5.4: CO2 and Ocean Acidification

In previous sections, you have learned that CO₂ concentrations are increasing in the atmosphere and that the ocean absorbs CO₂ from the atmosphere; in this section, we will learn what happens to the carbon dioxide that the ocean absorbs. Let’s start by exploring the relationship between pH and CO₂ using a week of data (July 2024) from the Pioneer MAB array. Answer the questions that follow the graph below.

Relationship between CO2 and pH at the Oceanographic Observatories Initiative Pioneer MAB array. Created by N. Gownaris. 

Exercise \(\PageIndex{1}\)

What kind of correlation do these data suggest between pH and aqueous CO₂?

Answer

A negative relationship - when CO2 is higher, pH is lower.

 

What is pH again?

pH is a measure of the acidity or basicity of a solution, specifically, it’s the negative logarithm (base 10) of the hydrogen ion concentration ([H+]). Lower pH values indicate higher concentrations of hydrogen ions (H+).  The “log” in pH means that each whole number on the pH scale represents a tenfold difference in hydrogen ion concentration.

At pH = 7 (pure water), hydrogen ion and hydroxide ion activity are equal. Solutions with pH values above 7 have lower concentrations of H+ ions and are considered basic or alkaline.  Solutions with pH values below 7 have higher concentrations of H+ ions and are considered acidic. The term acidification refers to a decrease in pH.

The pH of the ocean is slightly basic, with an average pH of about 8.1.  Even a small decrease in ocean pH, a phenomenon known as ocean acidification, can have significant impacts on marine life and ecosystems. Lowering the pH reduces the availability of calcium carbonate, a key building block for shells and skeletons, potentially leading to shell dissolution and other physiological problems in shell-building species. 

 

How the CO₂ Influences Ocean pH

How CO₂ Reacts with Water (H₂0)

When CO₂ gas dissolves in the ocean, it interacts with the water to produce a number of different compounds according to the reaction below:

CO₂ + H₂O ↔ H₂CO₃ ↔ H⁺ + HCO₃^– ↔ 2H⁺ + CO₃^2-

CO₂ reacts with water to produce carbonic acid (H₂CO₃), which then dissociates into bicarbonate (HCO₃^–) and hydrogen ions (H⁺). The bicarbonate ions can further dissociate into carbonate (CO₃^2-) and additional hydrogen ions (Figure \(\PageIndex{2}\)).

Figure \(\PageIndex{2}\) The fate of dissolved carbon dioxide in the oceans. Most of the carbon ends up in the form of bicarbonate (PW).

Most of the CO₂ dissolving or produced in the ocean is quickly converted to bicarbonate. Bicarbonate accounts for about 92% of the CO₂ dissolved in the ocean, and carbonate represents around 7%, so only about 1% remains as CO₂, and little gets absorbed back into the air. The rapid conversion of CO₂ into other forms prevents it from reaching equilibrium with the atmosphere, and in this way, water can hold 50-60 times as much CO₂ and its derivatives as the air.

 

CO₂ and pH

The equation above also illustrates carbon dioxide’s role as a buffer, regulating the pH of the ocean. Recall that pH reflects the acidity or basicity of a solution. The pH scale runs from 0-14, with 0 indicating a very strong acid, and 14 representing highly basic conditions. A solution with a pH of 7 is considered neutral, as is the case for pure water. The pH value is calculated as the negative logarithm of the hydrogen ion concentration according to the equation:

pH = -log₁₀[H⁺]

Therefore, a high concentration of H⁺ ions leads to a low pH and acidic condition, while a low H⁺ concentration indicates a high pH and basic conditions. It should also be noted that pH is described on a logarithmic scale, so every one point change on the pH scale actually represents an order of magnitude (10 x) change in solution strength. So a pH of 6 is 10 times more acidic than a pH of 7, and a pH of 5 is 100 times (10 x 10) more acidic than a pH of 7.

Carbon dioxide and the other carbon compounds listed above play an important role in buffering the pH of the ocean. Currently, the average pH for the global ocean is about 8.1, meaning seawater is slightly basic. Because most of the inorganic carbon dissolved in the ocean exists in the form of bicarbonate, bicarbonate can respond to disturbances in pH by releasing or incorporating hydrogen ions into the various carbon compounds. If pH rises (low [H⁺]), bicarbonate may dissociate into carbonate, and release more H⁺ ions, thus lowering pH. Conversely, if pH gets too low (high [H⁺]), bicarbonate and carbonate may incorporate some of those H⁺ ions and produce bicarbonate, carbonic acid, or CO₂ to remove H⁺ ions and raise the pH. By shuttling H⁺ ions back and forth between the various compounds in this equation, the pH of the ocean is regulated and conditions remain favorable for life.

