Compounds are broadly classified into acids, bases and salts. Many inorganic, biochemical reactions involve acids and bases. Many theories came to explain acids & bases. According to Arrhenius theory, acids give H+ ions, bases give OH- ions in their aqueous solutions. But this theory was failed to explain acidic behaviour of solutions
of CO₂, SO₃ and basic nature of solutions of NH₃, Na₂O etc. 
       Lowry - Bronstead introduced a new theory i.e. proton theory. According to them
Acid: A Substance that donates proton to another substance e.g. HCl, HNO₃.
Base: A Substance that accepts proton from another substance.
                             e.g.: NH₃, Cl^-, HSO₄ ^-
Conjugate acid - base pair: The acid-base pair which differs by a proton.
                             NH₃ + H₂O  NH₄⁺ + OH^-
                            H₂O & OH^-; NH₄⁺ & NH₃ are conjugate acid-base pairs.

Neutralisation: The transfer of a proton from acid to base.
e.g.: One proton of H₂O is transferred to NH₃ to form NH₄⁺
         A strong acid has a weak conjugate base and a strong base has a weak conjugate acid some compounds like H₂O, NH₃, ions like HCO₃^-, HSO₄^- can act as both acid or base, called amphiprotic.
        Acceptance or donation of protons takes place only in the presence of other substance.
This theory does not explain the behaviour of compounds like BF₃, BCl₃, AlCl₃. Water is levelling solvent. Water equalises H₃O⁺ concentration for acids, OH^- concentration for bases, is called 'levelling effect'. So relative strengths of acids or bases are studied in solvents like NH₃ or CH₃COOH. The order of strength of common acids in acetic acid is:
                           HClO₄ > HI > HBr > H₂SO₄ > HCl > HNO₃ > H₃PO₄ > H₂CO₃.   
  Acidic, basic behaviour is well explained by Lewis interms of electrons by his "Lewis theory".

According to Lewis theory
Acid: Any substance that accepts electron pair to form a coordinate covalent bond.
                                             eg: BF₃, BCl₃
Types of Lewis acids:
¤ All cations : CO⁺³, Fe⁺³
¤ Elements with electron sextet: O, S
¤ Molecules with multiple bonds: CO₂, SO₂
¤ Molecules with available d-orbitals: SF₄, SiF₄
¤ Molecules with incomplete octet: AlCl₃, BF₃
Base: Any substance that donates electron pair to form a coordinate covalent bond
                                               eg: H₂O, NH₃.
Types of Lewis bases:
¤ All anions: F^-, Cl^-
¤ Molecules with lone pairs: ROH, NH₃
¤ Molecules with multiple bonds: C₂H₂, C₂H₄
 All bronstead bases are also Lewis bases. For example NH₃ accepts a proton to form NH₄⁺ So it is bronstead base, NH₃ donate one electron pair so it is Lewis base. But all Lewis acids need not be Bronstead acids as they need not contain protons (eg: Fe⁺³, Cu⁺², SO₂)
    Lewis theory does not explain the strength of acids and bases. Reaction between HCl, H₂SO₄ and NaOH, KOH do not form any dative bond and the reaction is fast.
                       HClO₄ , H₂SO₄, HNO₃, HCl are strong acids.
    CH₃COOH, H₂CO₃ are weak acids. KOH, NaOH, CsOH are strong bases. NH₄OH, Cu(OH)₂ are weak bases. The strength of an acid or base can be known from their degree of ionization
(α = alpha). For strong acids or bases it is equal to 1. For weak acids or weak bases it is lessthan 1.

Ionization constant of weak acid (K_a)
      Let us suppose a weak acid HX is partially ionised in aqueous solution:
                             HX (aq.)  +  H₂O (l)     H₃O⁺(aq.)    +    X^- (aq.)
Initial concn.                 C                                    O                          O
Equilibrium concn.     c - cα                              cα                         cα

K [H₂O]²  =  K_w = [H⁺] [OH^-]
       The product of the concentration of H⁺ and OH^- ions in water or in aqueous solutions is known as ionic product of water. At 25°C, K_w is 1.0 × 10^-14 moles²/lit². Kw increases with the increase in temperature.

Significance of K_w:
¤ If [H⁺] or [OH^-] is known the other can be calculated from Kw
¤ For a neutral liquid, [H⁺] = [OH^-] =

 = 1.0 × 10^-7 moles/lit
¤ For acidic solutions [H⁺] > 10^-7 ( ... [H⁺] > [OH^-])
¤ For basic solutions [H⁺] < 10^-7 [... [H⁺] < [OH^-])
Dissociation constant of water (K):
               [H₂O]  H⁺  +  OH^- at 298 K
         
Concept of pH
¤ pH concept was introduced by S.P.L. Sorenson to express the acidity of a dilute solution accurately. "negative value of logarithm(to the base 10) of [H⁺] ion concentration (moles/lit) in a solution is known as pH".
              pH = -log₁₀ [H⁺]
              pOH = -log₁₀ [OH^-]            (Hint: p = -log 10)
              pH + pOH = 14 (at 298 K)
       Solutions can have pH from 0 to 14. Acidic solutions have pH from 0 to 7, that of basic solutions have 7 to 14 for neutral solutions pH = 7. pH = 0 means solution is strongly acidic
(e.g.: 1 N HCl). For more concentrated acidic solutions pH can be negative, for more concentrated basic solutions pH can be more than 14. If the pH of a solution remains constant, it is called buffer solution.

