According to Arrhenius concept of acid and bases,
1. Acid
2. Bases
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.
Some common examples of Lewis acids that can accept electron pairs include:
Limitations of Lewis Concept
Basic strength directly depends on the availability of lone pair for H.
Compare the basic strength of the following NH3, CH3NH2, (CH3)2NH, (CH3)3N
1. steric effect
2. Inductive effect
3. solvation 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 H3O+ ions. 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- 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.
pH is defined as the negative logarithm of hydrogen ion concentration.
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
Acid and base ionization constants
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: NH4OH/NH4Cl.
Buffer System in Blood: The buffer system present in blood is H2CO3+NaHCO3.
The Henderson-Hasselbalch equation is used to calculate the pH of a buffer solution.
For an Acidic Buffer:
For a Basic Buffer:
and pH = 14 - pOH
Where:
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).
For an acidic buffer solution containing a weak acid and its conjugate base:
where:
For a basic buffer solution containing a weak base and its conjugate acid:
1. pOH form:
2. Converting to pH:
For example, in the buffer NH4OH/NH4Cl, NH4+ (ammonium ion) is the conjugate acid of the base NH3, and Ka is the ionization constant of the reaction:
So for this basic buffer:
or equivalently,
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
4.In Bacteriological research, culture media are generally buffered to maintain pH required for the growth of the bacteria being studied.
Salts are the products of a neutralization reaction between an acid and a base.
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.
In an acid-base reaction, a conjugate acid-base pair consists of two species that differ by a single proton (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.
For the reaction:
A conjugate acid-base pair differs only by one proton. For example:
Conjugate acid-base pairs are essential for understanding acid-base reactions, buffer systems, and equilibrium in solutions.
The following is the strength order of some common acids, from strongest to weakest:
The conjugate bases of these acids have opposite strength orders, from weakest to strongest:
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.
1. Self ionization of water
The self-ionization of water is the process where water molecules dissociate into hydronium and hydroxide ions.The equilibrium constant (Kw) for water is:
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:
The equilibrium constant (Ka) is:
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:
The equilibrium constant (Kb) is:
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:
The equilibrium constant (Kh) is:
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:
The solubility product constant (Ksp) is:
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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