4.7.A.1 Acid-Base Reactions: Proton Transfer:

1. Definition of Acids and Bases:

i. Arrhenius Definition:
Acid: A substance that increases the concentration of hydrogen ions (H⁺) or protons when dissolved in water.
Base: A substance that increases the concentration of hydroxide ions (OH⁻) when dissolved in water.
Example:
Acid: HCl → H⁺ + Cl⁻
Base: NaOH → Na⁺ + OH⁻

ii.Bronsted-Lowry Definition:
Acid: A proton (H⁺) donor.
Base: A proton (H⁺) acceptor.
Example:
In the reaction

$\text{HCl} + \text{H₂O} \rightleftharpoons \text{Cl}⁻ + \text{H₃O}⁺$

, HCl contributes a proton (an acid), and H₂O accepts a proton (a base).

iii. Lewis Definition:
Acid: A species that can accept a pair of electrons to form a coordinate covalent bond.
Base: A substance that can donate a pair of electrons to form a coordinate covalent bond.
Example: In the reaction between boron trifluoride (BF₃) and ammonia (NH₃), BF₃ acts as the Lewis acid (electron pair acceptor), and NH₃ acts as the Lewis base (electron pair donor).

2. Proton Transfer Mechanism:

The proton transfer mechanism is the basis of many acid-base reactions, especially in aqueous solutions, where a proton, H⁺, which is simply a hydrogen atom without its electron, moves from one species to another. This mechanism can be understood by the roles of donors and acceptors.

i. Understanding H⁺ Ion Transfer:

The H⁺ ion is simply a bare nucleus—a single proton—since the hydrogen ionized loses its lone electron. Thus, it’s very reactive and can readily bond with other species. In aqueous solution, the H⁺ ion does not appear often in its bare form but is instead often present as the hydronium ion, H₃O⁺, bonded to a water molecule.

The transfer of H⁺ is accompanied by breaking up the bonds in which the proton is withdrawn from one molecule-the acid-and forms a bond where the proton is accepted by another molecule-the base. The reaction can often be written in an equilibrium form as follows:

$$\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-$$

In this, HA is a kind of acid, from whence the H⁺ ion has been removed, and A⁻ is a conjugate base.

ii. Role of Donors and Acceptors:

In the case of proton transfer, we class substances into two primary roles:

Proton Donor (Acid):
A proton donor is a substance that can release or donate a proton (H⁺) to another species.
The Bronsted-Lowry definition defines an acid as any substance that donates a proton in a reaction.
Example: In the dissociation of hydrochloric acid (HCl):

$$\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$$

Here, HCl is the proton donor.

Proton Acceptor (Base):
A proton acceptor is a species that can receive or accept a proton (H⁺) from another species.
As per the Bronsted-Lowry definition, a base is any species that accepts a proton in a reaction.
Example: Ammonia with water.

$$\text{NH}_3 + \text{H}_2\text{O} \rightarrow \text{NH}_4^+ + \text{OH}^-$$

iii. Water as a Proton Transfer Species:
Water is both an acid and a base. The very small fraction of water molecules ionizes in pure water to form hydronium (H₃O⁺) and hydroxide ions (OH⁻):

$$\text{H₂O}\rightleftharpoons {\text{H}}^{+}+{\text{OH}}^{-}$$

When H₂O acts as a proton donor, it forms OH⁻. Conversely, when it acts as

If H₂O acts as a proton donor, then it creates OH⁻, and if it acts as a proton acceptor, then it creates H₃O⁺.

In summary, proton transfer consists of the donation of a proton, H⁺, from acid to the base. Hence, it forms new bonds and shifts it around. The acid is the proton donor, and the base is the proton acceptor. The movement of the proton plays a crucial role in many chemical and biological processes, such as enzyme functions, acid-base control in the body, and acid-base reactions in aqueous solutions.

3. Conjugate Acid-Base Pairs:

A conjugate acid-base pair is two species that vary by one proton (H⁺).
Acid: loses a proton (H⁺).
Base: gains a proton (H⁺).
,When an acid releases a proton the species produced will be its conjugate base, and when a base accepts a proton, the species formed will be its conjugate acid.

Examples:
i. HCl (acid) → Cl⁻ (conjugate base)
ii. CH₃COOH (acid) → CH₃COO⁻ (conjugate base)
iii. NH₃ (base) → NH₄⁺ (conjugate acid)
iv. H₂O (acid) → OH⁻ (conjugate base)
v. H₂SO₄ (acid) → HSO₄⁻ (conjugate base)

Note: Strong acids have weak conjugate bases, and strong bases have weak conjugate acids.

4. Strength of Acids and Bases:

Strong Acids/Bases: Completely ionize in water.
Example: HCl (acid), NaOH (base).
Characteristics: High H⁺ or OH⁻ concentration, very low (acid) or high (base) pH.

Weak Acids/Bases: Partially ionize in water.
Example: CH₃COOH (acid), NH₃ (base).
Characteristics: Low H⁺ or OH⁻ concentration, moderate pH.

Ionization in Water:
The strong acids/bases fully ionize, whereas the weak acids/bases form an equilibrium state, almost completely ionizing.

5. pH and Buffers:

• pH and pOH

i. Relationship between [H⁺] and pH:
pH measures the acidity of a solution based on the concentration of hydrogen ions [H⁺].
pH is calculated as:

$$\text{pH} = -\log[H⁺]$$

High [H⁺] = low pH (acidic solution).
Low [H⁺] = high pH (basic solution).

ii. Calculating pH and pOH:
pOH is a measure of the basicity of a solution in terms of the hydroxide ion concentration [OH⁻].
pOH is defined as:   pOH = -log[OH⁻]
Relationship:

$$\text{pH} + \text{pOH} = 14$$

(In pure water at 25°C,

$[H⁺] = [OH⁻] = 1 \times 10^{-7}$

[H+]=[OH−]=1×10−7, so pH = 7 and pOH = 7).
Neutralization occurs when an acid reacts with a base to form water and a salt.

• Buffers and Buffering Capacity

Importance of Buffers to Maintain pH**
Buffer solutions resists a drastic change in their pH when acid or base solution is added little by little in small quantities
buffers are, therefore, importance in biological tissues such as that in bloods, which makes sure that life continues with good performance
How Buffers Function in Acids and bases:
Buffered solution is, therefore, weak acid-salt of this weak acid /weak base/ salt of a weak base.
When an acid is added, the buffer absorbs the H⁺ ions, and when a base is added, the buffer releases H⁺ ions to maintain pH.
Example: In blood, the carbonic acid (H₂CO₃) and bicarbonate (HCO₃⁻) buffer system helps maintain a pH of about 7.4.

Buffers effectively neutralize added acids or bases by shifting equilibrium to maintain constant pH.