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Arrhenius theory of Acids & Bases

According to Arrhenius concept of acid and bases,

1. Acid

  • An acid is a substance that releases H+ ions on dissolving in water.
  • Greater the number of H+ ions produced in the solution, the stronger is the acid.

2. Bases

  • A base is a substance that releases OHions on dissolving in water.
  • Similarly greater the number of OHions produced in the aqueous solution, the stronger is the base.

Acids, Bases & Salts | Chemistry Class 11 - NEET

Bronsted Lowry Theory of Acids and Bases

  • A Bronsted-Lowry acid must contain at least one hydrogen atom, typically attached to a very electronegative atom such as oxygen.
  • When a Brønsted-Lowry acid donates a proton, it loses one H from its molecular formula, and its charge decreases by one.
  • The resulting species is a base since the reversal of the loss of proton is the gain of a proton.
  • The species gaining the proton is acting as a proton acceptor or a Brønsted-Lowry base.
  • Furthermore, when an acidic substance loses a proton, it forms a base, called the conjugate base of the acid. When a basic substance gains a proton, it forms an acid called the conjugate acid of a base.

  • The strength of a base is measured by its ability to capture protons. In contrast, the strength of an acid is measured by its ability to donate protons.

  • Therefore, the stronger the acid, the weaker the conjugate base, and the stronger the base, the weaker is its conjugate acid.

Acids, Bases & Salts | Chemistry Class 11 - NEET

Acids, Bases & Salts | Chemistry Class 11 - NEET

Lewis Theory of Acid & Bases

Lewis Acid

  • Lewis Acids are the chemical species that have empty orbitals and are able to accept electron pairs from Lewis bases. 
  • This term was classically used to describe chemical species with a trigonal planar structure and an empty p-orbital. An example of such a Lewis acid would be BR3 (where R can be a halide or an organic substituent).
  • Water and some other compounds are considered as both Lewis acids and bases since they can accept and donate electron pairs based on the reaction.
  • All Bronsted-Lowry’s acids are Lewis acids while acids need not be Bronsted-Lowry’s acids.

Acids, Bases & Salts | Chemistry Class 11 - NEET

Examples of Lewis Acids

Some common examples of Lewis acids that can accept electron pairs include:

  1. H+ ions (or protons) can be considered as Lewis acids along with onium ions like H3O+.
  2. The cations of d block elements that display high oxidation states can act as electron-pair acceptors. An example of such a cation is Fe3+.
  3. Cations of metals such as Mg2+ and Li+ can form coordination compounds with water acting as the ligand. These aqua complexes can accept electron pairs and behave as Lewis acids.Acids, Bases & Salts | Chemistry Class 11 - NEET
  4. Carbocations given by H3C+ and other trigonal planar species tend to accept electron pairs.
  5. The Pentahalides of the following group 15 elements can act as Lewis acids – Antimony, Arsenic, and Phosphorus.
  6. Apart from these chemical compounds listed above, any electron-deficient π system can act as an acceptor of electron pairs – enones, for example.

Lewis Base

  • Atomic or molecular chemical species having a highly localized HOMO (The Highest Occupied Molecular Orbital) act as Lewis bases. 
  • These chemical species have the ability to donate an electron pair to a given Lewis acid in order to form an adduct, as discussed earlier.
  • The most common Lewis bases are ammonia, alkyl amines, and other conventional amines. 
  • Commonly, Lewis bases are anionic in nature and their base strength generally depends on the pKa of the corresponding parent acid. 
  • Since Lewis bases are electron-rich species that have the ability to donate electron-pairs, they can be classified as nucleophiles. 
  • Similarly, Lewis acids can be classified as electrophiles (since they behave as electron-pair acceptors).

Examples of Lewis Bases

  1. Examples of Lewis bases which have an ability to donate an electron pair are listed below.
  2. Pyridine and the derivatives of pyridine have the ability to act as electron pair donors. Thus, these compounds can be classified as Lewis bases.
  3. The compounds in which Oxygen, Sulphur, Selenium, and Tellurium (which belong to group 16 of the Periodic Table) exhibit an oxidation state of -2 are generally Lewis bases. 
  4. Examples of such compounds include water and ketones.
    Acids, Bases & Salts | Chemistry Class 11 - NEET
  5. The simple anions which have an electron pair can also act as Lewis bases by donating these electrons. Examples of such anions include H– and F–. Even some complex anions, such as the sulfate anion (SO42-) can donate pairs of electrons.
  6. The π-systems which are rich in electrons (such as benzene, ethyne, and ethene) exhibit great electron pair donating capabilities.
  7. Weak Lewis acids have strong conjugate Lewis bases. 
  8. Apart from this, many chemical species having a lone pair of electrons such as CH3– and OH– are identified as Lewis bases due to their electron pair donating capabilities.

