by Dr. Sarah D. Oktay
Managing Director UMass Boston Nantucket Field Station
In the last issue of Yesterday’s island/Today’s Nantucket we learned about pH and the importance of various acid concentrations to the development of wine profiles and soil. This week, we will talk about some normal pH ranges for different types of water. Hopefully you recall that the pH scale was invented by the Danish chemist Søren Sørensen in 1909. A Ph.D. from the University of Copenhagen, Sørensen was the director of the chemical department of the Carlsberg Laboratory, which was supported by the beer company of the same name, brewing being one of the oldest chemical industries. Chemists were aware of things being more acidic or basic, but they did not have a way of expressing it effectively and quantitatively each time until Sørensen invented the pH scale. Once again beer & brewing help contribute to our greater knowledge!
As a very quick reminder, pH indicates a sample’s acidity, but is actually a measurement of the potential activity of hydrogen ions (H+), usually in an aqueous sample. pH measurements typically fall on a scale from 0 to 14, with 7.0 considered neutral. Solutions with a pH below 7.0 are considered acids. Solutions with a pH above 7.0 are considered bases. For some odd reason, folks seem to think there are a lot of neutral pH 7 things in the world when it is actually relatively rare to find something completely neutral.
Lots of things affect the pH of water in natural systems. All organisms are subject to the amount of acidity of a stream, pond, a vernal pool or coastal habitat and function best within a given range. One factor which affects the pH is the amount of plant growth and organic material within a body of water. When this material decomposes, carbon dioxide is released. The carbon dioxide combines with water to form carbonic acid. Although this is a weak acid, large amounts of it will lower the pH. An important factor which affects pH is the amount of acid precipitation that falls in the watershed. Acid rain is caused by nitrogen oxides (NOx) and sulfur dioxide (SO2) in the air combining with water vapor. These pollutants are primarily from automobile and coal-fired power plant emissions.
Many people don’t realize that rain water is actually relatively acidic due to its trip through our lower atmosphere, where it has time to come in contact with carbon dioxide and other gasses that contribute to an average world wide pH for rainfall of about 5. The National Atmospheric Deposition Program has developed maps showing pH patterns, and they very clearly show the more acidic rainfall along the east coast.
Excessively high and low pHs can be detrimental for the use of water. High pH causes a bitter taste, water pipes and water-using appliances become encrusted with deposits, and it depresses the effectiveness of the disinfection of chlorine for those municipalities that still use chlorine for disinfection, thereby causing the need for additional chlorine when pH is high. Low-pH water will corrode or dissolve metals and other substances. The pH of water determines the solubility (amount that can be dissolved in the water) and biological availability (amount that can be utilized by aquatic life) of chemical constituents such as nutrients (e.g., phosphorus, nitrogen, and carbon) and heavy metals like lead, cadmium, and copper. For example, in addition to determining how much and what form of phosphorus is most abundant in the water, pH also determines whether aquatic life can use it. Heavy metals tend to be more toxic at lower pH because they are more soluble and more bioavailable.
Streams and ponds and other fresh water bodies can be more acidic or basic than rainfall, depending on the amount of leaves dissolved in them and the contribution from humans in addition to the amount of organic matter. Photosynthesis uses up hydrogen molecules, which causes the concentration of hydrogen ions to decrease and therefore the pH to increase. Respiration and decomposition processes lower pH. For this reason, pH is higher during daylight hours and during the growing season, when photosynthesis is at its peak. Plants and animals as they decomposed tend to make water more acidic. The tannins come from dissolved leaves that you’ll see when you look at a darkly stained pond (very orangey) like the one here at the Field Station. These leaves and organic matter from cells break down into a whole suite of important acids such humic, tannic and fulvic acids which means many of our ponds are pretty acidic. Our groundwater is also relatively acidic (4.2-6.0 pH) unless it has come in contact with a soil amendment or leaking septic system which might make it more basic (relatively unusual).
