the sizes of which correspond to the size of the cavity. Such structures, being in biological membranes, provide the transport of ions across membranes and are therefore called ionophores.
Natural ionophores that perform ion transport functions are antibiotics: valinomycin and nonactin.
Crown ethers and cryptands are models of natural ionophores. The first of them selectively interact with alkali metals, the second - with alkaline earth metals.
The simplest crown ethers have the general formula (CH 2 CH 2 O) n.
The stability of complexes with crown ethers is related to the size of metal ions and the size of the cycle. Li + binds more strongly to crown-4 (the number "4" indicates the number of oxygen atoms contained in the cycle of the crown ether molecule), Na + - to crown-5, K + - with crown-6, Cs + - with crown-8.
Cryptands - macrobicyclic ligands - most effectively bind ions of alkaline earth metals, they can even dissolve barium sulfate.
6. Complex compounds underlying qualitative reactions to ionsFe 2+ . Fe 3+ . Co 2+ . Ni 2+ . Hg 2+
A qualitative reaction to the Fe 2- ion is the interaction with potassium hex-cyanoferrate (III) (red blood salt). In this case, a blue precipitate is formed - potassium-iron (II) hexacyanoferrate (III) (turnboolean blue).
FeSO 4 (II) + K 3 (III) -> KFe (III) (blue) + K 2 SO 4
Qualitative reactions to Fe 3+ ion are:
Interaction with potassium hexacyanoferrate (II) (yellow blood salt).
This produces a blue precipitate. - potassium iron (II) hexacyanoferrate (III) (Prussian blue).
FeCl 3 + K 4 -> KFe + ZKS1
It should be noted that in this case, unlike the previous one, there is a redox process, in which iron (III) chloride acts as an oxidizing agent, since its redox potential [ф ° (Fe 3+ / Fe 2+) = + 0.77 V] is greater than the redox potential of the complex ion - hexacyanoferrate (II) (f ° 3- / 4 ~ = + 0.36 V), which is a reducing agent. Thus, the sediments of Turnbull blue and Prussian blue are identical not only in color, but also in chemical structure.
Interaction with potassium thiocyanate.
In this case, a red complex is formed - triaquotrithio-cyanato-iron (III).
3+ (yellow) + 3SCN - (red) + ЗН 2 О
A qualitative reaction to the Co 2+ ion is the interaction with ammonium thiocyanide, thus forming blue ammonium tetraisothiocyanatocobaltate (II), which is stable only in an organic solvent, for example, in amyl alcohol.
[Co (H 2 O) 4] 2+ + 4NCS - 2- (blue) + 4H 2 O
The qualitative reaction to the Ni 2+ ion is Chugaev's reaction - interaction with dimethylglyoxime, thus forming a chelate compound of bright red color - nickel dimethylglyoxime. The reaction is carried out in an ammonia solution. It is very sensitive, used in. toxicology and forensic medicine for the detection of nickel.
A qualitative reaction to the mercury (II) ion is its interaction with a solution of potassium iodide. First, an orange precipitate of mercury iodide (II) precipitates, which dissolves in an excess of potassium iodide to form a colorless complex compound - potassium tetraiodomercurate (II).
HgCl 2 + 2KI -> HgI 2 + 2KC1 HgI 2 + 2K1 (from 6 tests) -> K 2
A solution of this salt in a concentrated solution of caustic alkali is known as Nessler's reagent and is used as a sensitive reagent for ammonium and ammonia ions.
SO 4
Purpose: to obtain a complex copper sulfate-tetroamino salt from copper sulfate CuSO 4 ∙ 5H 2 O and a concentrated ammonia solution NH 4 OH.
