J. Lewis proposed a more general theory of acids and bases.
Lewis foundations these are electron pair donors (alcohols, alcoholate anions, ethers, amines, etc.)
Lewis acids - they are electron pair acceptors , those. compounds having a vacant orbital (hydrogen ion and metal cations: H +, Ag +, Na +, Fe 2+; halides of elements of the second and third periods BF 3, AlCl 3, FeCl 3, ZnCl 2; halogens; tin and sulfur compounds: SnCl 4 , SO 3).
Thus, the Brønsted and Lewis bases are the same particles. However, Bronsted basicity is the ability to attach only a proton, while Lewis basicity is a broader concept and means the ability to interact with any particle that has a low-lying free orbit.
The Lewis acid-base interaction is a donor-acceptor interaction, and any heterolytic reaction can be represented as the interaction of a Lewis acid and a Lewis base:
There is no single scale for comparing the strength of Lewis acids and bases, since their relative strength will depend on which substance is taken as the standard (for Bronsted acids and bases, this standard is water). To assess the ease of acid-base interaction according to Lewis, R. Pearson proposed a qualitative theory of “hard” and “soft” acids and bases.
Rigid bases have high electronegativity and low polarizability. They are difficult to oxidize. Their highest occupied molecular orbitals (HOMO) are of low energy.
Soft grounds have low electronegativity and high polarizability. They oxidize easily. Their highest occupied molecular orbitals (HOMO) are high energy.
Hard acids have high electronegativity and low polarizability. They are difficult to recover. Their lowest free molecular orbitals (LUMO) have low energy.
Soft acids have low electronegativity and high polarizability. They are easy to recover. Their lowest free molecular orbitals (LUMOs) are high energy.
The hardest acid is H +, the softest is CH 3 Hg +. The hardest bases are F - and OH - , the softest ones are I - and H - .
Table 5. Hard and soft acids and bases.
|
Intermediate | ||
|
H + , Na + , K + , Mg 2+ , Ca 2+ , Al 3+ , Fe 3+ , BF 3 , AlCl 3 , RC + = O |
Cu 2+ , Fe 2+ , Zn 2+ , R 3 C + |
Ag + , Hg 2+ , I 2 |
|
Foundations |
||
|
H 2 O, OH - , F - , ROH, RO - , R 2 O, NH 3 , RNH 2 |
ArNH 2 , Br - , C 5 H 5 N |
R 2 S, RSH, RS - , I - , H - , C 2 H 4 , C 6 H 6 |
Pearson's principle of hard and soft acids and bases (HICA principle):
Hard acids preferentially interact with hard bases, while soft acids preferentially interact with soft bases.
This is expressed in high reaction rates and in the formation of more stable compounds, since the interaction between orbitals close in energy is more efficient than the interaction between orbitals that differ significantly in energy.
The principle of IMCA is used to determine the predominant direction of competing processes (reactions of elimination and nucleophilic substitution, reactions involving ambident nucleophiles); for the targeted creation of detoxifiers and medicines.
Theories of acids and bases
Theories of acids and bases- a set of fundamental physico-chemical concepts that describe the nature and properties of acids and bases. All of them introduce definitions of acids and bases - two classes of substances that react with each other. The task of the theory is to predict the products of the reaction between the acid and the base and the possibility of its occurrence, for which the quantitative characteristics of the strength of the acid and base are used. The differences between theories lie in the definitions of acids and bases, the characteristics of their strength and, as a result, in the rules for predicting the reaction products between them. All of them have their own area of applicability, which areas partially intersect.
Acid-base interactions are extremely common in nature and are widely used in scientific and industrial practice. Theoretical ideas about acids and bases are important in the formation of all conceptual systems of chemistry and have a versatile influence on the development of many theoretical concepts in all major chemical disciplines.
Based on the modern theory of acids and bases, such sections of chemical sciences as the chemistry of aqueous and non-aqueous electrolyte solutions, pH-metry in non-aqueous media, homo- and heterogeneous acid-base catalysis, the theory of acidity functions, and many others have been developed.
