1.2.1 BRÖNSTED – LOWRY DEFINITIONS

The Brönsted−Lowry definition of acids and bases is based upon the transfer of one or more hydronium ions (H⁺) from an acid to a base. In 1923, Brönsted and Lowry independently defined as follows.

Acids are those molecules or ions which can donate one or more hydronium ions to compounds called bases. Example, hydrochloric acid (HCl), sulphuric acid (H₂SO₄) phosphoric acid (H₃PO₄) or even water (H₂O) and ammonium ions:

HCl ⇌  H⁺ + Cl⁻
H₂SO₄ ⇌  2H⁺ + SO₄²⁻
H₃PO₄ ⇌  3H⁺ + PO₄³⁻
H₂O  ⇌  H⁺ + OH⁻
NH₄⁺  ⇌  H⁺ + NH₃

Note that all Arrhenius acids are Brönsted acids, but all the Brönsted acids may not be Arrhenius acids.

Bases are those molecules or ions which can accept one or more hydronium ions from acids. For example, ammonia (NH₃), water (H₂O), hydroxide ( OH⁻), or anions like acetate (CH₃COO⁻), sulphide (S²⁻), hydrogensulphide (HS⁻), etc.:

NH₃ + H⁺ ⇌  NH₄⁺
H₂O + H⁺ ⇌  H₃O⁺
OH⁻ + H⁺ ⇌  H₂O
CH₃COO⁻  + H  ⇌ CH₃COOH
S²⁻ + H⁺ ⇌  HS⁻ ;  HS⁻ + H⁺ ⇌  H₂S (two steps)

Note that all Arrhenius bases are Brönsted bases, but all the Brönsted bases may not be Arrhenius bases

1.2.2 CONJUGATE ACIDS AND BASES – NEUTRALIZATION REACTIONS OF BRÖNSTED ACID AND BASES

 Neutralization is defined as a reaction in which a hydronium ion is transferred from a Brönsted acid HA to a Brönsted base B resulting in the formation of another Brönsted acid BH⁺ and another Brönsted base A⁻, called the conjugate base and conjugate acids of the original acid and base respectively:

 HA  +  B  ⇌  A⁻  +  BH⁺  (Neutralization of acid HA with base B)

For a complete neutralization of a polyprotic acids (containing more than one ionizable H⁺ ions) or a polyprotic base (capable of accepting more than one H⁺ ions), more than one hydronium ions must get transferred from the acid to the base.

In the above neutralization reaction, the products formed are related to and differ from the original  acid and base by a hydronium ion (H⁺) only. By definition, the species that differ from one another by one hydronium ion only (like A ⁻ and HA, or B and BH⁺) are called the conjugate base  and conjugate acid respectively. If we write the acid species first, the conjugate acid–base  pairs formed in the above reaction are  HA – A^– and BH⁺– B . Some examples of neutralization reactions are given in Table 1.2.1.

Table 1.2.1 Conjugate acid – base pairs and neutralization reactions

Acid       +             Base       ⇌     Conjugated Base      +  Conjugated Acid
  HA         +              B           ⇌                           A⁻            +             BH⁺
HCl        +             OH^–       ⇌                           Cl^–          +             H₂O
 HCl        +              NH₃       ⇌                           Cl^–          +             NH₄⁺
HX         +               CN ⁻     ⇌                           X^–           +              HCN
H₂PO₄^– +             OH⁻       ⇌                           HPO₄^2–   +             H₂O
H₂SO₄    +             Cl ⁻         ⇌                           HSO₄^–    +             HCl