Safety precautions:
1. Glass chemical glassware should be handled with care and should be checked for cracks before starting work.
2. Before starting work, you should check the serviceability of electrical appliances.
3. Heat only in heat-resistant dishes.
4. Carefully and economically use chemicals. reagents. Do not taste or smell them.
5. Work should be carried out in gowns.
6. Ammonia is poisonous and its vapors irritate the mucous membrane.
Reagents and equipment:
Concentrated ammonia solution - NH 4 OH
Ethyl alcohol - C 2 H 5 OH
Copper sulfate - CuSO 4 ∙ 5H 2 O
Distilled water
Graduated cylinders
Petri dishes
Vacuum pump (water jet vacuum pump)
Glass funnels
Theoretical justification:
Complex compounds are a substance containing a complexing agent, with which a certain number of ions or molecules are associated, called addends or legends. The complexing agent with the addends constitutes the inner sphere of the complex compound. In the outer sphere of complex compounds, there is an ion associated with a complex ion.
Complex compounds are obtained by the interaction of substances with simpler composition. In aqueous solutions, they dissociate to form a positively or negatively charged complex ion and the corresponding anion or cation.
SO 4 = 2+ + SO 4 2-
2+ = Cu 2+ + 4NH 3 -
Complex 2+ stains the solution in a cornflower blue color, while Cu2 + and 4NH3 taken separately do not give such staining. Complex compounds are of great importance in applied chemistry.
SO4 - dark purple crystals, soluble in water, but insoluble in alcohol. When heated to 1200C, it loses water and a part of ammonia, and at 2600C it loses all ammonia. When stored in air, salt decomposes.
Synthesis equation:
CuSO4 ∙ 5H2O + 4NH4OH = SO4 ∙ H2O + 8H2O
CuSO4 ∙ 5H2O + 4NH4OH = SO4 ∙ H2O + 8H2O
Mm CuSO4 ∙ 5H2O = 250 g / mol
Mm SO4 ∙ H2O = 246 g / mol
6g CuSO4 ∙ 5H2O - Xg
250 g CuSO4 ∙ 5H2O - 246 SO4 ∙ H2O
X = 246 ∙ 6/250 = 5.9 g SO4 ∙ H2O
Progress:
Dissolve 6 g of copper sulfate in 10 ml of distilled water in a heat-resistant glass. Heat the solution. Stir vigorously until complete dissolution, then add concentrated ammonia solution in small portions until a purple complex salt solution appears.
Then transfer the solution to a Petri dish or a porcelain dish and precipitate crystals of the complex salt with ethyl alcohol, which is poured in with a burette for 30-40 minutes, the volume of ethyl alcohol is 5-8 ml.
Filter the obtained crystals of complex salt on a Buchner funnel and leave to dry until the next day. Then weigh the crystals and calculate the% yield.
5.9g SO4 ∙ H2O - 100%
m sample - X
X = m sample ∙ 100% / 5.9g
Control questions:
1.What is the type of chemical bonds in complex salts?
2. What is the mechanism for the formation of a complex ion?
3.How to determine the charge of a complexing agent and a complex ion?
4.How does complex salt dissociate?
5. Make the formulas of complex compounds of dicyano - sodium argentate.
Laboratory work No. 6
Obtaining orthoboric acid
Target: to obtain orthoboric acid from borax and hydrochloric acid.
Safety precautions:
1. Glass chemical glassware should be handled with care and should be checked for cracks before use.
2. Before starting work, check the serviceability of electrical appliances.
3. Heat only in heat-resistant dishes.
4. Use chemicals carefully and sparingly. Do not taste or smell them.
5. Work should be carried out in gowns.
Equipment and reagents:
Sodium tetraborate (decahydrate) - Na 2 B 4 O 7 * 10H 2 O
Hydrochloric acid (conc.) - HCl
Distilled water
Hot plate, vacuum pump (water jet vacuum pump), beakers, filter paper, porcelain cups, glass rods, glass funnels.
Progress:
Dissolve 5 g of sodium tetraborate decahydrate in 12.5 ml of boiling water, add 6 ml of hydrochloric acid solution and leave to stand for a day.
Na 2 B 4 O 7 * 10H 2 O + 2HCl + 5H 2 O = 4H 3 BO 3 + 2NaCl
The formed precipitate of orthoboric acid is decanted, washed with a small amount of water, filtered under vacuum and dried between sheets of filter paper at 50-60 0 C in an oven.