Limitations of Lewis Concept

  1. It does not explain the behaviour of protonic acids such as HCl, H2SO4, HNO3 etc.
  2. It does not predict the magnitude of relative strength of acids and bases.

Basic Strength

Acids, Bases & Salts | Chemistry Class 11 - NEET

Basic strength directly depends on the availability of lone pair for H. 

Comparison of Basicity of Ammonia and Alkyl Amines 

Compare the basic strength of the following NH3, CH3NH2, (CH3)2NH, (CH3)3N

  • Factors which affect the basicity of Amines

1. steric effect 

2. Inductive effect 

3. solvation effect.

  • The base whose conjugate acid is more stable will be more acidic forming conjugate acid of the given basAcids, Bases & Salts | Chemistry Class 11 - NEET
  • Stability order of conjugate acid 

       Acids, Bases & Salts | Chemistry Class 11 - NEET

  • Therefore basic strength, (CH3)3N > (CH3)2NH > CH3NH2 > NH3
  • In Aqueous solution or in polar solvent  (CH3)2NH > CH3NH2 > (CH3)3N > NH3
  • In aqueous solution, the conjugate acids form H-bonds (intermolecular) with water molecules and stabilise themselves. Conjugate acid of 1° amine which has largest no. of H-atoms form maximum H-bond with water and is most stable. Consequently 1° amine is most basic.
  • Due to steric effect 1° amine is considered more basic as compared to 3° amine as lone pair is hindered by three alkyl group and less available for H .
  • Considering the combined effect of the three (Inductive, solvation and steric effect) we can conclude that 2° > 1° > 3° > NH3
  • Aromatic amines are least basic as their lone pair is in conjugation and less available for protonation.

Levelling Effect

The leveling effect is a concept in acid-base chemistry where strong acids or bases in a particular solvent appear to have the same strength due to the solvent's limiting properties. In water, any acid stronger than the hydronium ion (H3O+) donates a proton to water and only produces H3Oions. This means that acids like hydrochloric acid (HCl) , sulfuric acid (H2SO4), and nitric acid (HNO3) all seem equally strong in water, even though they have different intrinsic acid strengths. Similarly, bases stronger than hydroxide ion (OH-) are "leveled" to the strength of OH-OH in water.

The leveling effect depends on the solvent: in a different solvent, such as acetic acid, acids and bases may not exhibit the same strength as they do in water. This concept is essential for understanding why certain reactions may behave differently in various solvents and helps explain the limits of acid and base strength in specific environments.

The pH Scale

pH is defined as the negative logarithm of hydrogen ion concentration.

  • pH = – log [H+] and conversely [H+] = lO-pH
  • Total [H+] or [OH] in a mixture of two strong acids or bases = (ΣNV/ΣV)
  • Similarly, negative logarithm of hydroxyl ion concentration is pOH.

Acids, Bases & Salts | Chemistry Class 11 - NEET

  • pH value of an acid having H+ concentration less than 10-7, is always in between 6 and 7. For 10-8 N HCl solution. it is 6.958.Similarly, for 10-8 NaOH solution, the pH is 7.04 (because basic solutions always have pH 7.)
  • pH of solution is accurately measured by pH meter or emf method or roughly by pH paper or indicator paper.
  • pH values can be zero or even negative in highly concentrated acidic solutions. For example, a 1 N HCl solution has a pH of 0, while more concentrated solutions like 2 N, 3 N, or 10 N can have negative pH values.

Acids, Bases & Salts | Chemistry Class 11 - NEET

Dissociation of Acids & Bases

Dissociation of weak electrolytes, (of weak acids and weak bases)
Weak electrolytes are ionized to only a slight extent, the concentration of ions are relatively low; most solute molecules do not split into ions.
The extent to which the weak electrolytes dissociates is expressed by either ionization constants or degree of ionization.
A weak monoprotic acid
HA + H2O ⇄ H3O+ + A-
A weak base
B + H2O ⇄ BH+ + OH-
Equilibrium constants of ionization
Acids, Bases & Salts | Chemistry Class 11 - NEET

Acids, Bases & Salts | Chemistry Class 11 - NEET

Acid and base ionization constants
Acids, Bases & Salts | Chemistry Class 11 - NEET

Buffer Solution

Solution which resists the change in its pH value by addition of a small amount of acid or a base, is called buffer solution.

Acidic Buffer: It has a pH < 7.

Example: CH3COOH/CH3COONa, Boric acid/Borax.

Basic Buffer: Has a pH > 7.

Example: NHNH4OH/NH4Cl.