Some common conjugated acid-base pairs are given in Table 1.2.2,

Table 1.2.2 Common conjugated acid–base pairs

Acid (HA or BH⁺)     Conjugate Base (A^– or B)      Strength*
Hydrochloric acid (HCl)       Chloride (Cl)          Strong – Weak
Nitric acid (HNO­₃)                  Nitrate (NO₃^– )        Strong – Weak
Sulphuric acid (H₂SO₄)     Hydrogensulphate (HSO₄^–)  Strong – Weak
Hydrogensulphate (HSO_4^–)  Sulphate (SO₄^2–)      Strong –  Weak
Acetic acid (CH₃COOH)       Acetate (CH₃COO^–)     Weak – Strong
Hydrogen cyanide (HCN)    Cyanide (CN ⁻)           Weak – Strong ^––
Water (H₂O)                           Hydroxide (OH^–)       Weak – Strong
Hydronium ion (H₃O⁺)        Water (H₂O)              Strong – Weak
Ammonium ion (NH₄⁺)       Ammonia (NH₃)     Strong – Weak
Ammonia (NH₃)                   Amide (NH₂^–)         Weak – Strong
Carbonic acid (H₂CO₃)   Hydrogencarbonate (HCO₃^–)  Weak – Weak
Hydrogen carbonate (HCO₃^–) Carbonate (CO₃^2–)          Weak – Weak
Phosphoric acid (H₃PO₄)     Dihydrogenphosphate (H₂PO₄^–)  W-Weak
Dihydrogenphosphate(H₂PO₄^2–) Hydrogenphosphate(HPO₄^2–)W-Wk
Hydrogenphosphate (H₂PO₄^–)  Phosphate (PO₄^3–)  Weak – Strong

*Strength of acids and bases are discussed later.

The usefulness of Brönsted–Lowry concept is that:

1. Any protonic solvent like ammonia or sulphuric acid can be handled :

NH₃        +           NH₃                  ⇌           NH₄⁺      +             NH₂^–
H₂SO₄             +       H₂SO₄        ⇌               H₃SO₄⁺  +            HSO₄^–
(Conjugated Base   +    Conjugated Acid    ⇌   Acid  +   Base)

2. All hydronium ion transfer reactions (not usually called acid–base reactions) can be treated as such:

NH₄⁺  + S^2–  →   NH₃  +   HS^–
Acid₁        Base₂         Base₁        Acid₂

However, Brönsted–Lowry definition fails for non-hydronic (non-protonic) solutions like SO₂, POCl₃, BrF₃, etc.

1.2.3 CONJUGATED ACID–BASE PAIRS IN POLYPROTIC ACIDS

In case of polyprotic acid−base system, it is important to remember that the conjugated acid−base pairs differ by one hydronium ion only. Take for example, phosphoric acid (H₃PO₄). It can donate three hydronium ions successively to a base and form anions H₂PO₄^– (dihydrogenphosphate), HPO­₄^2– (hydrogenphosphate) and phosphate (PO₄^3–) respectively. In this system, there are three  conjugate acid–base pairs as shown below:

               H₃PO₄ – H₂PO₄^–,        H₂PO₄^– – HPO₄^2–  ,    and         HPO₄^2– – PO­₄^3–

The other acid−base pairs like H₃PO₄ – HPO₄^2– H₃PO₄ – PO₄^3–  or H₂PO₄^– –  PO₄^3– do not form conjugated acid – base pairs as they differ by two or three hydronium ions.

Similarly, for carbonic acid (H₂CO₃), the conjugated acid−base pairs are H₂CO₃ – HCO₃^– and HCO₃^– – CO₃^2–; the carbonic acid (H₂CO₃) – carbonate ion (CO₃^2–) is not a  conjugated acid – base pair.

1.2.3.1 Different types of Acids and Bases

Acids: The different types of monobasic  acids are  the following.

1. Neutral molecules like hydrochloric acid (HCl), nitric acid (HNO₃), acetic acid (CH₃COOH) or hydrogen cyanide (HCN), which ionize to form only one anion (chloride, nitrate, acetate and cyanide, respectively):

HCl(aq) ⇌  H⁺(aq) + Cl ⁻(aq)
HNO₃(aq) ⇌ H⁺(aq) + NO­₃ ⁻(aq)
CH₃CO₂H(aq) ⇌ H⁺ (aq) + CH₃CO₂⁻(aq)
HCN(aq) ⇌  H⁺(aq) + CN ⁻(aq)

