Acids, Bases, Solvents and Reactions in non-aqueous solvents

1. Acids, Bases, Solvents and Reactions in non-aqueous solvents Prof.Pote A.P. Department of Chemistry Jijamata College of Science and Arts, Bhende bk.

4. (a) Measuring the pH of vinegar. (b) Measuring the pH of aqueous ammonia.

5. A circle of shiny coins is created by the reaction between the citric acid of the lemon and the tarnish on the surface of the copper. Source: Fundamental Photos

6. Acids and Bases: i) Arrhenius Theory • In 1884, Arrhenius offered modern approach to acid & base concept. • He was awarded the Nobel Prize in chemistry in 1903. • According to his theory. • Acids: It is hydrogen compound which gives hydrogen ions (H+) , in aqueous solution , eq. HCl, HNO3. H2SO4 The ionization of acids like HCl, HNO3 can be represented as HCl H+ + Cl- HNO3 H+ + NO3-

7. Arrhenius Theory • Bases: It is hydroxide compounds, which gives hydroxide ions (OH-) in aqueous solution. • eq. NaOH, KOH. The ionization like NaOH, KOH can be represented as NaOH Na+ + OH- KOHK+ + OH- • The process of neutralization involves the union or combination of the H+ and OH-ions to forms water, thus • HCl + NaOHNaCl+ H2O • i.e. H+ + OH- H2O

8. Arrhenius Theory Advantages: • To determine acid base properties of aqueous solution. • It explains neutralization, strength of acid & base hydrolysis of salts, etc. Disadvantages / Limitation: • It defines acids & bases in terms of only aqueous solution & not in solution of other solvents like benzene, chloroform, acetone, etc. • Some compounds which do not gives OH ion in water but they exhibit basic properties Eq. CH3COO- ion, pyrite, ammonia, etc. • They cannot explain acidic nature of AlCl3, which does not having H+ ions. • The neutralization is limited only to reactions taking place in agues medium.

9. II) Bronsted – Lowry Theory Johannes Bronsted and Thomas Lowry proposed in 1923 that acid – base reactions is transfer of proton from one to another. Acid: It is proton donor i.e. it is a substance which is capable of donating a proton. Base: It is proton acceptor i.e. it is a substance which is capable of accepting a proton.

10. Bronsted – Lowry Theory Conjugate Acid – Base pair: • It is the pair of substances which is formed by loss or gain of proton from each other is called as conjugate acid – base pair. • Thus acid (HA) and base (B) are conjugate with one another, when related by the transfer of proton. • Every Bronsted acid has conjugate base, & every Bronsted base has conjugate acid. • CH3C00–is conjugate base obtained from Acetic acid CH3OOH. Similarly H2O is conjugate base of acid H3O+.

11. Molecular model: The reaction of an acid HA with water to form H3O+ and a conjugate base. Acid Base Conjugate Conjugate acid base

13. Bronsted – Lowry Theory Advantage: • Acids & bases are defined in terms or substance & not in term of their solution. • The concept recognizes acid-base behavior of the substance irrespective of nature of solvent. • It helps in explaining of hydrolysis of salt. Disadvantage / limitation: • There are many substances like BF3, AlCl3 which do not possess proton but still behave. As an acid. • Number of acid-base reaction, not involving transfer of a proton. • It cannot explains acid – base reaction in non – protein solvents like COCl2, SO2, N2O4, etc.

14. III) Lewis concept In 1923 Lewis proposed the concept of acid – bases. The definition was based on following four assumptions • Rapidity of neutralization: The process of combination of acid & bases is rapid. H3O+ + OH- 2H2O • Displacement of weaker acid or base: The acids or bases will displace a weaker acid or base from its compounds. Eg. Stronger base, OH-, displaces the weaker base acetate ion from acetic acid. Thus CH3COOH + OH- 2H2O + CH3COO-

15. Lewis concept • Titration using indicators: The acids & bases are titrated against one another by using substance indicator. Eg crystal violet gives yellow colour in acid and violet colour in base solution. • Acid – Base catalyst : Ex. 1) Acid catalyst – Fridel Craft, Halogenations. 2) Base catalyst - Aldol condensation.

16. Lewis concept Definitions : Acid – It is any substance or species that capable of accepting an electron pair. Ex. : i) All cations eg.- Cu2+, Fe3+, Ag+ ii) Electron deficient molecule, eg. BF3, AlCl3 Base – It is any substance or species that capable of donating an electron pair. Eg. i) all anions eg. OH-, F-, CN-, ii) Electron pair donor, eg. NH3, H2O, etc. iii) Olefinic compounds like ethylene (CH2=CH2), C=O, N=O,, acetylene.