Buffer System in Blood: The buffer system present in blood is HNaHCO
\text{H}_2\text{CO}_3 + \text{NaHCO}_3
H2CO3+NaHCO3.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation is used to calculate the pH of a buffer solution.

For an Acidic Buffer:Acids, Bases & Salts | Chemistry Class 11 - NEET

For a Basic Buffer: Acids, Bases & Salts | Chemistry Class 11 - NEET

and pH = 14 - pOH

Where:

  • pKa = - log Ka
  • pKb = - log Kb
  • Ka and Kb are the dissociation constants of the acid and base, respectively.
  • [Salt], [Acid] and [Base] represent the molar concentrations of the salt, acid, and base.

Buffer Capacity

It is defined as the number of moles of acid or base added to 1 L of a buffer solution to change the pH by one unit. It is denoted by (ϕ).

Acids, Bases & Salts | Chemistry Class 11 - NEET

For an Acidic Buffer Mixture (Henderson-Hasselbalch Equation)

For an acidic buffer solution containing a weak acid and its conjugate base:Acids, Bases & Salts | Chemistry Class 11 - NEET

where: 

  • p𝐾𝑎 = −log𝐾𝑎 (with Ka being the ionization constant of the weak acid),
  • [Salt] represents the concentration of the conjugate base (salt of the acid).
  • [Acid] represents the concentration of the weak acid.
  • For example, in the buffer CH3COOH/CH3COONa (acetate ion) is the conjugate base of acetic acid, CH3COOH.

For a Basic Buffer Mixture (Henderson-Hasselbalch Equation)

For a basic buffer solution containing a weak base and its conjugate acid:

1. pOH form:Acids, Bases & Salts | Chemistry Class 11 - NEET

2. Converting to pH:

  • Since pH + pOH = 14
  • pH=14−pOH
  • Using the relationship p𝐾𝑏 =14 −p𝐾𝑎 , we can write:Acids, Bases & Salts | Chemistry Class 11 - NEET
  • where: p𝐾𝑏 = −log𝐾𝑏 (with Kbeing the ionization constant of the weak base),
  • [Salt] represents the concentration of the conjugate acid (salt of the base), and
  • [Base] represents the concentration of the weak base.

For example, in the buffer NH4OH/NH4Cl, NH4+ (ammonium ion) is the conjugate acid of the base NH3NH, and K
\text{K}_a
Ka is the ionization constant of the reaction:Acids, Bases & Salts | Chemistry Class 11 - NEET

So for this basic buffer:Acids, Bases & Salts | Chemistry Class 11 - NEET

or equivalently,Acids, Bases & Salts | Chemistry Class 11 - NEET

Importance of Buffer Solution

1. In Biological processes

The pH of our blood is maintained constant inspite of various acid and base producing reactions going on in our body. The buffer action is due to the presence of carbonic acid , bicarbonate ion and carbon dioxide in the blood.

2. In Industrial processes

The use of buffers is an important part of many industrial processes, e.g.,

in electroplating, in the manufacture of leather, dyes, photographic materials.

3. In Analytical chemistry

  •  In the removal of acid radicals such as phosphate, oxalate and borate which interfere in the precipitation of radicals of group 3.
  • In complexometric titration
  • To calibrate the pH metres

4.In Bacteriological research, culture media are generally buffered to maintain pH required for the growth of the bacteria being studied.

Salts

Salts are the products of a neutralization reaction between an acid and a base.

Types of Salts

1. Normal salts These are obtained by complete neutralisation of an acid with a base, e.g., NaCI, K2SO4, etc.

2. Acidic salts These are formed by incomplete neutralisation of polybasic acids. e.g., NaHCO3, Na2SO4 etc.

3. Basic salts These are formed by incomplete neutralization of polyacidic base, e.g., Mg(OH)Cl, Bi(OH)2Cl, etc.

4. Double salts These are formed by the combination of two simple salts and exist only in solid state, e.g., Mohr salt or ferrous ammonium sulphate (FeSO4.(NH4)2SO4.6H2O], alum, etc.

5. Complex salts These are formed by the combination of simple salts or molecular compounds. These are stable in solid state well as in solutions.

6. Mixed salts These salts furnish more than one cation or more than one anion when dissolved in water, e.g., (CaNa₂(CO₃)₂) , NaKSO4, etc.

Conjugate Acid-Base Pair(CABP) 

In an acid-base reaction, a conjugate acid-base pair consists of two species that differ by a single proton \text{H}^(H+). When an acid donates a proton, it forms its conjugate base, while when a base accepts a proton, it forms its conjugate acid. This concept is central to the Bronsted-Lowry theory of acids and bases.

