4.8.A.1 Brønsted-Lowry Acid-Base: 

1. Definition of Acids and Bases:

i. Brønsted-Lowry Definition of Acids and Bases:

The Brønsted-Lowry definition of acids and bases is a broader definition that compares with the earlier theories. This theory states:

Acid: A substance that transmits a proton (H⁺ ) to another substance.
Base: A substance that accepts a proton (H⁺ ) from another substance.

This theory deals with the transfer of protons between substances in aqueous or non-aqueous solutions. It is not limited to only those substances that have the ability to get partially ionized in water.

Example: In the reaction between hydrochloric acid (HCl) and water (H₂O):      HCl→H++Cl−

HCl gives a proton to water, so HCl is the acid, and water is a base that accepts the proton.

H2O + H⁺ → H3O⁺
Water (H2O) accepts a proton, forming the hydronium ion (H3O⁺).

ii. Comparison with Other Acid-Base Theories:

a. Arrhenius Definition:

The Arrhenius theory is one of the earliest theories and focuses primarily on water-based reactions:

Acid: A substance which increases the concentration of H⁺ ions in aqueous solution.
Base: A substance which increases the concentration of OH⁻ ions in aqueous solution.

Example: For HCl in water:      HCl→H⁺ + Cl⁻
HCl is a Lewis acid because the H⁺ions are available in its solution in water on dissociation. Likewise, OH⁻ions are available because of dissociation of NaOH in the water making NaOH as a Lewis base.

Limitation: Only aqueous solution is covered with the help of the Arrhenius definition but does not serve to explain acid and base reactions for non-aqueous solvents or in non-aqueous solution without the use of H⁺and OH⁻ions involved.

iii. Lewis Definition:

The Lewis theory is still more general and goes beyond proton transfer. The theory bases on electron pairs:

Acid: substance that accepts a pair of electrons (electron pair acceptor).
Base: substance that donates a pair of electrons (electron pair donor).

This definition covers a broader spectrum of reactions that involve some containing no protons at all. The reaction is composed of some where the metal ions or other molecules are acids or bases.

Example: For example, in the reaction between boron trifluoride (BF₃) and ammonia (NH₃):    BF3​+NH3​→NH3​BF3​
BF₃ is a Lewis acid since it accepts an electron pair from ammonia, NH₃, a Lewis base.

Limitation: The Lewis theory is incredibly broad and relevant to most reactions, but at times it simply is not so easy to use when attempting to describe simple proton transfer as if it were occurring through the Brønsted-Lowry or Arrhenius definitions.

2. Conjugate Acid-Base Pairs:

The concept of conjugate acid-base pairs should be understood based on how acids and bases chemically interact. Based on the Brønsted-Lowry theory, an acid-base pair consists of two species which differ in only a single proton H⁺ ionic charge:

A species that forms as a base following the acceptance of one proton(H⁺)
It is a species following an acid’s donation of just one proton(H⁺).

In any acid-base reaction, the acid and the base react to give their respective conjugate base and conjugate acid. The conjugate acid and conjugate base are related by the transfer of a proton.

i. General Example:

a. Acid-Base Reaction:       HA⇌H++A−
HA is the acid, and A⁻ is its conjugate base.
H⁺ is the proton being transferred.

b. Base-Action:
A base (B) accepts a proton (H⁺), creating its conjugate acid (BH⁺):     B+H+⇌BH+
The base is B, and the conjugate acid is BH⁺.

ii. Key Takeaways:
a. Conjugate acid: The species that results when a base accepts a proton.
b. Conjugate base: The species that results when an acid donates a proton.
c. The conjugate acid and conjugate base are connected by the transfer of a single proton (H⁺).
d. In the reaction, on the left side of the equation, the acid is bonded to its conjugate base and on the right-hand side of the equation, the base is bonded to its conjugate acid.

3. Strength of Acids and Bases:

The Brønsted-Lowry theory describes strength in terms of their tendency to easily give away or take in protons (H⁺) in that medium. It is based on how totally acids and bases will dissociate, or break apart, in an aqueous solution.

Strong acids and bases fully dissolve, or totally dissociate in aqueous solutions.
Weak acids and bases only partially dissociate, or break apart in solution.

