Gemini and chatGPT Q & A

It's known that warmer water holds less dissolved gas. With ocean temperatures rising and more CO2 being released into the atmosphere, why is the ocean becoming more acidic, not less?

Gemini Answer

That’s a great question that highlights a common misconception! Here’s why a warming ocean doesn’t lead to less acidity:

CO2 Solubility: While it’s true that warmer water generally holds less dissolved gas, the increase in atmospheric CO₂ is far greater than the decrease in solubility due to warming. This means the ocean is still absorbing more CO₂ overall.
Chemical Reactions: When CO₂ dissolves in seawater, it doesn’t just stay as CO₂. It reacts with water to form carbonic acid (H₂CO₃). This acid then dissociates, releasing hydrogen ions (H⁺) and bicarbonate ions (HCO₃⁻). The increase in H⁺ ions is what increases the acidity (lowers the pH) of the ocean.
Ocean as a Carbon Sink: The ocean is a massive carbon sink, absorbing about 30% of the CO₂ released into the atmosphere. As atmospheric CO₂ levels rise due to human activities, the ocean absorbs more, leading to further acidification.

Think of it like this:

Imagine a glass of water. If you add a few drops of lemon juice, it becomes slightly acidic. Now, warm that glass up. It might hold slightly less gas, but it’s still acidic because the lemon juice (the acid) is still there. In this case, the increasing CO₂ in the atmosphere is like adding more and more lemon juice to the ocean, even if some of it escapes due to warming.

In summary:

Ocean warming and ocean acidification are both consequences of increased atmospheric CO₂.
While warmer water holds less gas, the increase in atmospheric CO₂ outweighs this effect.
The chemical reactions of CO₂ in seawater lead to increased acidity, regardless of slight changes in solubility due to warming.

P. Doesn’t bicarbonate act as a base?

A. Bicarbonate (HCO₃⁻) is tricky because it can act as both an acid and a base. This is called being amphiprotic.

Here’s why:

It can accept a proton (H⁺): When it reacts with an acid, bicarbonate can accept a proton to become carbonic acid (H₂CO₃). This is the behavior of a base.
It can donate a proton (H⁺): When it reacts with a base, bicarbonate can donate a proton to become a carbonate ion (CO₃²⁻). This is the behavior of an acid.

So, which is it more likely to be?

In most biological systems, like our blood or the ocean, bicarbonate acts primarily as a base. This is because it helps to neutralize acids and maintain a stable pH.

Think of it like this:

Imagine a teeter-totter. Bicarbonate is sitting in the middle. If you add something acidic, it tips towards being a base to balance it out. If you add something basic, it tips the other way to act like an acid.

Ultimately, whether bicarbonate acts as an acid or a base depends on the specific environment and what other substances are present.

P. If the ocean is absorbing more CO2 and producing more bicarbonate ions, why is it still becoming more acidic? Shouldn’t the bicarbonate ions be counteracting the increased acidity?

A. You’re absolutely right to think that! Bicarbonate ions do play a crucial role in buffering the ocean’s pH and counteracting the effects of increased hydrogen ions. However, it’s not a simple one-to-one balance, and here’s why:

Equilibrium Shift: When CO2 dissolves in seawater, it sets off a chain of reactions that form carbonic acid (H₂CO₃), which then dissociates into bicarbonate (HCO₃⁻) and hydrogen ions (H⁺). This is a reversible reaction, meaning it can go in both directions. As more CO2 is added, the equilibrium shifts, producing more H⁺ ions and more bicarbonate ions.
Limited Capacity: While bicarbonate acts as a buffer, its capacity to neutralize acid is limited. As more and more CO2 is absorbed, the system becomes overwhelmed, and the pH continues to drop despite the presence of bicarbonate.
Carbonate Availability: Another critical factor is the availability of carbonate ions (CO₃²⁻). These ions are essential for marine organisms to build their shells and skeletons. However, the increase in H⁺ ions reacts with carbonate ions to form more bicarbonate. This reduces the availability of carbonate, making it harder for these organisms to thrive.

