acid
N H 4 ^ *
conjugate base
base
H_{2}O OH conjugate acid Box D X 5 H_{2}SO_{4} H_{2}O
S O 4 ^ 2-
The Correct Answer and Explanation is :
To answer this, let’s break it down step by step.
Identifying the Acid, Base, Conjugate Acid, and Conjugate Base:
- Acid: An acid is a substance that donates a proton (H⁺).
- Base: A base is a substance that accepts a proton.
- Conjugate Base: The conjugate base is what remains after an acid donates a proton.
- Conjugate Acid: The conjugate acid is what is formed when a base gains a proton.
Given Substances and Their Roles
- NH₄⁺ (Ammonium ion) → Acid (donates H⁺ to form NH₃)
- NH₃ (Ammonia) → Conjugate Base (formed after NH₄⁺ loses a proton)
- H₂O (Water) → Acts as either an acid or a base (amphoteric)
- OH⁻ (Hydroxide ion) → Conjugate Base of water when it donates H⁺
- H₂SO₄ (Sulfuric acid) → Acid (strong acid that donates protons)
- SO₄²⁻ (Sulfate ion) → Conjugate Base of H₂SO₄
Correct Matching
| Substance | Role |
|---|---|
| NH₄⁺ | Acid |
| NH₃ | Conjugate Base |
| H₂O | Amphoteric (can act as acid or base) |
| OH⁻ | Conjugate Base |
| H₂SO₄ | Acid |
| SO₄²⁻ | Conjugate Base |
Explanation (300 words)
Acids and bases are classified according to the Bronsted-Lowry theory, which defines acids as proton (H⁺) donors and bases as proton acceptors. Every acid has a corresponding conjugate base, which is formed after the acid donates a proton. Similarly, every base has a conjugate acid, which is formed when the base gains a proton.
For example, ammonium ion (NH₄⁺) acts as an acid because it donates an H⁺ to form ammonia (NH₃), which is its conjugate base. In water, H₂O can act as either an acid or a base. When H₂O donates a proton, it forms hydroxide (OH⁻), which is its conjugate base. On the other hand, when H₂O accepts a proton, it forms H₃O⁺, which is its conjugate acid.
Sulfuric acid (H₂SO₄) is a strong acid, meaning it fully dissociates in water to release protons. Its conjugate base is sulfate ion (SO₄²⁻), which is formed after H₂SO₄ loses protons.
Understanding these concepts is crucial in acid-base chemistry, as it helps explain reactions such as neutralization, buffer solutions, and pH regulation in biological systems. For example, the bicarbonate (HCO₃⁻) system in blood acts as a buffer, maintaining pH balance by shifting between carbonic acid (H₂CO₃) and bicarbonate (HCO₃⁻) depending on the body’s needs.
By recognizing acids, bases, and their conjugate pairs, we can predict reaction directions, buffer capacities, and chemical equilibrium in various environments.