Evolution of ideas about acid-base interactions
Scientific ideas about the nature of acids and bases began to take shape at the end of the 18th century. In the works of A. Lavoisier, acidic properties were associated with the presence of oxygen atoms in the composition of the substance. The then known mineral and organic acids did indeed contain oxygen. This hypothesis quickly proved untenable when, thanks to the work of G. Davy and J. Gay-Lussac, a number of oxygen-free acids (for example, hydrogen halides, hydrocyanic acids) became known, while many oxygen-containing compounds do not exhibit acidic properties.
From the beginning of the 19th century, substances capable of interacting with metals with the release of hydrogen began to be considered acids (Yu. Liebig, 1839). Around the same time, J. Berzelius put forward an idea that explains the acid-base properties of substances by their electrical "dualistic" nature. So, he attributed electronegative oxides of non-metals and some metals (for example, chromium, manganese, etc.) to acids, and considered electropositive metal oxides to be bases. Thus, acidity or basicity is considered by Berzelius as a functional rather than an absolute property of a compound. Berzelius first attempted to quantify and predict the strength of acids and bases.
With the advent of the theory of electrolytic dissociation by S. Arrhenius (1887), it became possible to describe acid-base properties based on the products of electrolyte ionization. Thanks to the work of W. Ostwald, the theory was developed for weak electrolytes.
At the beginning of the XX century. American chemists G. Cady, E. Franklin and C. Kraus created the theory of solvosystems, which extended the provisions of the Arrhenius-Oswald theory to all solvents capable of self-dissociation.
The modern theories of acids and bases are based on the ideas of J. Bronsted and G. Lewis. There are quite successful attempts to create generalized theories (M. Usanovich, 1939), but they do not find wide application.
Liebig's hydrogen theory
Definitions. An acid is a substance that can react with a metal to release hydrogen. The concept of "foundation" in this theory is missing.
reaction products. When an acid reacts with a metal, a salt and hydrogen are formed.
Examples. Acid - HCl.
Reaction 2HCl + Zn = ZnCl 2 + H 2
Criteria for the progress of the reaction. Metals that are to the left of hydrogen in the activity series react with strong acids. The weaker the acid, the more active metal is needed for the reaction between them. Quantitative characteristics. Since the theory is rarely used, quantitative characteristics of the strength of the acid (and hence the prediction of the direction of the reaction) have not been developed within the framework of this theory.
Scope of applicability. Prediction of the interaction of hydrogen-containing substances with metals in any solvents.
specific features. In accordance with this theory, ethyl alcohol and ammonia are weak acids, as they are able to react with alkali metals:
2C 2 H 5 OH + 2Na \u003d 2C 2 H 5 ONa + H 2
2NH 3 + 2Na \u003d 2NaNH 2 + H 2
Arrhenius-Ostwald theory of electrolytic dissociation
Main article: Theory of electrolytic dissociation
For acid HA K = ·/
For base MOH K = ·/
For a reaction between an acid and a base to take place, the product of their dissociation constants must be greater than 10 -14 (the ion product of water).
Scope of applicability. It quite satisfactorily describes the reactions of sufficiently strong acids and bases with each other and the properties of their aqueous solutions. On the basis of ideas about the degree and constant of dissociation, the division of electrolytes into strong and weak was fixed, the concept of hydrogen index was introduced, the extension of which to alkaline media requires, however, additional assumptions (the introduction of the ionic product of water).
The theory can be used to describe the hydrolysis of salts and the reaction of acids and bases with salts, but this requires a very cumbersome apparatus - the proton theory (see below) is much more convenient.
The applicability of the Arrhenius-Ostwald theory is limited to aqueous solutions. in addition, it does not allow explaining the presence of the main properties of ammonia, phosphine and other compounds that do not contain hydroxo groups.
Proton Brønsted-Lauri theory
Main article: Protolytic theory of acids and bases
Model comparison
acid-base interaction
according to Lewis and Bronsted
Protolytic (proton) theory of acids and bases was proposed in 1923 independently by the Danish scientist J. Bronsted and the English scientist T. Lauri. In it, the concept of acids and bases was combined into a single whole, manifested in the acid-base interaction: A B + H + (A - acid, B - base). According to this theory, acids are molecules or ions that can be proton donors in this reaction, and bases are molecules or ions that add protons (acceptors). Acids and bases are collectively known as protoliths.