Strong vs Weak Acids and Bases

i. Strong Acids:
A strong acid is an acid that completely dissociates (ionizes) in water; that is, it donates protons (H⁺) very easily. For example,
Hydrochloric acid (HCl)

• Hydrochloric acid (HCl):
  $$\text{HCl}\to {\text{H}}^{+}+{\text{Cl}}^{-}$$
• Sulfuric acid (H₂SO₄):
  $${\text{H}}_2{\text{SO}}_4\to {\text{H}}^{+}+{\text{HSO}}_4^{-}$$
• Nitric acid (HNO₃):
  $${\text{HNO}}_3\to {\text{H}}^{+}+{\text{NO}}_3^{-}$$

These acids fully donate protons, thus, in their aqueous solutions, the concentration of hydrogen ions is very high making the solution strongly acidic.
ii. Weak Acids:
A weak acid is an acid which only partially dissociates in water
So, a weak acid is one which partially donates its protons to water less as compared to the strong acids
Examples of weak acids:

• Acetic acid (CH₃COOH):
  $$\text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+$$
• Formic acid (HCOOH):
  $$\text{HCOOH} \rightleftharpoons \text{HCOO}^- + \text{H}^+$$

These acids are not fully dissociative, thus the concentration of hydrogen ions, H⁺, is somewhat low, so the solution is less acidic.

iii. Strong Bases:
A strong base is one that completely dissociates in water, giving hydroxide ions (OH⁻).
Some examples of strong bases are:
Sodium hydroxide (NaOH):
a strong base, the base fully dissociates and the concentration of hydroxide ions (OH⁻) is high so the solution becomes very basic or alkaline.

iv. Weak Bases:
A weak base accepts protons only partially or yields a relatively small concentration of hydroxide ions (OH⁻).
Some examples of weak bases are:
Ammonia (NH₃):    NH3​+H2​O⇌NH4+​+OH−
Pyridine (C₅H₅N):  C5​H5​N+H2​O⇌C5​H5​NH++OH−

Because these bases do not accept a proton fully, they yield fewer hydroxide ions in the solution and are less basic.
v. Factors Affecting the Strength of Acids and Bases:

Factors that affect the strength of acids and bases include the following:

a. Bond Strength and Polarity (for Acids)
If the bond of hydrogen and the atom it is bonded with happens to be a weak bond, it will donate a proton more easily, therefore stronger acids.
Even if polarity plays a very significant role; acids which have a very polar bond are very much likely to dissociate.
For example;
HCl is a strong acid because the bond between H and Cl is relatively weak. Chlorine’s electronegativity enhances the polarity of the substance. This makes proton donation easier.
CH₃COOH; Acetic acid is weaker as the bond between H and oxygen is relatively strong, not so polar and, hence, less favorable for proton donation.

b. Electronegativity of the Atom Bonded to Hydrogen (for Acids)
The electronegativity of the atom that the hydrogen is bonded to determines how easily it can donate a proton. The more electronegative, the stronger the acid.
For example:
Hydrofluoric acid is a weak acid, even though fluorine is highly electronegative due to the fact that H is strongly bonded to F.

c. Size of the Atom (for Acids)
The larger the atom to which hydrogen is bonded, the smaller the bond and the stronger the acid. This occurs because larger atoms will better stabilize the negative charge that forms when they shed a proton.
Example:
Hydroiodic acid (HI) is a stronger acid than HCl because iodine is larger than chlorine, making the bond weaker and easier to break.
d. Ionic Character (for Bases)
In the case of bases, their ability to accept protons-or better said donate a pair of electrons-is largely an issue of their ionic character and its propensity to stabilize those ions which then form
Strong bases; NaOH tends to readily dissolve and liberates a huge quantity of OH⁻
Ammonia as well as the ammonia salts tend to be less as the proton tends to have relatively weak interactions with the nitrogen atoms.
e. Solvent Effects
Solvent: Acid and base strengths can also be solvent dependent. That is, acids and bases that are strong in water don’t necessarily play the game as well in other solvents.
Aprotic solvents, lacking H⁺ ions, may modify the acid or base ionizations so that the relative acid or base strength changes, typically rendering them seem weaker than it would in water.

f.Temperature
The temperature will also dictate the degree of acid and base dissociation. For instance, weak acids and bases prefer a higher temperature when they dissociate.

g. Acid or Base Concentration
Strength relies on the strength of the concentration of the acid/base in solution. If its concentration is extremely low, then even the strongest acid will act as a weak acid since it will contain a very few number of protons.
The strong acids completely dissociate irrespective of concentration. Weak acids will only partly dissociate but the extent of the dissociation is independent of concentration.