Acids, Bases and Salts

Chemistry EQUILIBRIUM ACIDS, BASES AND SALTS  ACIDS BASES AND SALTS Arrhenius Concept Arrhenius Acid: Substance which gives H+ ion on dissolving in water (H+ donor) Ex.  HNO3, HClO4, HCl, HI, HBr, H2SO4, H3PO4 etc.  H3BO3 is not Arrhenius acid. (it is a Lewis base) H+ ion in water is extremely hydrated (in form of H3O+, H5O2+, H7O3+, general form H+ (H2O)n The structure of solid HClO4 is studied by X-ray, it is found to be consisting of H3O+ and ClO4– HClO4 + H2O  H3O+ + ClO4– (better representation) Arrhenius Base: Any substance which releases OH– (hydroxyl) ion in water (OH– ion donor) OH– ion is present also in hydrated form of H3O2–, H7O4–, H5O3– . general form OH– (H2O)n First group elements (except Li.) form strong bases Limitation of Ostwald Dilution Law (1) It is not applicable for strong electrolyte (2) It is not applicable for saturated solution. Modified Arrhenius Concept It rectifies most of the above limitations (i) Water is weak electrolyte and ionizes to a very weak extent. H2O  H+ + H2O  H+ + OH– H3 O+ 1 Chemistry H2O + H2O  H3 O+ + OH– Above reaction is called Autoionization or salinization of water. (ii) Water is neutral in nature i.e. [H3O+] = [OH–] (iii) The substances which increase the H3O+ ion concentration act as acids and while those which increase OH– ion concentration act as bases. Ex. (a) SO2 + H2O  H2SO3  H3O+ + HSO3– Acid (b) NH3 + H2O  NH4OH NH4+ + OH– Base Basicity or Precocity of an Acid Types of bases Mono acidic base  Gives single OH- per molecule e.g. CsOH, RbOH NH4OH Di acidic base  Gives two OHper molecule e.g. Ba (OH)2, Ca(OH)2 Tri acidic basic  Gives three OHper molecule e.g. Fe(OH)3 It is number of H+ ions furnished by a molecule of an acid. An acid may be classified according to its basicity. Thus we may have, (i) Mono basic or Mono protic acids like HCl, HNO3, CH3COOH, HCN etc. (ii) Dibasic or Diprotic acids like, H2SO4, H2CO3, H2SO3, H2S etc. (iii) Tribasic or Triprotic acids like H3PO4,H3AsO4 etc. 2 Chemistry Acidity or Hydroxity of a Base Types of acids Mono basic acid  Gives single H+ per molecule e.g., HA (HNO3, HClO4,HCl) Di basic acid  Gives two H+ per molecule e.g. H2A (H2SO4, H3PO3) Tri basic acid  Gives three H+ per molecule e.g.H3A (H3PO4) It may be defined as the number of OH– ions furnished by a molecule of a base. A base can be, (i) Mono acidic or Monohydroxy like NaOH, NH4OH, AgOH etc. (ii) Di acidic or dihydroxy like Ba(OH)2, Mg(OH)2, Ca(OH)2, Sr(OH)2 etc. (iii) Tri acidic or trihydroxy like Fe(OH)3, Al(OH)3 etc. Strength of Acid or Base (i) Strength of Acid or Base depends on the extent of its ionisation. Hence equilibrium constant Ka or Kb respectively of the following equilibria give a quantitative measurement of the strength of the acid or base. (ii) HA + H2O  H3O+ + A– ; Ka = [H3 O+ ][A− ] [HA] (iii) Similarly B + H2O  BH+ + OH– ; Kb= [BH+ ][OH− ] [B] here H2O is solvent. Amphoteric: Substances which can act as acid as well as base are known as amphoteric HCl + H2O  H3O+ + Cl– base NH3 + H2O  NH4+ + OH– acid 3 Chemistry Amphiprotic: An amphiprotic molecule (or ion) can either donate or accept a proton, thus acting either as an acid or a base. Water, amino acids, hydrogen carbonate ions and hydrogen sulfate ions are common examples of amphiprotic species. Since they can donate a proton, all amphiprotic substances contain a hydrogen atom. An acid–base reaction always proceeds in the direction of formation of the weak acid and the weak base. In the equilibrium, HA + H3O+ +  H2O A– Strong acid Strong base Weak acid Weak base In general "The conjugate base of a strong acid is always a weak base and the conjugate base of a weak acid is always a strong base." A number of organic compounds containing oxygen, can accept protons and thus act as bases. Ex. ̈ + H2SO4 C2H5 OH (a) Ethanol C2H5 H2 + HSO4– (Oxonium ion)  + (C2H5)2 O  + HCl  (C2H5) ̈ 𝑂: (b) 2 H + Cl– Ethyl ether Oxonium ion Bronsted lowery concept does not differ appreciably from the Arrhenius theory for aqueous solution only. Autoionization or Autoproteolysis or Self ionisation H2O + H2O  H3O+ + OH– acid base NH3 + NH3  NH4+ + NH2– acid base H2SO4 + H2SO4  H3SO4+ + HSO4– acid base A limitation of the Bronsted Lowery theory is that the extent to which a dissolved substance can act as an acid or a base depends largely on the solvent. (a) HClO4 + HF  Acid Base H2F+ + ClO4– ( HClO4 acts as a acid in HF ) Acid Base 4 Chemistry (b) HNO3 behaves as base in HClO4 and HF H3O+ + NO2+ + 2HF2– HNO3 + 4HF  (base) (acid) Urea is weak acidic in liquid NH3 (c) NH2CONH2 + NH3  NH4+ + NH2CONH– Acid Base Acid Base Note: H2SO4 also acts as base in HF solvent. Classification of Bronsted - Lowery Acids and Bases