The oceans are strongly buffered by a complex interaction between calcium carbonate (what seashells are made of) and many other salts and dissolved solids. For millions of years they have maintained a pH of about 8.2. The oceans absorbed and buffered the natural contributions that can change pH. That careful balance has survived over time because of a near equilibrium among the acids emitted by volcanoes and the bases liberated by the weathering of rock. One consequence of this changing chemistry is a reduction in the saturation state of seawater with respect to calcite and aragonite: two common types of calcium carbonate formed by marine organisms. Evidence is now mounting that such changes in seawater chemistry can have consequences for a wide variety of marine organisms. Many calcifiers—including commercially important species such as oysters, mussels, and sea urchins as well as corals—exhibit reduced calcification rates in response to elevated CO2 levels. In contrast, growth rates of some seagrasses and nitrogen-fixing cyanobacteria appear to be enhanced by increased CO2.
Ocean acidification is a relatively newly discovered threat for ocean life that has been emerging and accelerating over the past few years. Ocean acidification is the name given to the ongoing decrease in the pH of the Earth’s oceans, caused by the uptake of anthropogenic carbon dioxide (CO2) from the atmosphere. About a quarter of the carbon dioxide in the atmosphere goes into the oceans, where it forms carbonic acid (HCO3-)
The oceans are becoming more acidic (less basic) as the result of excess carbon dioxide entering the atmosphere. As the amount of carbon has risen in the atmosphere there has been a corresponding rise of carbon going into the ocean. Between 1751 and 1994 surface ocean pH is estimated to have decreased from approximately 8.25 to 8.14, representing an increase of almost 30% in “acidity” (H+ ion concentration) in the world’s oceans. This ongoing acidification of the oceans poses a threat to the food chains connected with the oceans.
Last week at the Lighthouse school I did a very simple experiment with the third graders that explains this concept extremely easily and clearly. I brought in one of my favorite items, a Sodaclub/Sodastream soda maker which injects pressurized carbon dioxide gas into water held in reinforced liter bottles for instant seltzer water at a fraction of the cost and a big saving to the environment (no plastic bottles to discard). You can add all kinds of flavorings and make your own carbonated beverage. I call it “pop” in one of my many obvious Okie mannerisms (to see what carbonated beverages are called in various states check out http://popvssoda.com/). Anyway, back to class, this device is easily portable and kids love it and it helps me explain ocean acidification in less than a minute! We take tap water, measure the pH using fine scale pH paper; then carbonate the water and measure it again. Normally the pH drops by at least 1 unit (10 times the acidity) in moments. We are able to experiment with warm or cool water to look at other properties of gasses in liquids.
When scientists first started recording these more acidic conditions in the oceans (less basic), they thought this meant all shelled organisms would have a tougher time making their shells; but it appears (according to work done by University of North Carolina researcher Justin Ries) that there are winners and losers. Researchers grew 18 different species of economically and ecologically important marine calcifiers (creatures that make their shells out of calcium carbonate) at various levels of CO2 predicted to occur over the next several centuries. When CO2 combines with water, it produces carbonic acid, raising the overall amount of carbon in seawater but reducing the amount of the carbonate ion used by organisms in their calcification.
Seven species (crabs, lobsters, shrimp, red and green calcifying algae, limpets and temperate urchins) showed a positive response, meaning they calcified at a higher rate and increased in mass under elevated CO2. Ten types of organisms (including oysters, scallops, temperate corals and tube worms) had reduced calcification under elevated CO2, with several (hard and soft clams, conchs, periwinkles, whelks and tropical urchins) seeing their shells dissolve. One species (mussels) showed no response. This is, of course, a worry for us on Nantucket because we depend on many of these fisheries. There may be relatively simple ways to offset this pH change locally and scientists on the Cape and Islands are investigating how effective efforts like returning shells back to the harbors to buffer the water might be in balancing out the added carbon dioxide input. Of course the coral reefs that have been suffering from bleaching do not need the added stress of a dissolving infrastructure.
I’d prefer to end on an upbeat note and tie back into the start of our story, so what is your guess on the typical pH of beer? Why of course Søren Sørensen would know that it is 4.5. For the blog of the week this week, I went to The New York Times Green Blog (http://green.blogs.nytimes.com) for an article on ocean acidification that ties in much of the recent information concerning large events in the past relative to ocean pH, check out http://www.sciencemag.org/content/335/6072/1058.abstract for more.