The essence of the acid-base interaction is the transfer of a proton from an acid to a base. In this case, the acid, having transferred a proton to the base, itself becomes a base, since it can again attach a proton, and the base, forming a protonated particle, becomes an acid. Thus, in any acid-base interaction, two pairs of acids and bases are involved, called Brönsted conjugated: A1 + B2 A2 + B1.
The same substance, depending on the conditions of interaction, can be both an acid and a base (amphoteric). For example, when interacting with strong acids, water is a base: H 2 O + H + H 3 O +, and when reacting with ammonia, it becomes an acid: NH 3 + H 2 O NH 4 + + OH -.
Solvosystem theory
Main article: Solvosystem theory
The theory of solvosystems is an extension of the Arrhenius-Ostwald theory to other ionic (in particular, protic) solvents. Proposed by American chemists G. Cady, E. Franklin and C. Kraus
Definitions. An ionic solvent is a solvent that self-dissociates into a cation and an anion. The cation is called the lyonium ion, and the anion is called the lyate ion. Protic solvent - a solvent capable of autoprotolysis, that is, the transfer of an H + ion from one molecule to another:
2HL ↔ H 2 L + + L -
These are solvents containing a sufficiently polar bond involving hydrogen and an unshared electron pair on some other non-metal (most often nitrogen, oxygen or fluorine).
Note: in this definition, the proton theory is "protected", because autoprotolysis is an acid-base reaction according to Breasted-Lowry. It also "hardwired" the Lewis theory, since it explains the reasons for the formation of lyonium ions.
The H 2 L + ion is called the lyonium ion, and L - is the lyate ion.
Acids are substances that form a lyonium ion in a given solvent.
Bases are substances that form a lyate ion in a given solvent.
Salts are substances that dissociate in a given solvent to form a cation and anion that are not lyonium and lyate.
reaction products. In the reaction of an acid with a base (neutralization reaction, a salt and a solvent are formed.
Examples.
Quantitative characteristics and criteria for the reaction The strengths of acids and bases are characterized by their dissociation constant.
The dissociation constants depend on the solvent. Protic solvents with high autodissociation constants ("acidic solvents", such as HF) differentiate acids (in which acids become weak and differ in strength) but level bases (all bases become strong, turning into a lyate ion). Protic solvents with low auto-dissociation constants ("basic solvents, eg NH3") differentiate bases but neutralize acids (which become strong when converted to lyonium).
The reaction proceeds from strong acids to weak acids.
Scope of applicability. Allows you to predict acid-base reactions in any solvent. Management of acid-base processes using a solvent. Extends to non-aqueous solutions the concept of pH (pH) as the concentration of lyonium ions. Describes the basic properties of substances that do not contain OH groups.
However, for many problems the theory is too cumbersome.
Specific features Some acid-base reactions in this theory can be turned upside down, for example
KOH (acid) + HCl (base) = KCl (solvent) + H 2 O (salt)
Lewis electronic theory
Main article: Lewis theory
In the theory of Lewis (1923), on the basis of electronic representations, the concept of acid and base was further expanded. Lewis acid - a molecule or ion that has vacant electron orbitals, as a result of which they are able to accept electron pairs. These are, for example, hydrogen ions - protons, metal ions (Ag +, Fe 3+), oxides of some non-metals (for example, SO 3, SiO 2), a number of salts (AlCl 3), as well as substances such as BF 3, Al 2 O 3 . Lewis acids that do not contain hydrogen ions are called aprotic. Protic acids are considered as a special case of the class of acids. A Lewis base is a molecule or ion capable of being an electron pair donor: all anions, ammonia and amines, water, alcohols, halogens. Examples of chemical reactions between Lewis acids and bases:
- AlCl 3 + Cl - → AlCl 4 -
- BF 3 + F - → BF 4 -
- PCl 5 + Cl - → PCl 6 - .