17. Lewis concept Merits : 1)The definition of acids & bases are depends upon electron configuration and distribution in substances, 2)Acid and bases are in terms of only in an element & not in term of their solutions. 3)A variety of compounds and ions are included under this theory. 4)This theory does not lower the status of any previous definition of acids and bases. Demerits : 1) It is very wide concept. 2) It fails to explain strength of acid & bases. 3) Lewis concept assumes that combination of acid and base is rapid reaction, but this is not for all cases.

18. V) Lux – Flood concept It explains acid base characters in terms of the oxide ions. • Acids : It is an oxide ion acceptor. • Base : It is an oxide ion donor.

19. Strength of Acid – Base The strength of Protonic acid can be determined by its ability to gives up proton. The acid HA is ionizes in water as HA H+ + A- The equilibrium constant for this reaction is according to law of mass action. [H+] + [A-] K = [HA] The reaction is carried out in aqueous medium, thus HA + H2O H+ + A- Thus equilibrium constant for this reaction is [H3O+] + [A-] K = [HA] [H2O]

20. strength of Protonic acid • Generally [H2O] is very large and thus it is taken constant. On rearranging the above equation, we have • [H3O+] + [A-] • K[H2O] = Ka = • [HA] [H2O] • Where Ka is known as Dissociation or Ionization constant.. • The large value of Ka indicates that the acid is strong, it is always expressed as pKa (pKa = -logKa), • for strong acids pKa value are smaller while for weak acids, it is higher.

21. strength of Protonic base It can be determine as shown B + H2O BH+ + OH- [BH+] + [OH-] K = [B] [H2O] Generally [H2O] is very large and thus it is taken constant. On rearranging the above equation, we have [BH+] + [OH-] K [H2O] = Kb = [B] Where the large value of Kb indicates that the base is strong, It is always expressed as pKb (pKb = -log Kb), for strong bases pKb value are smaller, while for weak bases, it is higher.

23. Trends in strength of Hydracids Hydracids : The compounds in which protons is directly attached to the non –metal X is called as hydracids and shown by general formula HX. Strength of hydracids : • Strength of H – X bond: The acids will be stronger when H–X bond is weaker. Eg. HF is weaker acid than HI, where H–F bond energy is 135 KCal / mole & H–I is 72 KCal / mol. • Electronegativity of X: Which increase in electronegativity of X, there is increase of strength of acid. Because high electronegative causes lost of proton. • eg. LiH is weaker acid than HF.

24. Trends in strength of Oxiacids Oxyacids : The compounds in which protons is attached to the non–metal X through oxygen called as Oxyacids, shown by general formula HOX or XOM (OH)n. Strength of Oxyacids : • Size of X : As size of X increases, its electronegativity decreases. Hence release of proton becomes more difficult and will be result decreases strength of acid. Eg. HOCl is strong acidic than HOI. It is due to increasing in size from Cl to I. • Electronegativity of X : As electronegativity deceases, strength of acid is also decreases. Eg. NaOH is basic while HOClis acidic because Cl is more electronegative than Na.

25. Addition of oxygen : In the structure of Oxyacids if additional oxygen atom are attached to X, then the acidity of Oxyacids is increases. Eg. HOCl4 > HOCl3 > HOCl2 > HOCl When negative charge on acid is increases after successive dissociation, then acidity of oxyacids are increases. Eg. H3PO4 > H2PO41- > HPO42- > PO43-

26. Strength of Lewis acid and base a) Electronic effect : Greater the electronegativity of substituent, higher the Lewis acidity and lower its Lewis basicity. Eg. 1) strength of acids BF3 > BH3 > B(CH3)3 Fluorine is highly electronegative, it attracts bonding pair of electrons towards it self making boron electron deficient. Hence BF3 is strong Lewis acids. While in case of B(CH3)3, methyl group is electron donor, it gives electron density to the boron atom, & make it slightly electron rich, hence B(CH3)3 is weak Lewis acid. Eq. 2) strength of bases NF3 < NH3 < N(CH3)3 In case of NF3, Fluorine is more electronegative atom it attract electron pair toward it self. It make nitrogen electron deficient. Hence NF3 is weak Lewis base. While in N(CH3)3 is strong Lewis base. In short, hence greater the electronegativity of substituent, higher the higher the Lewis acidity and decreases Lewis basicity.

27. Strength of Lewis acid and base b) Steric effect : This steric effect is quite large in Lewis acids. Strength of base is increase due to steric effect along with inductive effect.

28. Solvent • The solvent is a substance which can dissolve other substances. • Water is the most widely used solvent. • The number of substance dissolve in water and most of reactions carried out in water. • Another non-aqueous solvents are anhydrous liquid ammonia (NH3), liquid sulphur dioxide (SO2), anhydrous hydrogen sulphide (H2S). • The acidic and basic properties of solute depends upon the properties of the solvent.