To obtain purer crystals, orthoboric acid is recrystallized. Calculate theoretical and practical output
Control questions:
1. Structural formula of borax, boric acid.
2. Dissociation of borax, boric acid.
3. Make up the formula of sodium tetraborate acid.
Laboratory work No. 7
Obtaining copper (II) oxide
Target: to obtain copper oxide (II) CuO from copper sulfate.
Reagents:
Copper (II) sulfate CuSO 4 2- * 5H 2 O.
Potassium and sodium hydroxide.
Ammonia solution (p = 0.91 g / cm 3)
Distilled water
Equipment: technochemical scales, filters, glasses, cylinders, vacuum pump(water jet vacuum pump) , thermometers, electric stove, Buchner funnel, Bunsen flask.
Theoretical part:
Copper (II) oxide CuO is a black-brown powder, at 1026 0 С it decomposes into Cu 2 O and О 2, almost insoluble in water, soluble in ammonia. Copper (II) oxide CuO occurs naturally in the form of a black earthy weathering product of copper ores (melaconite). In the lava of Vesuvius, it was found crystallized in the form of black triclinic tablets (tenorite).
Artificially, copper oxide is obtained by heating copper in the form of shavings or wire in air, at a red heat temperature (200-375 0 С) or by calcining carbonate nitrate. The copper oxide obtained in this way is amorphous and has a pronounced ability to adsorb gases. When calcined, at a higher temperature, a two-layer scale forms on the copper surface: the surface layer is copper (II) oxide, and the inner layer is red copper (I) oxide Cu 2 O.
Copper oxide is used in the production of glass enamels, to give a green or blue color, in addition, CuO is used in the production of copper-ruby glass. When heated with organic substances, copper oxide oxidizes them, converting carbon and carbon dioxide, and hydrogen into ode and being reduced in the process into metallic copper. This reaction is used in the elementary analysis of organic substances to determine the content of carbon and hydrogen in them. In medicine, it also finds use, mainly in the form of ointments.
2. Prepare a saturated solution from the calculated amount of copper sulfate at 40 0 С.
3. Prepare a 6% alkali solution from the calculated amount.
4. Heat the alkali solution to 80-90 0 С and pour the copper sulfate solution into it.
5. The mixture is heated at 90 ° C for 10-15 minutes.
6. The precipitated precipitate is allowed to settle, washed with water until the ion is removed SO 4 2- (sample BaCl 2 + HCl).
Laboratory work No. 5
Theoretical part
Complex (coordination) compounds- these are compounds in which at least one of the covalent bonds is formed by the donor-acceptor mechanism.
All coordination compounds are composed of inner sphere(complex particle), and in the case of cationic and anionic coordination compounds - and from external sphere... There is an ionic bond between the inner and outer spheres of the coordination compound.
Inner sphere (complex particle) consists of a central atom (a metal complexing atom) and ligands.
In the formula for complex compounds, the inner sphere is enclosed in square brackets. The inner sphere has no charge in neutral complexes, is positively charged in cationic ones, and negatively in anionic coordination compounds. The charge of the inner sphere is the algebraic sum of the charges of the central atom and the ligands.
Central atom- this is most often an ion of a d - element: Ag +, Cu2 +, Hg2 +, Zn2 +, Ni2 +, Fe3 +, Pt4 +, etc.
The coordination number of the central atom- the number of covalent bonds between the complexing agent and the ligands.
As a rule, the coordination number is twice the charge of the central atom. In most complex compounds, the coordination numbers are 6 and 4, less often 2, 3, 5 and 7.
Ligands- anions or molecules bound to the central atom by covalent bonds formed by the donor-acceptor mechanism. Ligands can be polar molecules (H2O, NH3, CO, etc.) and anions (CN–, NO2–, Cl–, Br–, I–, OH–, etc.).
Ligand dentition Is the number of covalent bonds with which this ligand is connected to the complexing agent.
Ligands are divided into monodentate (H2O, NH3, CO, CN–, NO2–, Cl–, Br–, I–, OH–), bidentate (C2O42-, SO42-, etc.) and polydentate.
For example, in the anionic complex compound K3: the inner sphere is 3–, the outer sphere is 3K +, the central atom is Fe3 +, the coordination number of the central atom is 6, the ligands are 6CN–, their dentate is –1 (monodentate).