General theory of Usanovich
The most general theory of acids and bases was formulated by M. Usanovich in 1939. The theory is based on the idea that any acid-base interaction is a salt formation reaction. According to this theory, an acid is a particle that can remove cations, including a proton, or add anions, including an electron. A base is a particle that can accept a proton and other cations or donate an electron and other anions."(formulation 1964). Unlike Lewis Usanovich, the concepts of "acid" and "base" are based on the sign of the charge of the particle, and not the structure of the electron shell.
Usanovich's theory actually cancels one of the fundamental principles of classical chemistry - the notion of classes of acids and bases: " acids and bases are not classes of compounds; acidity and basicity are functions of a substance. Whether a substance is an acid or a base depends on the partner.» .
The disadvantages of Usanovich's theory include its too general nature and insufficiently clear definition of the wording of the concepts of "acid" and "base". The disadvantages also include the fact that it does not describe nonionic acid-base transformations. Finally, it does not allow for quantitative predictions.
According to Lewis, the acidic and basic properties of organic compounds are measured by the ability to accept or donate an electron pair, followed by the formation of a bond. An atom that accepts an electron pair is an electron acceptor, and a compound containing such an atom should be classified as an acid. An atom that provides an electron pair is an electron donor, and a compound containing such an atom is a base.
For example: Lewis acids include BF 3 , ZnCl 2 , AlCl 3 , FeCl 3 , FeBr 3 , TiCl 4 , SnCl , SbCl 5 , metal cations, sulfuric anhydride - SO 3 , carbocation. Lewis bases include amines RNH 2 , R 2 NH, R 3 N, ROH alcohols, ROR ethers, RSH thiols, RSR thioethers, anions, compounds containing π-bonds (including aromatic and heterocyclic compounds.
5.3. The concept of hard and soft acids and bases (the principle of HICA, the Pearson principle)
The general approach to dividing acids and bases into hard and soft can be characterized as follows.
Hard acids- Lewis acids, in which acceptor atoms are small in size, have a large positive charge, high electronegativity and low polarizability. The molecular orbital of hard acids, to which donor electrons pass, has a low energy level.
Soft acids - Lewis acids containing acceptor atoms of large size with a small positive charge, with low electronegativity and high polarizability. The molecular orbital of soft acids, which accepts donor electrons, has a high energy level.
Rigid bases- donor particles in which donor atoms have high electronegativity and low polarizability. Valence electrons are held firmly, the product is oxidized with difficulty. The orbital whose pair of electrons is transferred to the acceptor has a low energy level. Donor atoms in hard bases can be oxygen, nitrogen, fluorine, chlorine.
Soft grounds- donor particles, in which donor atoms have low electronegativity and high polarizability, they are easily oxidized; valence electrons are held weakly. An orbital whose pair of electrons is transferred to an acceptor has a high energy level. Donor atoms in soft bases are carbon, sulfur, and iodine atoms. Table 4
According to Pearson's principle of hard and soft acids and bases (HMCA), Lewis acids are divided into hard and soft. Hard acids- acceptor atoms with small size, large positive charge, high electronegativity and low polarizability.
Soft acids- acceptor atoms of large size with a small positive charge, with a small electronegativity and high polarizability.
The essence of LCMO is that hard acids react with hard bases and soft acids react with soft bases. For example: when sodium ethoxide interacts with isopropyl iodide, ethoxide - C 2 H 5 O ion - as a hard base will react with a hard acid, which is a proton in - position. The elimination reaction will be predominant. 
Lewis acid - a molecule or ion that has vacant electron orbitals, as a result of which they are able to accept electron pairs. For example, hydrogen ions are protons, metal ions (Ag +, Fe 3+), oxides of some non-metals (SO 3, SiO 2), a number of salts (AlCl 3), substances like BF 3, Al 2 O 3. Lewis acids that do not contain hydrogen ions are called aprotic. Protic acids are considered as a special case of the class of acids.
A Lewis base is a molecule or ion capable of being an electron pair donor: all anions, ammonia and amines, water, alcohols, halogens.
Examples of chemical reactions between Lewis acids and bases:
AlCl 3 + Cl - → AlCl 4 -
BF 3 + F - → BF 4 -
PCl 5 + Cl - → PCl 6 - .
The ionic potential is the ratio of the electronic charge of an ion to its effective radius.