Think of it like this:

Imagine a seesaw with bicarbonate on one side and hydrogen ions on the other. As you add more hydrogen ions (from CO2), bicarbonate tries to balance it out. But if you keep adding more and more hydrogen ions, eventually, the bicarbonate side won’t be able to keep up, and the seesaw will tip towards the acidic side.

In summary:

Bicarbonate does act as a buffer to counteract ocean acidification.
However, its capacity is limited, and the increasing CO2 levels are overwhelming the system.
The increase in H⁺ ions also affects the availability of carbonate ions, crucial for marine organisms.

So, while bicarbonate is helping to mitigate the effects of ocean acidification, it’s not a complete solution, and the ongoing absorption of CO2 continues to pose a significant threat to marine ecosystems.

chatGPT Answer

Your question highlights an important interplay between ocean warming and carbon dioxide (CO₂) dynamics. At first glance, it seems logical to think that warmer oceans would release more CO₂ into the atmosphere, potentially reducing ocean acidity. However, the relationship is more nuanced due to the complex chemistry and physics involved.

Key Points to Consider:

CO₂ Solubility and Temperature:
- As water warms, its ability to dissolve gases, including CO₂, decreases. This is why warmer oceans might release some CO₂ back into the atmosphere. However, this effect doesn’t entirely offset other factors at play.
Ocean-Atmosphere CO₂ Exchange:
- While warmer water might release more CO₂, the oceans still act as a major carbon sink due to the sheer volume of CO₂ emitted from human activities. This excess CO₂ is absorbed regardless of temperature to some extent, driven by the difference in CO₂ concentrations between the atmosphere and the ocean.
Ocean Acidification Process:
- When CO₂ is absorbed by seawater, it reacts with water to form carbonic acid, which then dissociates into bicarbonate and hydrogen ions. The increase in hydrogen ions lowers the pH, making the water more acidic.
- Even if warming reduces CO₂ absorption efficiency slightly, the sheer volume of CO₂ entering the oceans from human activities is so large that the acidification process continues.
Feedback Loops:
- Warmer oceans might also disrupt biological and physical processes that regulate CO₂, such as the ability of phytoplankton to photosynthesize effectively or changes in ocean circulation patterns that bring carbon-rich deep waters to the surface.
Regional Variability:
- The interplay between ocean warming, CO₂ absorption, and pH is not uniform across the globe. Local factors such as ocean currents, upwelling, and biological activity influence whether a region of the ocean is a net source or sink of CO₂.

Conclusion:

While ocean warming slightly reduces CO₂ solubility, the massive influx of anthropogenic CO₂ into the atmosphere continues to drive ocean acidification globally. Thus, the warming ocean is not enough to counterbalance the overall increase in acidity caused by human activities. The result is a warming ocean that is simultaneously becoming more acidic, a dual stressor for marine ecosystems.

P. Doesn’t bicarbonate act as a base?

A. Yes, bicarbonate ions (HCO₃⁻) are considered a weak base in many chemical contexts, but their behavior is more nuanced because they can act as both a base and an acid. This property makes bicarbonate ions an important amphiprotic species in chemistry.

Why Bicarbonate is a Base:

As a Base: Bicarbonate can accept a proton (H⁺) to form carbonic acid (H₂CO₃), which then can decompose into water (H₂O) and carbon dioxide (CO₂).
HCO₃⁻ + H⁺→ H₂CO₃

This reaction demonstrates its basic properties when it interacts with an acid.

Why Bicarbonate is Amphiprotic:

As an Acid: Bicarbonate can also donate a proton (H⁺) to form carbonate (CO₃²⁻):
HCO₃⁻→ H⁺ + CO₃²⁻

This occurs in the presence of a strong base, highlighting its acidic properties.
Buffering Role in Seawater and Blood:
- In ocean chemistry, bicarbonate is crucial in buffering pH levels. It can react with hydrogen ions to reduce acidity or with hydroxide ions to mitigate alkalinity.
- In human blood, bicarbonate is part of the carbonic acid-bicarbonate buffer system that helps maintain pH homeostasis.