Therefore, they can donate only one hydronium ion to a base, and their neutralization reactions can be represented as (monobasic means reacting with one equivalent of a base):

HCl (aq) + NaOH (aq) → NaCl (aq) + H_2­O(l)
HNO₃ (aq) + KOH (aq) → KNO₃ (aq) + H₂O(l)
CH₃COOH (aq) + NH₃ (aq) → CH₃COONH₄ (aq)

2. Cations like ammonium ion (NH­₄⁺), which can donate a hydronium ion (H⁺) to other bases:

NH₄⁺(aq) ⇌ H⁺ (aq) + NH₃(aq)

Bases: A base B is defined as a monoprotic base if one molecule of B can accept  only one H⁺ ion and get neutralized completely. It forms only one protonated species [BH]⁺, which may be a cation or a neutral molecule. (A protonated species is formed by the covalent attachment of one or more hydrons to another species.). Some examples are given below.

3. Hydroxides of monovalent cations e.g., sodium hydroxide (NaOH), potassium hydroxide (KOH), thallium(I) hydroxide (TlOH) or ammonium hydroxide (NH₄OH),
4. Molecules containing a lone pair of electrons which can be donated to a hydronium ion, e.g., water, ammonia (NH₃), organic mono amines [ethyl amine (CH₃CH₂­NH₂), aniline (C₆H₅NH­₂), etc.
5. All anions having one unit of negative charge, such as acetate (CH₃COO ⁻), cyanide (CN ⁻), halides (X ⁻), hydrogensulphate (HSO₄⁻), etc.
   

   1.2.4 ACIDS

   1.2.4.1 Monoprotic or Monobasic Acids

   A monoprotic or monobasic Brönsted acid can be represented as HA and ionizes to give only one H⁺ ion and one anion A^–. As one molecule of a monobasic acid ionizes to give only one hydronium ion, it can form only one series of salts which contain the anion A^–. Examples include HCl, nitric acid, acetic acid, HCN, formic acid HCOOH, etc.

   1.2.4.2 Diprotic or Dibasic Acids

   A diprotic or dibasic acid can be represented as H₂A. They can donate, or ionize stepwise to give two hydronium ions per molecule and form two distinct series of salts containing the anions HA⁻ and A²⁻. For example sulphuric acid (H₂SO₄), which ionizes stepwise in aqueous solutions to give two H⁺(aq) ions per molecule and forms two series of salts containing hydrogensulphate (HSO₄⁻) and sulphate (SO₄²⁺) ions:

               H₂SO₄ (aq) ⇌ H⁺ (aq) +  HSO₄⁻(aq)
   H₂SO₄ (aq) + NaOH(aq) → Na⁺(aq) +  HSO₄⁻(aq) + H₂O(l)

          SO₄⁻(aq) ⇌ H⁺ (aq)  + SO­₄^{2 −} (aq)
   HSO₄⁻(aq) + NaOH(aq) → Na⁺(aq)  + SO­₄^{2 −}(aq) + H₂O(l)

   Overall ionization reaction and neutralization reaction can be written as

                  H₂SO₄ (aq) ⇌ 2H⁺ (aq) +  SO₄²⁻ (aq)
   H₂SO₄ (aq) + 2NaOH(aq) → 2Na⁺(aq) +  SO₄²⁻(aq) + 2 H₂O(l)

    Oxalic acid [(COOH)₂], succinic acid [(CH₂COOH)₂]  and tartaric acid [(CHOHCOOH)₂, or C₄H₆O₆) are some other common dibasic acids.

