Concepts of Acids and Bases

Chemistry Acids, Bases& Salts Concepts of Acids and Bases Arrhenius Theory: This concept was presented in 1884. According to this theory all substances which give H+ ions when dissolved in water are called acids, while those which ionise to give OH- ions are called bases. The H+ ions do not exist as such and exist in combination with molecules of H2O as H3O+ ions (Known as hydronium ion). H+ + H2O H3O+ HCl + H2O H3O+ + Cl– e.g. HA + H2O H3O+ + A– Acid H2SO4 + 2H2O 2H3O+ + SO4–2 Acid BOH + H2O B+ + OH– Base NaOH + H2O Na+ + OH– Base NH4OH NH4+ + OH– Base A) Limitations of Arrhenius theory: The concept does not explain the acidic or basic properties of acids or bases in nonaqueous solvents. It fails to explain acidic nature of non - protic compounds like SO2, P2O5, CO2, NO2 etc. Which do not have hydrogen in their molecules to furnish H+ ions. It fails to explain acidic nature of certain salts like AlCl3 etc., in aqueous solutions. 1 Chemistry Example: Which of the following is not the characteristics of an acid– (a) Turns blue litmus to red (b) Turns phenolphthalein pink from colourless (c) Decompose carbonates (d) Oxy compounds of non–metals Solution : Correct answer is (B) Refer (1.1) preliminary classification of acid & bases. Option (b) indicates characteristic of bases. Example : Arrhenius theory of acid–base is not applicable in (a) aqueous solution (b) In presence of water (c) Non–aqueous solutions (d) None of the above Solution : Correct answer is (C), since Arrhenius theory is only applicable to aqueous media. B) Acid Base Concept of Bronsted and Lowry : The theory was given by Bronsted, a Danish chemist and Lowry, an English chemist independently in 1923. According to it an acid is a substance, molecule or ion which has a tendency to release the proton (protogenic) and similarly a base has a tendency to accept the proton (protogenic). e.g. HCl + H2O H3O+ + Cl– In the reaction, HCl acts an as acid because it donates a proton to the water molecule. Water, on the other hand, behaves as a base by accepting a proton. 2 Chemistry Bronsted and lowry theory is also known as proton donor and proton acceptor theory. Other examples: CH3COOH + H2O H3O+ + CH3COO– NH4+ + H2O H3O+ + NH3 NH3 + H2O NH4+ + OH– In the reaction (i) and (ii) water is acting as a base, while in reaction (iii) it is acting as an acid. Thus water can donate as well as accept H+ and can act as both acid and base. Example : Suitable reason (s) for higher strength of an acid or base will be – (a) Higher value of Ka or Kb (b) Higher extent of ionisation (c) (a) and (b) both (d) Larger number of replaceable H atoms. Solution : Correct answer is (C), both Ka or Kb and degree of ionisation are the measure of strength of an acid or base. Example : The basicity of phosphorous acid is – (a) 1 (b) 2 (c) 3 (d) 4 Solution : Correct answer is (B) since Phosphorous acid is H3PO3 which has two replaceable H+ ions. The species like H2O, NH3, CH3COOH which can act as both acid and base are called amphiprotic. Moreover according to theory, an acid on losing a proton becomes a base, called conjugate base, while the base by accepting proton changes to acid called conjugate acid. Here CH3COO– ion is conjugate base of CH3COOH–, while H3O+ ion is conjugate acid of H2O. 3 Chemistry     some common conjugate List of Acid–Base pairs: 4 Chemistry Classification of Bronsted Lowry Acid & Bases: Type Molecular Cationic Anionic Acid HCl, HNO3, HClO4, H2SO4, H3PO4, CH3COOH, HBr, H2O etc. NH  4 , N 2 H 5 , PH  4 , + +2 Na , Ba (All cations) [Fe(H2O)6]3+, [Al(H2O)6]+3, [Al(H2O)6]3+, –  3 HS , HSO , H 2SO  4 , HSO  4 , HCO3 , 2 4 HPO , all amphiprotic anions 5 Base NH3, N2H4, Amines, H2O, Alcohol, Ethers [Fe(H2O)5 OH]2+ [Al(H2O)5 OH]2+ Cl–, Br–, OH–, HSO  4 CH–, CO32  , SO 24  , NH  2 , CH3COO H– Chemistry 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 nonaqueous solvents (such as Glacial acetic acid, Hydrogen halides, SO2 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 important in deciding its be haviour. 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, Na2O) is an oxide ion (O2–) donor and an acid (like SiO2, CO2 or P4O10) is an oxide ion (O2–) acceptor. e.g. Base Acid (a) CaO + SiO2 CaSiO3 (b) CaO + P4O10 4Na3PO4 (iv) Substances are termed amphoteric if they show a tendency of losing as well as accepting an oxide ion. e.g. ZnO, Al2O3 Example : The aprotic solvent is – (a) H2O (b) C6H6 (c) HF Solution : Correct Answer is (b) 6 (d) NH3 Chemistry Example : Which of the following is the strongest conjugate base (a) Cl– (b) CH3COO– (c) kk (d) kl+ Solution : Correct answer is (b), since CH3COOH is weakest acid among HCl, CH3COOH, and HNO2 Example : The strongest Bronsted base in the following anions is (a) ClO– (b) (c) (d) Solution : Correct answer is (a), since HClO is weakest acid among HClO, HClO2, HClO3 and HClO4. Example : Urea can show weak acidic nature in liquid NH3 solution . Write appropriate equation and label as acid & base. Solution : NH2CONH2 + NH3 + NH2CONH– Urea Acid base (c) Lewis theory : The theory was given by G.N.Lewis in 1938. According to it, an acid is a species which can accept a pair of electrons, while the base is one which can donate a pair of electrons. It is also known as electron pair donor and electron pair acceptor theory. e.g. (i) FeCl3 and AlCl3 are Lewis acids, because the central atoms have only six electrons after sharing and need to have more electrons. (ii) NH3 is a Lewis base as it has a pair of electrons which can be easily donated. Lewis acids:- CH3+, H+, BF3, AlCl3, FeCl3, etc. 7 Chemistry Lewis bases :- NH3, H2O, R-O-R, R - OH, CN-, OH- etc. (i) Characteristics of species which can act as Lewis acid :  Molecules with vacant orbitals : Lewis acids are electron deficient molecules such as BF3, AlCl3, GaCl3 etc.  Molecules in which the central atom has incomplete octet : The central atom of the halides such as TiCl4, SnCl4, PCl3, PF5, SF4, TeCl4. etc., have vacant d-orbitals. These can, therefore, accept an electron pair and act as Lewis acids.  Simple cation : All cations are expected to act as Lewis acid, since they are electron deficient in nature. .. Ag + 2 N H 3   H 3 N  Ag  NH 3  + + Fe2+ + 6CN -   Fe  CN 6  4–  Molecules having a multiple bond between atoms of dissimilar electronegativity : Typical examples of molecules belonging to this class of acids are CO2, SO2, and SO3. (ii) Characteristics of species which can act as Lewis bases :  Netural species having at least one lone pair of electrons : For example, ammonia amines, alcohols etc., act as Lewis bases as they contain a pair of electrons. .. : NH3 , R  N H2 , R   H, H   H, R   R  Negatively charged species or simple anions: For example chloride (Cl-),cyanide (CN-) 8 Chemistry hydroxide (OH-) ions etc. act as Lewis bases.  Multiple bonded compounds: The compounds such as CO, NO, Ethylene, acetylene etc. can act as Lewis bases. 9