Bronsted - Lowery acids and bases can be (i) Molecular (ii) Cationic (iii) Anionic Table - 1 Acid Type Base Molecular HCl, HNO3 , HClO4 , NH3 , N2 H4 , Amines, H2 SO4 , H3 PO4 , H2 O, Alcohol, Ethers etc. H2 O etc. Cationic NH4+ , N2 H5 + , PH4+ , [Fe(H2 O)5 OH]2+ Na+ , Ba2+ (All cations) [Al(H2 O)5 OH]2+ etc. [Fe(H2 O)6 ]3+ , [Al(H2 O)6 ]3+ etc. Anionic HS − , HS𝑂3− , H2 P 𝑂4− , HS𝑂4− HCO3 − , HPO4 2− etc. all amphiprotic anions − Cl− , Br − , OH − , HSO− 4 , CN 𝐶𝑂32− , 𝑆𝑂42− , 𝑁𝐻3− , CH3 COO− etc. all anions. Reactions in Non-Aqueous Solvents (i) Solvents like C6H6, CCl4, THF (Tetrahydrofuran), DMF (N, N-dimethyl formamide) etc. are used in organic chemistry. In inorganic chemistry reactions are generally studied in water. However a large number of non–aqueous solvents 5 Chemistry (such as Glacial acetic acid, Hydrogen halides, SO 2 etc.) have been introduced in inorganic chemistry. (ii) The physical properties of a solvent such as M.P., B.P., Dipole moment and Dielectric constant are of importance in deciding its behaviour. Classification of Solvents There are two types of solvents (A) Protonic (protic) (B) Aprotic (A) Protonic or Protic Solvents (i) They are characterized by the presence of a transferable hydrogen and the formation of "Onium" ions Autoionization taking place in them. (a) H2O + H2O  H3O+ + OH– (b) NH3 + NH3  NH4+ + NH2– (c) 3HX  H2X + HX2– (d) 2H2SO4 H3SO4+ + HSO4– (ii) Protonic solvents may be Ex. (a) Acidic (Anhydrous sulphuric acid, liquid HF, Glacial acetic acid etc.) (b) Basic (liquid NH3) (c) Amphiprotic (H2O, proton containing anions) (B) Aprotic Solvents Such solvents do not have replaceable hydrogen in them. These can be classified into three categories (a) Non-polar or very weakly polar, no dissociated liquids, which do not solvate strongly. Ex. CCl4, hydrocarbons, C6H6, C6H12 etc. (b) Non-ionised but strongly solvating, generally polar solvents. 6 Chemistry Ex. Acetonitrile CH3CN, DMF, DMSO (dimethyl sulfoxide), THF and SO2. (c) Highly polar, autoionizing solvents. Ex. Inter halogen compounds (BrF3, IF5 and trichloro phosphine) (a) 2 BrF3  (b) 2 IF5  (c) 2Cl3PO  BrF2+ + BrF4– IF4+ + IF6– Cl2PO+ + Cl4PO– Levelling Solvents (i) The Bronsted - Lowery theory can be extended to acid - base reactions in nonaqueous solvents. It can be used in differentiating the acid strength of a particular acid and in titration of weak bases. (ii) In water solvent, mineral acids appear to be equally strong because of their complete ionisation, water is called here a levelling solvent because it levels all the acids to the same strength. (iii) If instead of water solvent, we take mineral acids in pure acetic acid solvent (which is poor proton acceptor as compared to water) it is found acids become weak and can be differentiated. HCl + CH3COOH  Cl– + CH3COOH2+ Acid Base Base Acid In above example acetic acid and Cl– ions both compete for protons and the former Ex. being a poor proton acceptor does it much less effectively than water. Thus, HCl in acetic acid solvent appears to be a much weaker acid than that in water. (iv) Mineral acids in acetic acid solvent follow the following order of their strengths. HNO3 < HCl < H2SO4 < HBr < HClO4 (v) A weak base like acetamide or acetanilide in aqueous medium cannot be titrated with acids. If however, the weak base is taken in glacial acetic acid solvent, the former behaves as a strong base and can be titrated. This is because acetic acid (which acts as a better proton donor) exerts a levelling effect on the base. 7 Chemistry Lux - Flood Concept (1939 & 1947)* (i) The proton plays an important role in explaining the acid-base behaviour in the Bronsted-Lowery concept. Lux observed that acid - base reactions are also feasible in oxide systems without the aid of protons. (ii) Above approach was extended by Flood and applied to non-protonic systems, which were not covered by the Bronsted Lowery concept. (iii) According to this concept a base (like CaO, BaO or Na2O) is an oxide ion (O2–) donor and an acid (like SiO2, CO2 or P4O10) is an oxide ion (O2–) acceptor. Ex. Base Acid (a) CaO (b) 6Na2O + P4O10  4Na3PO4 (iv) Substances are termed amphoteric if they show a tendency of losing as well as + SiO2  CaSiO3 accepting an oxide ion. Ex. ZnO, Al2O3, BeO, Ga2O3 Lewis Concept (Electronic Concept)  An acid is a molecule/ion which can accept an electron pair with the formation of a coordinate bond. Acid e– pair acceptor Ex. Electron deficient molecules: BF3, AlCl3 Cations: H+, Fe2+, Na+ Molecules with vacant orbitals: SF4, PF3  A base is any molecule/ion which has a lone pair of electrons which can be donated. Base (One electron pair donate) Ex. Molecules with lone pairs: NH3, PH3, H2O, CH3OH 8