   A dibasic acid can be a neutral molecule (see above) or a cation of a diacidic base, e.g., hydrazinediium ion (N₂H₆²⁺):

    (N₂H₆²⁺)(aq) ⇌ 2H⁺(aq) + N­₂H₄(aq)

   1.2.4.3 Triprotic or tribasic acids

   A triprotic or tribasic acid can be represented as H₃A. They can donate or ionize stepwise, to give three hydronium ions per molecule and form three distinct series of salts containing the anions H₂A⁻ and HA²⁻ and A³⁻. For example, phosphoric (H₃PO₄), which ionizes stepwise in aqueous solutions to give three H⁺ (aq) ions per molecule and forms three series of salts containing dihyrogenphosphate (H₂PO₄⁻) and hydrogenphosphate (HPO₄²⁻) and phosphate (PO₄ ³⁻) ions:

                  H₃PO₄ (aq) ⇌ H⁺ (aq) +  H₂PO₄⁻(aq)
   H₃PO₄ (aq) + NaOH(aq) → Na⁺(aq) +  H₂PO₄⁻(aq) + H₃O⁺(aq)

                  H₂PO₄⁻(aq) ⇌ H⁺ (aq)  + HPO­₄^{2 −} (aq)
   H₂PO₄⁻(aq) + NaOH(aq) → Na⁺(aq)  + HPO­₄^{2 −}(aq) + H₃O⁺(aq)

        HPO₄²⁻ (aq) ⇌ H⁺ (aq)  + PO­₄^{3 −} (aq)
   HPO₄²⁻ (aq) + NaOH(aq) → Na⁺(aq)  + PO­₄^{3 −}(aq) + H₃O⁺(aq)

   Overall ionization reaction and neutralization reaction can be written as

                  H₃PO₄ (aq) ⇌ 3H⁺(aq) +  PO₄³⁻(aq)
   H₃PO₄ (aq) + 3NaOH(aq) → Na₃PO₄(aq) + 3H₂O(l)

   1.2.4.4 Polyprotic or Polybasic Acids

   A polyprotic acid can be represented as H_nA. In aqueous solutions, it ionizes to or donates more than one hydronium ion [H⁺(aq)], stepwise, to a base. Based on the value of  n,  these acids  are further classified as diprotic or dibasic acids, triprotic or tribasic acids and polyprotic acids.

1.2.5 BASES

 1.2.5.1 Monoacidic or Monoprotic Bases

Neutralization Reactions of Monoprotic Bases: A monoprotic base can accept only one hydronium ion (H⁺). The neutralization of one mole of a monoacidic or monoprotic base requires one mole of a monobasic acid (HCl):

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)
CH₃COONa  + HCl(aq) → CH₃COOH(aq) + NaCl(aq)
NH₃(aq) + HCl(aq) → NH₄⁺(aq) + Cl⁻(aq)

When one mole of a monoacid base is treated with a monobasic acid like hydrochloric acid (HCl), the base gets completely neutralized  by one mole of hydrochloric acid:

NaOH + HCl → NaCl + H₂O
 KOH + HCl → KCl + H₂O
 NH₄OH + HCl → NH₄Cl + H₂O

When a neutral molecule like ammonia or water accepts a hydronium ion, it forms the cation NH₄ ⁺ (ammonium) or  H₃O ⁺ ( hydrated hydronium ion) respectively:

 B(aq) + H⁺(aq)  ⇌ BH ⁺(aq)
 NH₃(aq) + H⁺(aq)  ⇌ NH₄⁺(aq)
 H₂O(l) + H⁺(aq) ⇌ H₃O⁺(aq)

When an anion A⁻ carrying one unit of negative charge, like hydroxide ion (OH⁻), or acetate (CH₃COO⁻)  accepts  a hydronium ion, a neutral molecule is formed:

OH⁻(aq) + H⁺(aq) ⇌ H₂O(l)
CH₃COO⁻ + H⁺(aq) ⇌  CH₃COOH
A⁻ (aq) + H⁺(aq) ⇌ HA(aq)

1.2.5.2 Diprotic or Diacidic Bases

A molecule of a polyprotic bases (B) can accept more than one hydronium ion  and form more than one hydronated species (BH_nⁿ⁺), Bases\d on the charge on the base, the hydronated species may be a cation, a neutral molecule or even an anion. Based on the value of n, a polyprotic base can be further classified as follows.