29. properties of the solvent • Melting Point & Boiling Point : The M. P. & B. P. of solvent gives range of temperatures in which solvent remains in liquid state. Thus solvents which are liquids at room temperature at 1 atmospphericpressureare most useful solvent. Eg. Water 0-1000C, the gases like ammonia and sulphur dioxide can be used as solvent at low temperature. • Dipole moment :Polar solutes dissolve in solvent with high dipole moment. Higher the dipole moment of solvent, more is its capacity to dissolve the polar solute. • Dielectric constant : When molecule placed in an electric field, would orient themselves to neutralize the field. A molecule with high dipole moment will orient in an electric field considerably more then molecule with lower dipole moment, have greater dielectric constant value. At the room temperature water has high dipole moment and dielectric constant, thus water is better solvent than ammonia.

30. properties of the solvent Lewis Acid–Base Character: • When ionic compound dissolves in polar solvent, compound is ionized into cation & anion. • Generally cations of solute compound are better solvated than anion because smaller size of cation companied to anion. • eg. Cations like Na+, Ca2+. Fe3+ is, more solvated than anions like Cl-, Br-, I-, etc. • The solvation of cation is due to interaction between cation and negative end of solvent. In such interaction cation act as acceptor of electrons i.e. Lewis acid. And solvent act as donor of electrons i. e. Lewis base. • The solvation of anion, interaction between anion and positive end of solvent molecule is takes place. Here anion acts as donor of electron (Lewis Base). And solvent act as acceptor of electrons (Lewis Acid).

31. Solvation • Solvation, also sometimes called dissolution, is the process of attraction and association of molecules of a solvent with molecules or ions of a solute. • As ions dissolve in a solvent they spread out and become surrounded by solvent molecules. Solvation Energy :- Energy is liberated during solvation called as Solvation Energy. Hydration Energy:- When water is used as solvent then as solvent then energy evolved is called as hydration energy.

32. properties of the solvent • Protonic Acidity or Basicity : They are three types of solvent : • Acidic or photogenic solvent : These solvent are having strong tendency to donate protons. EqHCl, H2SO4, HNO3, CH3COOH • Basic of protophillic solvent : These solvents are having strong tendency to accept protons. Eq. NH3, amines, aniline, etc. • Amphoteric or Amphiprotic solvent : These type of solvent can accept as well as donate protons. Eq H2O, CH3COOH . etc. • Auto Dissociation of Solvent: The solvents like H2O, NH3, HCl, HF, H2SO4 are having tendency to undergoing self dissociation or auto dissociation. 2H2O = H3O+ + OH- 2HCl = H2Cl+ + Cl-

33. properties of the solvent Non–Protonic Solvent / Aprotic solvent: These solvents do not contain proton in their formula. They do not have any tendency of donating or accepting protons. Eq CCl4, C6H6, CHCl3, BF3, etc. The non-Protonic solvent are of three types : • Non–Polar, Non–Dissociated Liquid : These type of liquid is non–polar, low dipole moment & low dielectric constant. They do not have any tendency to donate or self ionizing capacity. These type of solvent are useful for non–polar or covalent solute. Eq. CCl4. • Polar, Non–Dissociated Liquid : These solvent is non–protonic. They do not have self–ionizing capacity. some ionic compound dissolved in this type of solvent. Eq. Methyl cyanide (CH3CN), tetrahydrofuran, dimethyl formamide. • Highly Polar, Self–Ionizing Liquid : These solvent are highly polar, high dielectric constant, high dipole moment. all ionic compounds are soluble in this solvent. Eq. H2O, SO3, BrF3, etc.

34. Hard & soft acids and bases Hard acids and hard bases • Small atomic / ionic radius • High oxidation state • Low polarizability • High electronegativity (bases) • Hard bases have highest-occupied molecular orbitals (HOMO) of low energy, • Hard acids have lowest-unoccupied molecular orbitals (LUMO) of high energy. • Examples of hard acids are: H+, light alkali ions (Li through K all have small ionic radius), Ti4+, Cr3+, Cr6+, BF3. • Examples of hard bases are: OH–, F–, Cl–, NH3, CH3COO–, CO32–. • The affinity of hard acids and hard bases for each other is mainly ionic in nature.

35. Hard & soft acids and bases Soft acids and soft bases • large atomic / ionic radius • low or zero oxidation state • high polarizability • low electronegativity • soft bases have HOMO of higher energy than hard bases, • soft acids have LUMO of lower energy than hard acids. • Examples of soft acids are: CH3Hg+, Pt2+, Pd2+, Ag+, Au+, Hg2+, Hg22+, Cd2+, BH3. • Examples of soft bases are: H–, R3P, SCN–, I–. • The affinity of soft acids and bases for each other is mainly covalent in nature.