The nomenclature of complex compounds ( IUPAC )
When writing the formula for a complex particle (ion), first write the symbol of the central atom, then the ligands in alphabetical the order of their symbols, but first the anionic ligands, and then the neutral molecules. The formula is enclosed in square brackets.
In the name of the coordination compound, the cation is indicated first (for all types of compounds), and then the anion. Cationic and neutral complexes have no special endings. In the names of anionic complexes, the ending –at is added to the name of the central atom (complexing agent). The oxidation state of the complexing agent is indicated by a Roman numeral in parentheses.
The names of some ligands: NH3 - ammine, H2O - aqua (aqua), CN– - cyano, Cl– - chloro, OH– - hydroxo. The number of identical ligands in a coordination compound is indicated by a prefix: 2– di, 3– three, 4– tetra, 5– penta, 6– hexa.
Cl diamminesilver (I) chloride or
diamminesilver (I) chloride
K2 potassium tetrachloroplatinate (II) or
potassium tetrachloroplatinate (II)
Diammintetrachloroplatin (IV)
Classification of coordination compounds
There are several classifications of coordination compounds: according to the charge of the complex particle, the type of ligands, the number of complexing agents, etc.
Depending on the charge of the complex particle, coordination compounds are divided into cationic, anionic, and neutral.
V cationic complexes the inner sphere is formed only by neutral molecules (H2O, NH3, CO, etc.), or by molecules and anions at the same time.
Cl3 hexaaquatalese (III) chloride
SO4 tetraamminecopper (II) sulfate
Cl2 tetraammine dichloroplatinum (IV) chloride
V anionic complexes the inner sphere is formed only by anions, or anions and neutral molecules at the same time.
K3 potassium hexacyanoferrate (III)
Na sodium tetrahydroxoaluminate (III)
Na sodium diaquatetrahydroxoaluminate (III)
Neutral (electrically neutral) complexes are formed with simultaneous coordination to the central atom of anions and molecules (sometimes only molecules).
Diamminedichloroplatinum (II)
Tetracarbonylnickel (0)
Depending on the type of ligands, the coordination compounds are subdivided into: acido complexes (ligands are acid residues CN–, NO2–, Cl–, Br–, I–, etc.); aquacomplexes (ligands are water molecules); ammino complexes (ligands are molecules ammonia); hydroxo complexes (OH– groups are ligands), etc.
Dissociation and ionization of coordination compounds
Cationic and anionic coordination compounds in solution completely dissociate by ionic bond into inner and outer spheres:
K4 → 4K + + 4–
NO3 → + + NO3–
Complex ions are ionized (dissociated) stepwise as weak electrolytes:
+ ⇄ + + NH3
+ ⇄ Ag + + NH3
Formation of coordination compounds
The formation of complex particles (ions) in solutions from metal ions-complexing agent and ligands occurs stepwise:
Ag + + NH3 ⇄ +
NH3 ⇄ +
and is characterized by stepwise formation constants:
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Ag + + 2NH3 ⇄ + 
The larger the numerical value of βn, the stronger (more stable) the complex ion.
Obtaining coordination compounds
Coordination connections are most often obtained in the following ways.
1. Interaction of ions of a metal complexing agent (usually a solution of a salt of a given metal) with ligands (a solution of a salt, acid, base, etc.):
FeCl3 + 6KCN → K3 + 3KCl
Fe3 + + 6CN– → 3–
2. Full or partial replacement of some ligands in the coordination compound with others:
K3 + 6KF → K3 + 6KSCN
3– + 6F– → 3– + 6SCN–
β6 1.70 103 1.26 1016
A new coordination compound is formed if its formation constant is greater than the formation constant of the initial coordination compound.
3. Substitution of the complexing metal in the coordination compound while preserving the ligands. As in the previous case, this transformation is possible if a more stable coordination compound is formed in this case.
