pH will differ across thermoclines, chemoclines and in water with different salinity levels, though how salinity affects pH will also depend on the chemical composition of the water. 

Environmental processes that influence pH in water do so because of the chemical reaction between water and carbon dioxide. When dissolved in water, CO₂ reacts with H₂O to form carbonic acid (H₂CO₃). Carbonic acid can then break down into bicarbonate (HCO₃-), releasing a hydrogen ion (H⁺). Increase in the hydrogen ion concentration of water lowers the pH, making it more acidic. The resulting bicarbonate ion may gain back a hydrogen ion, and the exchange can continue to shift pH back and forth depending on the concentrations of each element.

 
CO₂ is also responsible for the pH of rain. As precipitation falls through the atmosphere, the same reaction between CO₂ and H₂O takes place, lowering the pH of rain to a slightly acidic 5.6.  

The same chemical reactions described above take place when rain, rivers or other sources of surface water pass over or through soil, rock and other natural materials that contain carbonates. The presence of carbonate rock, like limestone, or other carbonate materials can change the pH of water it comes in contact with, changing the pH of the water bodies that eventually receive that water. Other substances can also alter pH when runoff or precipitation encounters them. Volcanic ash, decomposing plant matter and airborne particulates are just a few examples.  

pH is an important parameter to monitor when studying harmful algal blooms (HABs). Algae consumes CO₂ during photosynthesis, reducing carbonic acid formation and making water more basic. The usual diurnal cycle of plant activity may create small fluctuations in pH due to photosynthesis but not enough to adversely affect aquatic life. With an overabundance of algae, however, pH may quickly and dramatically increase, interfering with a water body’s ability to buffer these changes and placing stress on aquatic organisms. 

When discussing influences on pH, it’s important to touch on the role of alkalinity as well. The USGS defines alkalinity as “a measure of the ability of the water body to neutralize acids and bases and thus maintain a fairly stable pH level.” This is also known as its “buffering capacity.” Highly alkaline solutions will be more resistant to the addition of acidic substances; their pH will change, but more slowly than a solution with a lower alkalinity. For example, if acidic runoff entered a high-alkaline lake, the acidity of the lake might not change much; a lower-alkaline lake would become more acidic in response to the same runoff.  

A water body becomes alkaline when runoff or precipitation collects certain chemicals, such as calcium carbonate (CaCO₃), from rocks or soil before entering the water. When acid is introduced to a solution with a high concentration of carbonates or bicarbonates, compounds in these carbonates and bicarbonates will neutralize the acid, and the pH of the solution will remain stable. If enough acid is introduced, the balance of hydrogen and hydroxide ions will change and pH will decrease. But, for solutions with higher alkalinity, this process occurs more slowly due to the neutralizing effect, or buffer, of bicarbonates. 

Water bodies with higher alkalinity (and therefore, higher buffering capacity) are generally able to better support aquatic life, as organisms can’t adapt well to sudden, dramatic changes in pH. A poorly buffered soil or water body will experience more dramatic fluctuations in pH when acidic or basic substances enter or pass through it.  
 

Human activity can release chemicals into water, air and soil that influence pH. If not properly monitored and treated, industrial runoff, effluent discharge and agricultural runoff can all include chemicals that change pH when they enter surface water bodies.  

Mining operations also produce chemicals that can enter water through runoff or seep into soil and groundwater. Additionally, mining operations can create acid mine drainage (AMD), which occurs when water reacts with rocks containing sulfur-bearing minerals, such as pyrite (FeS₂). This reaction produces sulfuric acid, which can leach heavy metals from the rocks. When mining operations excavate landscapes, they create opportunities for exposed rock to come in contact with water—whether rainfall, snowmelt or flooding. This increases the likelihood that sulfur-bearing minerals will react with water and create AMD. 

Acid rain is one of the most well-known influences on the pH of water bodies. Acid rain can form after natural events such as volcanic eruptions, but fossil fuel emissions are a far more common cause.  

As discussed under “Natural Influences,” rain already has a slightly acidic pH of 5.6 due to the reaction between H₂O and CO₂ as rain falls through the atmosphere. When industrial operations release sulfur dioxide (SO₂) and nitrogen oxides (NO_X) into the atmosphere, these chemicals react with water to lower the already-acidic pH of rain. It then becomes acid rain, defined as having a pH value between 4.2 and 4.4.  

In certain environments, like deserts, sulfuric and nitric acids may settle on surfaces as dry deposition, descending in gas or particulate form in the absence of rain. This doesn’t eliminate the threat of pollution; when rain or runoff washes over areas where dry deposition has settled, the water will collect these acids and carry them into surface water bodies or groundwater. 

As CO₂ in the atmosphere increases due to fossil fuel emissions, more and more of it dissolves into the ocean. The resulting creation of carbonic acid and its dissociation into bicarbonate and hydrogen ions is responsible for ocean acidification across the globe. According to the National Oceanic and Atmospheric Administration (NOAA), the pH of seawater has dropped by 0.1 pH units since the onset of the industrial revolution. As pH is a logarithmic scale, even a 0.1 decrease in pH represents a big difference; NOAA estimates this shift indicates a 30 percent increase in ocean acidity. You can read more on ocean acidification from NOAA here. 