Expressed by the ratio Z/r, where Z is the charge, r - ion radius. Used to characterize the interaction of an ion in a crystal lattice or in solution
Rigid bases include donor particles with high electronegativity, low polarizability, and difficult to oxidize. The compound firmly holds its electrons, its molecular orbital, the pair of electrons of which is transferred to the acceptor, has a low energy level. Soft bases. These include donor particles with low electronegativity, high polarizability, and rather easily oxidized. They weakly hold their valence electrons, their molecular orbitals, have a high level of energy (electrons are removed from the atomic nucleus).
Hard acids. includes Lewis acids, in which acceptor atoms are small in size, have a large positive charge, high electronegativity and low polarizability. The molecular orbital has a low energy level. Soft acids. includes Lewis acids containing acceptor atoms of large size with a small positive charge, with low electronegativity and high polarizability. The molecular orbital has a high energy level. The essence of the principle of HICA is that hard acids preferentially react with hard bases, and soft acids with soft bases. high reaction rates, the formation of more stable compounds
Ticket number 2 1. Halogens. Degrees of oxidation. Disproportionation of halogens. Comparison of oxidizing power. Hydrogen halides and hydrohalic acids. HF features. Halides of metals and non-metals, their interaction with water. Halogen oxides.
In the ground state, halogen atoms have the electronic configuration nsnp5. fluorine to a smaller radius, large values of ionization energy and electronegativity. The electron affinity of fluorine is less than that of chlorine. fluorine oxidation states -1, 0.
Halogen compounds in positive oxidation states exhibit oxidizing properties.
Halogens are the most active non-metals. Fluorine interacts with almost all simple substances, with the exception of light inert gases. From fluorine to iodine, the oxidizing ability decreases, and the reducing ability increases. Chlorine reacts with oxides of some metals: magnesium, aluminum, iron.
2MgO + 2С12 = 2MgCl2 + 02
Bromine is a strong oxidizing agent. In the aquatic environment, it oxidizes sulfur to sulfuric acid:
ZBr2 + S + 4H20 = bHBr + H2SO4
potassium manganate - to permanganate:
2K2Mp04 + Vg2 = 2KMp04 + 2KVg
The oxidizing properties of iodine are less pronounced than other halogens. Iodine is not able to oxidize not only oxygen, but also sulfur. Iodides have reducing properties. Under the action of chlorine, bromine, hydrogen peroxide and nitric acid, it is oxidized in an aqueous medium to iodic acid HNiu3:
3I2(solid) + 10HNO3(100%) = 6НIO3 + 10NO2 + 2Н20
Under standard conditions, hydrogen halides are colorless gases with a pungent odor. for HF, the melting and boiling points. The abnormally high melting and boiling points of hydrogen fluoride are explained by the enhancement of intermolecular interaction due to the formation of hydrogen bonds between HF molecules. Solid hydrogen fluoride consists of
zigzag polymer chains. For HCl, HBr, HI, the formation of hydrogen bonds is not typical due to the lower electronegativity of the halogen atom. Aqueous solutions of HC1, HBr, and HI behave like strong acids. hydrofluoric HF and hydrochloric HC1 acids do not interact with concentrated sulfuric acid, but HBr and HI are oxidized by it:
2НВг + H2S04(koh4.) = Br2t + S02 + 2H20
8HI + H2S04(koh4.) = 4I2 + H2S + 4H20
Alkali and alkaline earth metal halides are ionic substances. They are soluble in water and have high melting and boiling points.
Hypohalogenitic acids HCO are known only in dilute aqueous solutions.
Hypohalogenic acids are weak. acidic properties in the series HSiu-HBrO-Niu are weakened, while the basic ones increase. Iodous acid is already an amphoteric compound.
Solutions of hypohalites have a strongly alkaline reaction, and passing CO2 through them leads to the formation of acid:
NaCIO + H20 + C02 = NaHC03 + NSO
Hypohalogenic acids and their salts are strong oxidizing agents:
Of the oxo acids HXO2, hydrochloric acid HClO2 is known.
HClO2 is a medium strength acid.