SO4 + CuSO4 → SO4 + ZnSO4
2+ + Cu2 + → 2+ + Zn2 +
β4 2.51 109 1.07 1012
experimental part
Experience 1. Obtaining and destruction of hydroxo complexes
Pour 1 ml of solutions of zinc salts into two test tubes and aluminum(sulfates, chlorides or nitrates). Add 0.1 mol / L NaOH or KOH solution dropwise to each of the tubes until precipitation of the corresponding hydroxides forms. Write the reaction equations in ionic-molecular form, indicate the color of the precipitates.
ZnSO4 + 2NaOH → Zn (OH) 2¯ + Na2SO4
AlCl3 + 3NaOH → Al (OH) 3¯ + 3NaCl
Check the solubility of the precipitates obtained in a 2 mol / L solution of sodium or potassium hydroxide. Note your observations. Write the reaction equations in ionic-molecular form, indicate the color of the resulting solutions.
Zn (OH) 2 + 2NaOH → Na2
Al (OH) 3 + NaOH → Na or Na3
To destroy hydroxo complexes, add dropwise a 2 mol / L acid solution (HCl, H2SO4 or HNO3) to the resulting solutions. Note that as the acid is added, the solutions become cloudy or precipitates of the corresponding hydroxides are observed, which then dissolve in the excess acid. Write the reaction equations in ionic-molecular form.
Na2 + 2HNO3 → 2NaNO3 + Zn (OH) 2¯ + 2Н2О
Zn (OH) 2 + 2HNO3 → Zn (NO3) 2 + 2H2O
Na + HCl → Al (OH) 3¯ + NaCl ¯ + Н2О
Al (OH) 3 + 3HCl → AlCl3 + 3H2O
Experiment 2. Obtaining tetraamminecopper (II) sulfate and its destruction (qualitative reaction for the Cu2 + ion)
Pour 2 ml of copper sulfate solution into a test tube and add dropwise a 2 mol / L ammonia solution until a precipitate of hydroxymedium (II) sulfate (CuOH) 2SO4 is formed. Record the color of the precipitate formed. Place the coefficients and write the reaction equation in ionic-molecular form.
2CuSO4 + 2NH3 + 2H2O → (CuOH) 2SO4 + (NH4) 2SO4
Add concentrated ammonia solution to the test tube until the precipitate (CuOH) 2SO4 is completely dissolved. Write down the color of the tetraamminecopper (II) sulfate solution. Place the coefficients and write the reaction equation in ionic-molecular form.
(CuOH) 2SO4 + 6NH3 + (NH4) 2SO4 → 2SO4 + 2H2O
Divide the resulting solution of tetraammine copper (II) sulfate into two test tubes. Add a 2 mol / L sulfuric acid solution to the first tube, and sodium sulfide solution to the second. Note the color change of the solution in the first tube and the color of the formed precipitate in the second tube. Place the coefficients and write the reaction equations in ionic-molecular form. Under the formulas, indicate the color of the colored starting materials and reaction products.
SO4 + 2H2SO4 + 4H2O → SO4 + 2 (NH4) 2SO4
SO4 + Na2S → CuS + Na2SO4 + 4NH3
Experiment 3 Dissociation of complex compounds
Pour 3-5 drops of potassium chloride solution into a test tube and add a small amount (on the tip of a spatula) of crystalline sodium hexanitrocobaltate (III) Na3. Note the formation of a yellow precipitate of K2Na. This reaction is qualitative for potassium ions.
KCl → K + + Cl–
2 K + + Na + + 3– → K2Na¯
Pour 3-5 drops of iron (III) chloride solution into another tube, and then add 2-3 drops of thiocyanate solution ammonium or potassium. Note the color change in the solution. This reaction is qualitative for the Fe3 + ion.
FeCl3 → Fe3 + + 3Cl–
Fe3 + + 6SCN– ⇄ 3–
Carry out the appropriate qualitative reactions for K + and Fe3 + ions in a solution of potassium hexacyanoferrate (III) K3. Note your observations.
Which of the two equations for the dissociation of K3 in aqueous solution given below:
K3 → 3K + + 3–
K3 → 3K + + Fe3 + + 6CN–
consistent with your observations?
Formulate a conclusion about the nature of the dissociation of complex (coordination) compounds in aqueous solutions.