Most organisms are sensitive to changes in pH and have limited ranges they can live in, so monitoring pH is an important indicator of the overall health of a water body. The solubility of most chemicals is pH dependent and determines the availability of those chemicals to living organisms.  

pH can also indicate the presence of pollutants. Tracking changes in pH can help identify chemicals entering water bodies through agricultural runoff, treated wastewater, mining runoff or other industrial effluent. Low pH could also point to acid rain; acidic surface water may warrant monitoring for air pollution, or testing the soil to determine if rain carried chemicals into nearby water bodies after passing over contaminated land. 

pH is also relevant to the study and detection of harmful algal blooms (HABs). Fresh water can be slightly acidic (with a pH as low as 6), depending on the concentration of dissolved carbon dioxide (CO₂). The carbon dioxide combines with water to form a small amount of carbonic acid (H₂CO₃) and this process lowers the pH. Conversely, consumption of excess CO₂ can raise the pH of water. As HABs grow, algae consumes more and more CO₂. Therefore, a rapid rise in pH can indicate the formation of HABs. 

pH is also important to consider in relation to ammonia and ammonium–for the assessment of environmental health and also for commercial aquaculture. pH values below 5 in natural waters are generally too acidic for most aquatic life. Freshwater fish species survive best in the pH range of 6 to 9. pH affects the ammonia/ammonium equilibrium in water. At a pH of 6.5, almost all ammonia is in the form of ammonium. However, as the pH becomes more basic, ammonium becomes harmful ammonia. Fish can tolerate a moderate amount of ammonium, but even a small amount of ammonia is detrimental to their health. The lethal dose of ammonia for trout, for example, is only 0.2 mg/L. It’s important that outdoor aquaculture operations–whether established in ponds, lakes, harbors or coastal nearshore environments—monitor pH closely to ensure small changes in ammonia don’t wipe out their harvest.  

pH is also important to monitor in recreational water bodies for public health. While humans can tolerate a wider range of pH levels compared to many other organisms, pH outside of range 4-11 can cause skin and eye irritation for humans.  

On a global scale, pH in coastal waters is a sign of increasing ocean acidification. The pH of oceans worldwide has remained relatively stable for millennia. But since the industrial revolution, seawater has become more acidic as it absorbs carbon dioxide from the atmosphere. NOAA estimates a current average ocean pH of 8.1–a 30 percent increase in acidity in the last 200 years. Scientists are forecasting a drop to 7.8 in the next 100 years if current rates of carbon emissions continue. While many organisms can tolerate a pH as low as 6, acidic water slows and prevents the formation of CaCO₃. Ocean acidity will present problems for calcium carbonate-forming organisms like corals and mollusks–and potentially other organisms as well if acidification continues at the current rate. That kind of change would lead to unprecedented impacts on aquatic life. It’s important to continue monitoring ocean acidity to understand how this global shift is affecting our seas. 

The matrix groundwater flows through can affect its pH, so monitoring pH in groundwater may be an indicator of the chemical composition of formations surrounding a well and, therefore, what substances are flowing into the aquifer. 

Just like surface water bodies, groundwater is subject to the influences of acid rain and pollutants from industrial activities, because rain and runoff that flows through surface water can also seep into soil and groundwater. pH levels in groundwater can point toward which substances are influencing pH and whether those chemicals are naturally occurring or from human activity.  

Additionally, monitoring pH during drilling or well development is a helpful measure to assess contamination. A change from initial pH levels can be a sign that a substance is leaching into the well. If groundwater becomes basic during well development, that could indicate the presence of cement. 

Because pH influences the solubility and availability of chemicals in water, monitoring pH during water and wastewater treatment can inform decision making around process adjustments.  

pH plays an important role in contaminant removal. For example, chlorine is more effective in the disinfection process when water is in the neutral pH range of 6.5 to 7.5. Many pathogen-removal processes are pH dependent, and many EPA regulations regarding standards for disinfection include requirements for pH. As for influent, monitoring pH in source water helps drinking water facilities determine the most effective approach to treat water coming into their plant.  

Ensuring specific pH levels in finished water is essential for effective corrosion control. Acidic drinking water can corrode plumbing and appliances. While the EPA defines a pH between 6.5 to 8.5 as healthy drinking water, a pH between 7.2 and 7.6 for finished water ensures optimal corrosion control. Outside of this range, pH can contribute to lead leaching from brass fittings and dissolving in distribution infrastructure. It can also cause formation of lead-carbonate scale.   

pH also impacts the effectiveness of filtration. Natural organic matter, volatile organic compounds (VOCs) and synthetic organic compounds (SOCs) are all influenced by pH. These reactions can make filters more or less effective depending on the type of filtration material used. 

Paying close attention to pH is critical for effective chemical dosing. As mentioned above, pH influences the solubility and availability of chemicals for specific reactions and processes. Changing acidity may inhibit treatment processes, so it’s important to measure pH at chemical dosing points to ensure pH is conducive to the desired chemical process. 

Especially important to keep in mind, however, is the logarithmic nature of the pH scale. Because each number represents a change in pH to a factor of ten, even small changes in pH can have dramatic effects on the efficiency of chemical dosing. It is therefore essential to clean, calibrate and maintain pH sensors according to manufacturer instructions to avoid costly overdosing or underdosing in response to potentially inaccurate measurements.  

At times, acidic or alkaline substances may be added to influence pH during the treatment process. For example, certain ion compounds, such as ferric acid, may be required for phosphate removal, but increase the acidity of water–a reaction that must be corrected later in the process. 

pH also has an impact on biochemical processes. It can influence microorganism activity, an important factor for phases of water and wastewater treatment involving bacteria, such as nitrification. The optimal pH for nitrification is between 8 and 9. Free ammonia (NH₃) and free nitrous acid (HNO₂) slow the nitrification process by interfering with nitrifying bacteria. Nitrification diminishes the level of bicarbonates and increases carbonic acid; this causes a drop in pH. If pH drops during the process, ammonia oxidation also declines. Aeration helps mitigate this change by removing carbon dioxide from the tank. Any treatment process that involves bacterial activity could be influenced by these chemical reactions in a similar way.