 

CO₂ and Ocean Acidification

In recent years there has been rising concern about the phenomenon of ocean acidification. As described in the processes above, the addition of CO₂ to seawater lowers the pH of the water. As anthropogenic sources of atmospheric CO₂ have increased since the Industrial Revolution, the oceans have been absorbing an increasing amount of CO₂, and researchers have documented a decline in ocean pH from about 8.2 to 8.1 in the last century. This may not appear to be much of a change, but remember that since pH is on a logarithmic scale, this decline represents a 30% increase in acidity. It should be noted that even at a pH of 8.1 the ocean is not actually acidic; the term “acidification” refers to the fact that the pH is becoming lower, i.e. the water is moving towards more acidic conditions.

Figure \(\PageIndex{3}\) presents data from observation stations in and around the Hawaiian Islands. As atmospheric levels of CO₂ have increased, the CO₂ content of the ocean water has also increased, leading to a reduction in seawater pH. Some models suggest that at the current rate of CO₂ addition to the atmosphere, by 2100 ocean pH may be further reduced to around 7.8, which would represent more than a 120% increase in ocean acidity since the Industrial Revolution.

Figure \(\PageIndex{3}\) Changes in atmospheric CO₂ (red), seawater CO₂ (green) and pH (blue) in the Hawaiian Islands (NOAA PMEL).

 

When carbon dioxide enters into  the ocean, it influences shell-forming organisms in two ways.

 

The first is through a change in pH. As mentioned above, carbon dioxide (CO₂) interacts with water (H₂O) to form carbonic acid (H₂CO₃). However, carbonic acid rapidly dissociates in water to form a bicarbonate ion (HCO₃⁻), releasing a hydrogen ion (H+) in the process. This increase in free hydrogen ions leads to a reduction in pH, leading to more acidic waters than can dissolve calcium carbonate (CaCO₃) shells (Figure \(\PageIndex{4}\)). You can see the decline in pH (blue line) that mirrors the increase in atmospheric (red line) and oceanic (green line) CO2 in Figure \(\PageIndex{3}\). Note a special feature of this graph: it has a secondary y-axis. Scientists sometimes use a secondary y-axis when the parameters they want to show relationships between exist on very different scales. 

 

Exercise \(\PageIndex{2}\)

What range of values are shown on the axis for CO2? What range of values are shown on the axis for pH?

Answer

The CO2 (primary, or left) y-axis ranges from 275 to 425 while the pH (secondary, or right) y-axis ranges from 8.03 to 8.33

 

 

Figure \(\PageIndex{5}\) The results of an experiment placing the calcium carbonate shells of pterapods in seawater with a pH of 7.8, the projected ocean pH for the year 2100 under current rates of acidification. The top row shows the shells before the experiment, and the bottom row shows the dissolution of the shells after 45 days of exposure (NOAA).

 

The second way that CO₂ influences shell-forming organisms is through the ocean’s carbonate buffering system. The ocean carbonate buffering system is a natural mechanism that helps regulate the ocean’s pH, preventing it from becoming too acidic or too basic. It works through a series of chemical reactions involving carbon dioxide (CO₂), water (H₂O), carbonic acid (H₂CO₃), bicarbonate (HCO₃⁻), and carbonate (CO₃²⁻).

To maintain equilibrium, the additional hydrogen ions produced from the interaction between CO₂ and H₂O bind with carbonate ions, forming bicarbonate and increasing the pH. As a result, a decline in pH leads to a reduction in the availability of free carbonate ions, making it harder for organisms to form calcium carbonate shells (Figure \(\PageIndex{6}\)). In summary, some organisms can’t make shells because there are fewer carbonate ions and organisms struggle to extract enough material to build strong shells. Some organisms may still build shells, but at a greater energy expense, which can impact survival and reproduction. In seawater with lower pH existing shells can dissolve, especially in young or weaker organisms.

 

Figure \(\PageIndex{6}\) Proportions of carbon compounds in the ocean at various pH levels. As the ocean pH declines, the proportion of carbonate ions also declines, reducing rates of shell formation (NOAA).

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Review your Knowledge of the Ocean Carbonate System