Bicarbonate in Ocean Chemistry:

In the ocean, bicarbonate forms as CO₂ reacts with water and dissociates:

CO₂ (gas) + H₂O (liquid) ⇋ H₂CO₃ ⇋ H⁺ + HCO₃⁻

Here, bicarbonate acts as a major reservoir of dissolved inorganic carbon, buffering the ocean’s pH against dramatic shifts.

In summary, bicarbonate is primarily a weak base but can act as an acid depending on the chemical environment, making it a versatile and essential component in many natural systems.

P. If the ocean is absorbing more CO2 and producing more bicarbonate ions, why is it still becoming more acidic? Shouldn’t the bicarbonate ions be counteracting the increased acidity?

A. Yes, bicarbonate ions (HCO₃⁻) play a crucial role in buffering hydrogen ions (H⁺) in the ocean. However, the buffering capacity of bicarbonate has limits, and the current rate of ocean acidification is pushing this system toward its threshold. Let’s break this down:

How Bicarbonate Ions Buffer Hydrogen Ions:

The Carbonate-Bicarbonate Buffer System: The equilibrium between carbon dioxide (CO₂), carbonic acid (H₂CO₃), bicarbonate (HCO₃⁻), and carbonate (CO₃²⁻) maintains ocean pH:

CO₂ (gas) + H₂O ⇋ H₂CO₃ ⇋ H⁺ + HCO₃⁻ ⇋ 2H⁺ + CO₃²⁻

- Bicarbonate as a buffer: When H⁺ ions are added (increased acidity), bicarbonate reacts with them to form carbonic acid:
  HCO₃⁻ + H⁺ → H₂CO₃
  
  Carbonic acid can then break down into water and CO₂, which can outgas to the atmosphere.
- Similarly, if the ocean becomes too basic (alkaline), bicarbonate can donate H⁺ to stabilize pH.
Dominance of Bicarbonate:
- Bicarbonate makes up the majority (~85%) of the dissolved inorganic carbon in seawater, acting as the primary buffer for small pH changes.
- This system is why the ocean’s pH has been relatively stable over geological timescales.

Why the Buffering System is Strained by Ocean Acidification:

Excess CO₂ Input: Human activities release vast amounts of CO₂, much of which dissolves in the ocean. As more CO₂ enters, it reacts with water to form more carbonic acid, leading to an increase in H⁺ ions:

CO₂ + H₂O → H₂CO₃ → H⁺ + HCO₃⁻

This excess overwhelms the buffering system. Even though bicarbonate neutralizes some of the H⁺ ions, the increased input skews the equilibrium toward greater acidity.
Shifting Carbonate Chemistry: As more H⁺ ions are produced, carbonate ions (CO₃²⁻) are converted into bicarbonate to balance the reaction:
H⁺ + CO₃²⁻ → HCO₃⁻

This reduces the availability of carbonate ions, which are critical for marine organisms like corals and shellfish that rely on calcium carbonate (CaCO₃) to build their skeletons and shells.
Limits of Buffering:
- While bicarbonate can counteract some H⁺ ions, there’s a tipping point where the system cannot keep up with the rapid CO₂ influx.
- Ocean pH has already dropped from about 8.2 to 8.1 since the Industrial Revolution—a small but significant change on the logarithmic pH scale.

Conclusion:

Bicarbonate ions do counterbalance hydrogen ions and act as an essential buffer in the ocean. However, the sheer volume of CO₂ being absorbed is overwhelming this natural buffering system. This results in:

A gradual decrease in pH (ocean acidification).
A decline in carbonate ion availability, jeopardizing marine life that depends on calcium carbonate structures.

In essence, the buffering capacity of bicarbonate ions helps slow down the rate of acidification but cannot completely offset the effects of the current CO₂ surge.