Among the chemical species :(i) H 2 CO 3(ii) $HCO_3^-$iii) $CO_3^{2-}$(iv) Cl-the pH of natural waters is controlled mainly by

Question

Prepp Team · Accepted Answer

Understanding pH Control in Natural Waters

Natural waters, such as lakes, rivers, and oceans, maintain a relatively stable pH despite the addition of acidic or basic substances. This stability is primarily due to the presence of buffering systems. A buffer system is a mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resists changes in pH.

The Carbonate Buffer System

The most important buffering system in natural waters is the carbonate system. This system involves the equilibrium between dissolved carbon dioxide ($CO_2$), carbonic acid ($H_2CO_3$), bicarbonate ions ($HCO_3^-$), and carbonate ions ($CO_3^{2-}$).

The relevant species given in the question are:

(i) $H_2CO_3$ (Carbonic acid)
(ii) $HCO_3^-$ (Bicarbonate ion)
(iii) $CO_3^{2-}$ (Carbonate ion)
(iv) $Cl^-$ (Chloride ion)

Let's examine how the first three species function as a buffer system:

Carbon dioxide from the atmosphere dissolves in water to form carbonic acid:

$CO_2(aq) + H_2O(l) \rightleftharpoons H_2CO_3(aq)$

Carbonic acid then dissociates in two steps:

$H_2CO_3(aq) \rightleftharpoons HCO_3^-(aq) + H^+(aq)$

$HCO_3^-(aq) \rightleftharpoons CO_3^{2-}(aq) + H^+(aq)$

These reactions show the interconnectedness of $H_2CO_3$, $HCO_3^-$, and $CO_3^{2-}$. This system can neutralize both added acids and bases:

If an acid (adding $H^+$) is added, $HCO_3^-$ and $CO_3^{2-}$ can react with the added $H^+$ ions, shifting the equilibrium to the left and consuming the $H^+$, thus minimizing the decrease in pH.
If a base (adding $OH^-$) is added, $H^+$ ions from the dissociation of $H_2CO_3$ and $HCO_3^-$ can react with the $OH^-$ ions (forming water), or $H_2CO_3$ can react with $OH^-$ to form $HCO_3^-$. These reactions replenish the consumed $H^+$ or consume the added $OH^-$, minimizing the increase in pH.

Role of Chloride Ion

The chloride ion ($Cl^-$) is the conjugate base of a strong acid, hydrochloric acid (HCl). The conjugate base of a strong acid is a very weak base and does not significantly react with water or $H^+$ ions to affect pH. Therefore, $Cl^-$ acts as a spectator ion in most natural water systems and does not play a significant role in buffering or controlling pH.

Conclusion

Based on the roles of these chemical species in water chemistry, the pH of natural waters is mainly controlled by the carbonate system, which includes carbonic acid ($H_2CO_3$), bicarbonate ions ($HCO_3^-$), and carbonate ions ($CO_3^{2-}$). The chloride ion ($Cl^-$) does not contribute to pH buffering.

Therefore, the species mainly controlling the pH of natural waters are (i), (ii), and (iii).

Species	Chemical Formula	Role in Natural Waters
Carbonic Acid	$H_2CO_3$	Weak acid, part of buffer system
Bicarbonate Ion	$HCO_3^-$	Amphoteric species, key buffer component
Carbonate Ion	$CO_3^{2-}$	Weak base, part of buffer system
Chloride Ion	$Cl^-$	Spectator ion, does not buffer pH

Revision Table: pH Control in Natural Waters

Concept	Description
pH Buffering	Resistance to change in pH upon addition of acid or base.
Carbonate System	Main buffer system in natural waters involving $H_2CO_3$, $HCO_3^-$, $CO_3^{2-}$, and dissolved $CO_2$.
Species Involved	$H_2CO_3$, $HCO_3^-$, $CO_3^{2-}$ are the active buffering species.
Spectator Ions	Ions that do not significantly participate in chemical reactions controlling pH, like $Cl^-$.

Additional Information: Natural Water Buffering Systems

While the carbonate system is the most significant buffer in natural waters, other systems can also contribute, especially in specific environments. These include:

Phosphate system: Important in nutrient-rich waters, although usually at lower concentrations than the carbonate system.
Organic buffers: Weak organic acids and bases present in dissolved organic matter can also provide some buffering capacity.
Silicate system: Plays a long-term role in geological buffering through mineral weathering.

The effectiveness of a buffer system is highest when the pH of the water is close to the pKa values of the weak acid components of the system. For the carbonate system, the relevant pKa values are around 6.35 (for $H_2CO_3/HCO_3^-$) and 10.33 (for $HCO_3^-/CO_3^{2-}$). Natural waters typically have a pH range where the bicarbonate ion ($HCO_3^-$) is the dominant species, making the $H_2CO_3/HCO_3^-$ and $HCO_3^-/CO_3^{2-}$ equilibria crucial for pH control.

Prepp Team · Answer

Only (i) and (ii)

Prepp Team · Answer

Only (ii) and (iv)

Prepp Team · Answer

(i), (ii), (iii) and (iv)