Marine Science (9693) | Ocean acidification

Welcome to Chapter 9.3: Ocean Acidification

Hi future Marine Scientist! This chapter dives into one of the most serious consequences of human activity on the ocean: Ocean Acidification (OA). It might sound like a chemistry lesson, but don't worry! We will break down the essential chemical reactions and explain exactly why this phenomenon is so damaging to iconic marine life like coral reefs and shellfish.

Understanding OA is crucial because it connects global climate change (Topic 9.2) directly to marine chemistry (Topic 1) and biological impacts (Topic 5). Let's get started!

Section 1: The Ocean as a Carbon Sink

Ocean acidification starts with one simple fact: the ocean absorbs carbon dioxide ($\text{CO}_2$) from the atmosphere.

What is a Carbon Sink? (Syllabus 9.3.1)

The term carbon sink refers to any natural reservoir that absorbs and stores carbon-containing chemical compounds for an indefinite period.

• The ocean is the largest active carbon sink on Earth. It absorbs roughly 25% to 30% of the anthropogenic (human-caused) $\text{CO}_2$ released into the atmosphere, primarily from burning fossil fuels.
• Benefit: This uptake helps to limit the increase in atmospheric carbon dioxide concentrations, slowing the rate of global warming.
• Cost: When $\text{CO}_2$ is absorbed by the ocean, it changes the water's chemistry, leading to acidification.

Key Takeaway: The ocean acts as Earth's chemical sponge, soaking up excess $\text{CO}_2$, which is good for the atmosphere, but bad for the seawater chemistry itself.

Section 2: The Chemistry of Acidification

When $\text{CO}_2$ dissolves into seawater ($\text{H}_2\text{O}$), it triggers a series of chemical reactions that result in a decrease in pH (an increase in acidity) and a decrease in the availability of essential building blocks for marine life.

The Three Key Reactions (Syllabus 9.3.2)

Here is the step-by-step process of how absorbed $\text{CO}_2$ lowers ocean pH:

Step 1: Carbon Dioxide Forms Carbonic Acid

When atmospheric $\text{CO}_2$ dissolves into water, it forms Carbonic Acid ($\text{H}_2\text{CO}_3$).

$$ \text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3 $$

Carbonic acid is a weak acid, meaning it doesn't stay whole for long.

Step 2: Carbonic Acid Releases Hydrogen Ions ($\text{H}^+$)

The carbonic acid quickly dissociates (breaks apart) into a Hydrogen Ion ($\text{H}^+$) and a Hydrogen Carbonate Ion (Bicarbonate) ($\text{HCO}_3^-$).

$$ \text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^- $$

The key impact here is the production of free Hydrogen Ions ($\mathbf{H}^+$). A higher concentration of $\text{H}^+$ ions is the definition of increased acidity, which corresponds to a decrease in pH.

Did you know? The ocean's pH scale is naturally alkaline (around 8.1). "Acidification" means the pH is moving towards the acidic end (7.0), not that it has become an acid.

Step 3: Hydrogen Ions Steal Carbonate Ions ($\text{CO}_3^{2-}$)

The newly formed $\text{H}^+$ ions (the chemical signature of acidification) react aggressively with the Carbonate Ions ($\mathbf{CO}_3^{2-}$) already present in the seawater. Carbonate ions are essential for marine life to build shells.

When the $\text{H}^+$ ions meet $\text{CO}_3^{2-}$ ions, they quickly combine to form more Hydrogen Carbonate (Bicarbonate).

$$ \text{H}^+ + \text{CO}_3^{2-} \rightleftharpoons \text{HCO}_3^- $$

The net result of this reaction is twofold:

1. pH Drop: The overall concentration of free $\text{H}^+$ still increases (acidifying the water).
2. Carbonate Ion Depletion: The available supply of $\mathbf{CO}_3^{2-}$ ions is drastically reduced because they are being consumed to form bicarbonate.

Analogy for Step 3: Imagine Carbonate Ions ($\text{CO}_3^{2-}$) are LEGO blocks needed to build shells. When we add too much $\text{CO}_2$, it produces Hydrogen Ions ($\text{H}^+$) which act like bullies, grabbing those LEGO blocks and turning them into useless Bicarbonate ($\text{HCO}_3^-$). This leaves fewer blocks for corals and molluscs to build their homes.

Section 3: Impact on Hard Corals and Shelled Organisms (Calcification)

Ocean acidification particularly harms organisms that rely on calcification—the process of forming shells or skeletons out of Calcium Carbonate ($\mathbf{CaCO}_3$). These organisms are called calcifiers (Syllabus 9.3.3).