Buffer solution: The solution whose pH remains constant upon the addition of few ml of strong acid or strong base.
Acid buffer solution: The solution which is formed by mixing weak acid and its salt with a strong base
                           e.g.: CH₃COOH  +  CH₃COONa
Base buffer solution: The solution which is formed by mixing weak base and its salt with a strong acid.
                            e.g.: NH4OH  +  NH4Cl
Working of acid buffer: When few ml of strong acid is added to acid buffer
     (eg.: CH₃COOH + CH₃COONa), H⁺ ions of added acid react with CH₃COO^- ions to form a weak acid.
                          CH₃COO^- + H⁺→ CH₃COOH

      As there is no change in H+ concentration, pH of the buffer remains same. When few ml of strong base is added to the buffer, OH^- ions react with H⁺ ions of acid buffer to form H₂O.
                 CH₃ COOH + OH- → CH₃COO- + H₂O
      As there is no change in OH- concentration, pH of the buffer remains same.
Working of base buffer: When few ml of strong acid is added to base buffer.
                      (e.g.: NH₄OH + NH₄Cl), H+ ions of added acid reacts with OH^- ions of base to form weak ionised water. As there is no change in H⁺ ion concentration, pH of the buffer remains same.
                       NH₄OH + H⁺→ NH₄⁺ + H₂O

Buffer capacity of a buffer is maximum when pH = pK_a or pOH = pK_b [[salt] = [Acid] or [base]]                
A buffer solution shows its effective action only in the range of pH = pK_a ± 1
(e.g.: CH₃COOH + CH₃ COO Na has pK_a = 4.76. So it works well from pH 3.76 to 5.76). A buffer with greater  value is a good buffer.

Applications of buffer solutions:
¤ Buffers are used in softening of hard water.
¤ Buffers play an important role in biochemical reactions. Blood is a buffer
          (H₂CO₃ + NaHCO₃) with pH of 7.4
¤ Buffers are widely used in enzyme catalysed reactions, Chemical analysis and industrial synthetic processes.
Hydrolysis of Salts:
        When acid reacts with base, a salt is formed. When a salt is added to water, it gets hydrolysed. "Anion or Cation or both of a salt react with water to give OH^- ions or H⁺ ions or both is known as salt hydrolysis".
Hydrolysis of NH₄Cl:
      NH₄Cl salt is formed from a weak base NH₄OH and a strong acid HCl. So NH₄⁺ (conjugate ion) is strong acid, which undergoes cationic hydrolysis to give H⁺ ions.
         Nature of NH₄Cl solution is acidic. pH < 7
     Hydrolysis Constant (K_h) for NH₄Cl salt K_h=  
          pH = 7 -   [pK_b + log c]
Hydrolysis of CH₃COONa: CH₃COONa Salt is formed from a weak acid CH₃COOH and a strong base NaOH. So CH₃COO- is strong base, which undergoes anionic hydrolysis to give OH^- ions. Nature of CH₃COONa solution is alkaline. pH > 7.
                           K_h =   
                  pH = 7 +  

[pK_a + log c]
Hydrolysis of CH₃COO NH₄: CH₃COONH₄ Salt is formed from a weak acid CH₃COOH and a weak base NH₄OH. It involves both anionic as well as cationic hydrolysis. The nature of the solution may be neutral or acidic or basic depending upon the relative degrees of ionization of the weak acid and weak base.
                                      
                     pH = 7 +  [pK_a - pK_b]        or       pH =  [pK_w + pK_a - pK_b]

Hydrolysis of NaCl: NaCl is formed from strong base NaOH and strong acid HCl. So Na⁺ is weak acid and Cl^- is weak base both involves no hydrolysis. Nature of the solution is neutral and its pH = 7.
¤ Common ion effect and solubility product are very important Phenomena in analytical chemistry, purification of salts, and in gravimetric analysis etc.
Common ion Effect: The suppression of ionization of a weak electrolyte, when a solution (having common ion with weak electrolyte) of strong electrolyte is added to the solution of weak electrolyte.
pplications of Common ion effect:
¤ In precipitation purification of NaCl by passing HCl gas
¤ In controlling H+ ion concentration in buffer solution
¤ In Analytical Chemistry
¤ Precipitation of cations of groups in qualitative analysis.
¤ In grallimetric analysis.

Solubility: The amount of solute which can be dissolved by 100 grams of a Solvent at a given temperature. A Solute dissolves in a Solvent when hydration energy is greater than lattice energy. The solubility increases with increase of temperature, when the heat of solution is endothermic.
Solubility Product: The product of molar concentrations of cations and anions of salt
present in a saturated solution. Solubility product of a Salt Ax By can be explained as follows.
                                 A_xB_y  xA⁺ + yB^-
     Solubility Product (K_SP) = [A^y+]^x [B^x- ]^y

Applications of Solubility Product:
¤ In the precipitation of III group hydroxides, sulphides of II and IV group cations.
¤ In the manufacture of sodium bicarbonate.
solubility product helps in predicting precipitation on mixing two solutions. If the ionic product is greater than solubility product, precipitation takes place, If it is less than that of K_SP, no precipitate is formed. If ionic product is exactly equal to K_SP, precipitation just starts.