Diprotic or Diacidic Base: A diprotic or a diacidic base B can accept two hydronium ions [H⁺(aq)] per molecule and form  salts containing BH⁺ and BH₂²⁺ ions. A diprotic or diacidic base can be:

1. Hydroxide of divalent metal cation, e.g., calcium hydroxide (Ca(OH)₂) or barium hydroxide (Ba(OH)₂).
2. Organic compounds containing two basics groups, e.g., ethylene diamine (NH₂CH₂CH₂NH₂), p-phenylene diamine or 1,4-diaminobenzene (C₆H₄(NH₂)₂).
3. All A²⁻ anions (having two units of negative charge), like sulphide S²⁻, carbonate (CO₃²⁻), etc.

The typical reactions of diprotic or diacidic bases with a monobasic acid like hydrochloric acid are given below.

Ca(OH)₂ + 2HCl  → CaCl₂ + 2H₂O
Ba(OH)₂ + 2HCl → BaCl₂ + 2H₂O
NH₂CH₂CH₂NH₂ + 2HCl → (NH₃CH₂CH₂NH₃]²⁺  +  2Cl ⁻
 HS ⁻ (aq) + H⁺(aq) ⇌ H₂S (aq)
 CO₃^{2 −} (aq) + H⁺(aq) ⇌ HCO₃ ⁻ (aq)
HCO₃ (aq) + H⁺(aq) ⇌ H₂CO₃ (aq)

Note that the anion HS ^– is named as hydrogensulphide (as one word) whereas H₂S is named as hydrogen sulphide (two words).

1.2.5.3 Polyprotic  or Polyacid Bases

The above definition of a diprotic base can be extended further. Thus, a triprotic base B can accept, stepwise, three protons per molecule and form the ultimate species BH³⁺. Examples of triprotic bases include the following.

1. Hydroxides of trivalent metals like iron(III) hydroxide, Fe(OH)₃, or aluminium hydroxide, Al(OH)­₃.
2. Organic compounds containing three basic amino groups, e.g., 1,2,3-triaminopropane.
3. Anions having three units of negative charge, e.g., phosphate, PO₄^{3 −}.

The phosphate ion being a triprotic Brönsted base, can accept three H⁺ ions successively and stepwise, and forms     hydrogenphosphate (HPO₄²⁻), dihydrogenphosphate (H₂PO₄⁻) and phosphoric acid (H₃PO₄):

 H⁺ (aq)  + PO­₄^{3 −} (aq) ⇌ HPO₄²⁻ (aq)
H⁺ (aq)  + HPO­₄^{2 −} (aq) ⇌ H₂PO₄⁻(aq)
H⁺ (aq) +  H₂PO₄⁻(aq) ⇌ H₃PO₄ (aq)

Overall complete protonation reaction of phosphate ions with three H⁺(aq) ions can be written as:

PO₄^{3 −} (aq) + 3H⁺(aq) ⇌ H₃PO₄ (aq)

1.2.6 AMPHOTERIC COMPOUNDS

Amphoteric compounds are those compounds which react or get neutralized with both acids as well as bases and, in both the cases, form a salt and water. Therefore, an amphoteric compound can function both as an acid as well as as a base. Thus, zinc hydroxide (Zn(OH)₂) is an amphoteric hydroxide. It reacts with acids like hydrochloric acid or sulphuric acid and gets neutralized to form salts (zinc chloride or zinc sulphate) and water. On the other hand, it reacts with a base like sodium hydroxide,  and gets neutralized to form another salt, sodium zincate (Na₂ZnO₂),  and water. The reactions are:

 Acid       +   Base      →      Salt    +  Water
2HCl      +  Zn(OH)₂   →   ZnCl₂  +  2H₂O
  Zn(OH)₂ +  2NaOH  →  Na₂ZnO₂ + 2H₂O

Some other common amphoteric compounds are oxides and hydroxides of groups 3 to 16 e.g., ZnO, CuO, NiO, Fe₂O₃, Al₂O₃, Al(OH)₃, PbO, Pb( OH)₂, SnO and Sn(OH)₂ etc.

 

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