The Need for Calcium Carbonate

Calcifiers build their structures using calcium ions ($\text{Ca}^{2+}$) and carbonate ions ($\text{CO}_3^{2-}$):

$$ \text{Ca}^{2+} + \text{CO}_3^{2-} \rightleftharpoons \text{CaCO}_3\ \ (\text{Calcium\ Carbonate}) $$

Impact on Hard Corals (Syllabus 9.3.3)

Hard corals (such as reef-building corals) produce massive limestone skeletons that form the reef structure.

• Reduced Growth Rate: With less $\text{CO}_3^{2-}$ available (due to OA), corals must expend significantly more energy to extract the ions they need. This slows down their growth rate and makes their skeletons weaker and more brittle.
• Erosion Risk: In highly acidic conditions, the increased concentration of $\text{H}^+$ ions can actually cause existing $\text{CaCO}_3$ skeletons to dissolve, a process known as corrosion. Reefs may start to break down faster than corals can rebuild them.

Impact on Shelled Organisms (Syllabus 9.3.3)

Molluscs and other shelled organisms also suffer, especially during their early life stages when shells are thin and rapidly growing.

• Difficulty Forming Shells: Organisms like oysters, clams, sea snails, and important plankton species (like pteropods, often called 'sea butterflies') struggle to form their protective shells.
• Larval Mortality: Juvenile and larval stages are often the most vulnerable. If they cannot form shells efficiently, they are unlikely to survive to adulthood, leading to population crashes.
• Practical Link (9.3.4): A core practical activity involves placing empty mollusc shells in solutions of varying pH. When the pH is low (acidic), the shell mass decreases significantly due to the dissolution of the calcium carbonate structure. This simulates the impact of severe ocean acidification.

Key Takeaway: Ocean acidification undermines the foundational building process (calcification) for key marine organisms like corals and shellfish, threatening entire food webs.

Section 4: The Ocean's Natural Defence: Buffering

While we say the ocean is "acidifying," the change in pH is relatively slow compared to freshwater systems reacting to the same amount of $\text{CO}_2$. This is due to the ocean's remarkable buffering capacity (Syllabus 9.3.1).

What is Buffering?

A buffer is a solution that resists changes in pH when small amounts of acid ($\text{H}^+$) or base are added.

• Seawater contains a large natural reservoir of Bicarbonate Ions ($\mathbf{HCO}_3^-$) and Carbonate Ions ($\mathbf{CO}_3^{2-}$). These are known as the ocean carbon chemistry system.
• When extra $\text{H}^+$ ions are introduced (from dissolved $\text{CO}_2$), the natural carbonate ions ($\text{CO}_3^{2-}$) in the water absorb these $\text{H}^+$ ions, slowing the drop in pH (as seen in Step 3 above).
• This process is why the ocean helps limit the increase in atmospheric $\text{CO}_2$ concentrations—the chemical equilibrium shifts to absorb the excess gas.

The ocean's buffer system is powerful, but because human emissions have been so high and sustained, we are gradually overwhelming this natural system, causing a measurable, long-term shift towards lower pH across all surface waters.

Review of Impact Sequence

If you are asked to explain the full cycle of ocean acidification, remember the sequence of events:

1. Atmospheric $\text{CO}_2$ increases (human activity).
2. Ocean absorbs excess $\text{CO}_2$ (acting as a carbon sink).
3. $\text{CO}_2$ reacts with $\text{H}_2\text{O}$ to form $\text{H}_2\text{CO}_3$.
4. $\text{H}_2\text{CO}_3$ dissociates, releasing $\mathbf{H}^+$ ions (decreasing pH).
5. The excess $\mathbf{H}^+$ ions combine with available $\mathbf{CO}_3^{2-}$ ions.
6. Availability of $\mathbf{CO}_3^{2-}$ decreases, hindering $\text{CaCO}_3$ formation (calcification) in corals and shelled organisms.

Key Takeaway: The ocean’s natural buffer system slows acidification, but it does so by sacrificing the very ions ($\text{CO}_3^{2-}$) that calcifying organisms need to survive.

Summary: Core Concepts for Examination

To ensure you ace the OA questions, focus on these critical syllabus links:

• Cause: Increased atmospheric $\text{CO}_2$ dissolving in seawater.
• Chemical Result 1 (Acidity): Formation of $\text{H}^+$ ions ($\text{CO}_2 \rightarrow \text{H}_2\text{CO}_3 \rightarrow \text{H}^+$). This lowers the pH.
• Chemical Result 2 (Building Blocks): Consumption of $\mathbf{CO}_3^{2-}$ ions by $\text{H}^+$ ions ($\text{H}^+ + \text{CO}_3^{2-} \rightarrow \text{HCO}_3^-$).
• Biological Effect: Hard corals and shelled organisms struggle or fail to form protective $\text{CaCO}_3$ skeletons/shells, leading to reduced growth and dissolution.

Keep practicing those chemical equations—they are the backbone of this topic!