General Reaction

  • AcidH
    \text{H}^+
    H+ + Conjugate Base
  • Base + H+Conjugate Acid

Example

For the reaction:Acids, Bases & Salts | Chemistry Class 11 - NEET

HCl
  • HCl (acid) donates a proton to form  ClCl−, its conjugate base.
  • NH\text{NH}_3NH(base) accepts a proton to form NH4+, its conjugate acid.

Key Point

A conjugate acid-base pair differs only by one proton. For example:

  • HSO\text{HSO}_4^-HSO4 is the conjugate base of H2SO4 because they differ by one H\text{H}^+H+.
  • SOSO42−, however, is not the conjugate base of H2SObecause they differ by two protons.

Conjugate acid-base pairs are essential for understanding acid-base reactions, buffer systems, and equilibrium in solutions.

Relative strength of Acids/Bases

In acid-base chemistry, any species and its conjugate species have opposite strengths. A strong acid will have a weak conjugate base, and a strong base will have a weak conjugate acid. This relationship helps determine the direction and extent of acid-base reactions.

Example: Strength Order of Acids

The following is the strength order of some common acids, from strongest to weakest:

HClOAcids, Bases & Salts | Chemistry Class 11 - NEET

Strength Order of Conjugate Bases

The conjugate bases of these acids have opposite strength orders, from weakest to strongest:

Acids, Bases & Salts | Chemistry Class 11 - NEET

Ionic Equilibrium

In a solution, an ionic equilibrium exists between unionized electrolyte molecules and the ions resulting from their ionization. This dynamic balance is essential in determining the pH, conductivity, and reactivity of electrolyte solutions.

Types of keq

1. Self ionization of water

The self-ionization of water is the process where water molecules dissociate into hydronium and hydroxide ions.Acids, Bases & Salts | Chemistry Class 11 - NEETThe equilibrium constant (K_wKw) for water is:Acids, Bases & Salts | Chemistry Class 11 - NEET

2. Acid dissociation constant

The acid dissociation constant applies to the ionization of a weak acid in water.

For an acid HA dissociating as follows:Acids, Bases & Salts | Chemistry Class 11 - NEET

The equilibrium constant (Ka) is:Acids, Bases & Salts | Chemistry Class 11 - NEET

3. Base dissociation constant 

The base dissociation constant applies to the ionization of a weak base in water.

For a base B reacting with water:Acids, Bases & Salts | Chemistry Class 11 - NEET

The equilibrium constant (Kb) is:Acids, Bases & Salts | Chemistry Class 11 - NEET

4. Salt hydrolysis

Salt hydrolysis occurs when a salt reacts with water, typically producing acidic or basic solutions depending on the nature of the salt.

For example, the hydrolysis of NH4+ (from NH4Cl) in water:Acids, Bases & Salts | Chemistry Class 11 - NEET

The equilibrium constant (K_hKh) is: Acids, Bases & Salts | Chemistry Class 11 - NEET

5. Sparingly soluble salt

For a sparingly soluble salt, the solubility product applies when the salt partially dissolves in water.

For a salt like AB, which dissociates as: Acids, Bases & Salts | Chemistry Class 11 - NEET

The solubility product constant (K_{sp}Ksp) is:Acids, Bases & Salts | Chemistry Class 11 - NEET

Factors affecting the degree of ionisation and Bronsted and Lowry Concept of Acids and Bases

Factors affecting the degree of ionization

1. Temperature - With the rise in temperature, the degree of dissociation of an electrolyte in solution increased. Thus, Degree of dissociation α Temperature

2. Dilution : On the increasing of dilution, the degree of dissociation increases. But at infinite dilution, their is no effect on the degree of dissociation.

3. Concentration of the solution : 

Degree of dissociation ∝ 1/conc. of the sol.  ∝ 1/amount of solute in given vol. ∝ Amount of solvent

4. Nature of Solvent : Higher the dielectric constant of a solvent, more is its dissociation power or ionising power. Thus ,

Degree of ionisation or dissociation of an electrolyte ∝  dielectric constant of solvent. 

Dielectric constant : The dielectric constant of solvent is a measure of its tendency to weaken the forces of attraction between oppositely charged ions of the given electrolyte or the force of attraction applied by solvent molecules on solute molecule is defined as Dielectric constant of solvent.

Note : Water is the most powerful ionizing solvent as its dielectric constant is highest.

5. Presence of Common Ion : In the presence of a strong electrolyte having common ion, the degree of dissociation of an electrolyte decreases. eg. Ionisation of CH3COOH is suppressed in the presence of HCl due to common H  ions.

6. Nature of Electrolyte : At constant temperature, electrolytes ionize to a different extent in their solutions of same concentration.

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