HXO3 oxo acids are more stable than hypohalogenic acids. Chloric HClO3 and bromic HBIO3 acids were obtained in solutions with concentrations below 50%, while iodic HClO3 was isolated as an individual substance. Solutions of HclO3 and HBrO3 are obtained by the action of dilute H2SO4 on solutions of the corresponding salts, for example:
Ba(СlO3)2 + H2SO4 = 2НСО3 + BaSO4
Iodine acid is produced by the oxidation of iodine with fuming nitric acid.
acid, hydrogen peroxide solution:
I2 + 5H202 = 2HI3 + 4H20.
HXO3 are strong acids. In the series HClO3-HBrO3-HI3, there is a slight decrease in the strength of acids.
Perchloric acid HC104. is released in the form of HC104*H20 hydrates. Bromic acid HBrO4 is known only in solutions.
Liquid HF is composed of HF polymer chains.
The halogen-oxygen bond is fragile, which is caused by strong mutual repulsion of atoms with a high
Electronegativity. halogen oxides are unstable. Oxygen difluoride OF2 can be obtained
2F2 + 2NaOH =OF2 + 2NaF + H20
Oxygen difluoride is a strong oxidizing-fluorinating agent.
When an electric discharge is passed through a cooled mixture of fluorine and oxygen, another fluoride, 02F2, can be obtained.
Chlorine oxide (I) C120 It is obtained
3HgO + 2C1 2 = Hg30 2 Cl 2 + Cl 2 O
The connection is extremely unstable.
2. Titanium, zirconium, hafnium. Comparison of redox properties. Interaction of metals with solutions of acids and alkalis. Difference of Ti compounds from Zr and Hf. Reactions of Ti 2+ and Ti 3+ compounds. E 4+ compounds: oxides, a- and b-forms of acids. Halides, their hydrolysis. Salts of oxocations. halide complexes.
Ionization in the transition from titanium to zirconium is markedly reduced.
Only the first element of the group, titanium, exhibits high chemical activity. Hafnium has lanthanide contraction. +4 oxidation state is characteristic, most compounds are covalent. In the Ti - Zr-Hf series, the stability of compounds with the highest oxidation state increases. So, oxides of TiO, Ti203, Ti02 and fluorides TiF2, TiF3, TiF4 are stable for titanium, and only dioxides Zr02, Hf02 and tetrafluorides ZrF4, HfF4 are stable for zirconium and hafnium. Titanium is more prone to low oxidation states +2, +3 than its heavy analogs. Compounds of zirconium(III) and hafnium(III) do not exist in aqueous solutions. oxidation states, the basic and reducing properties are enhanced
For titanium, the coordination number is typically 6 and, more rarely, 4; zirconium and hafnium 7 and 8.
The reaction with halogens begins at low heating, always forming MX4 tetrahalides.
Unlike zirconium and hafnium, titanium reacts with hydrochloric and dilute sulfuric acids when heated.
2Ti + 6HC1 = 2TiCl3 + 3H2T
Titanium also dissolves in concentrated hydrofluoric acid to form green solutions.
2Ti + 6HF = 2- + Ti2+ + ZN2T
Ti + 6HF + 02 = H2 + 2H20
Extremely slowly, titanium dissolves in dilute and concentrated nitric acid, as well as in aqua regia - the reaction is prevented by the formation of a layer
Ti + 4H2S207 - Ti(S04)2 + 2S02T + 4H2S04
When heated, titanium powder slowly dissolves in concentrated solutions and alkali melts:
Ti + 2NaOH + H20 = Na2Ti03 + 2H2
Zirconium and especially hafnium are more resistant to acid oxidation. do not react with any of the dilute acids except hydrofluoric. zirconium and hafnium react vigorously only with a mixture of nitric and hydrofluoric acids:
3M + 4HN03+ 21HX = 3H3[MX7] + 4NO + 8H20
The interaction of zirconium and hafnium with hydrofluoric acid and concentrated sulfuric acid proceeds more slowly:
M + 7HF = H3 + 2H2T
M + 5H2S04 = H2 + 2S02t + 4H20
Concentrated HN03 improves the corrosion resistance of metals. Zirconium and hafnium do not react with alkalis.
