Buffers are everywhere!

Buffer solutions and the buffering effect they produce are extremely important in many practical applications of chemistry. The reason for this is that many chemical processes are quite sensitive to the pH; the extent of the reaction, its rate, and even the nature of the products can be altered if the pH is allowed to change.  Such a change will tend to occur, for example, when the reaction in question, or an unrelated parallel reaction, consumes or releases hydrogen- or hydroxide ions.

Buffers in biochemistry and physiology

Many reactions that take place in living organisms fall into this category. Most biochemical processes are catalyzed by enzymes whose activities are highly dependent on the pH; if the local pH deviates too far from the optimum value, the enzyme may become permanently deactivated.

• For example, the oxidation of glucose to carbon dioxide and water — the fundamental energy-producing process of aerobic metabolism — releases H⁺ ions.  The multiple steps of this process are catalyzed by enzymes whose activity is strongly affected by the pH.  If the region of the cell where these reactions take place were not buffered, the very act of living would soon reduce the local pH below the range in which the enzymes are active, halting energy production and killing the cell.
• Blood is strongly buffered (mainly by bicarbonate) to maintain its pH at 7.4±0.3; pH values below 7.0 or above 7.8 cause death within minutes.
• The interiors of most cells are buffered near pH 7.0 (by phosphates and proteins)

Buffers in industry

Buffers are emloyed in a wide variety of laboratory procedures and industrial processes:

• solutions for the calibration of pH meters
• growth of bacteria in culture media
• fermentation processes, including winemaking and the brewing of beer
• controlling the colors of dyes used in coloring fabrics
• shampoos (maintaining pH at or below 7 to conteract the alkalinity of soaps that can cause irritation)
• other personal care products such as baby lotions (keeping the pH around 6 to inhibit the growth of bacteria)
• pH control in printing inks to assure their proper penetration into the paper

Buffers in the environment

• Natural waters (lakes and streams) able to support aquatic life are buffered by the sediments they are in contact with. This buffering is not always able to offset the effects of acid rain.
• Seawater is buffered mainly by bicarbonate and borates.  This allows the oceans to absorb about half of the CO₂ emitted to the atmosphere by human activities.
• Soils, in order to remain productive must maintain a near-neutral pH, not only for plants but also for microorganisms which fix- and recycle nitrogen.  Soils rendered highly acidic by acid rain can keep essential nutrients such as phosphates in forms such as insoluble Ca₃(PO₄)₂ that prevents their uptake by plants.

Carbon dioxide, however, is the major source of acidity in the unpolluted atmosphere.  As will be explained farther on, CO₂ makes up 0.032 vol-% of dry air, and dissolves in water to form carbonic acid:   CO₂(g) + H₂O(l) → H₂CO₃(aq).

Thus all rain is “acidic” in the sense that it contains dissolved CO₂ which will reduce its pH to 5.7.

The term acid rain is therefore taken to mean rain whose pH is controlled by pollutants which can lower the pH into the range of 3-4, resulting in severe damage to the environment.

Origin of acid rain

Acid rain originates from emissions of SO₂ and various oxides of nitrogen (known collectively as "NO_x") that are formed during combustion processes, especially those associated with the burning of coal. Incineration of wastes, industrial operations such as smelting, and forest fires are other combustion-related sources; these also release particulate material into the atmosphere which plays an important role in concentrating and distributing these acid-forming substances.

Origins of NO_x

Nitrogen oxides are formed naturally by soil bacteria acting on nitrate ions, an essential plant nutrient and itself a product of natural nitrogen fixation. Ammonia, a product of bacterial decomposition of organic matter, is eventually transformed into NO₃^– by bacteria into nitrogen oxides.

The quantity of fixed nitrogen produced industrially to support intensive agriculture now exceeds the amount formed naturally, and has become the major source of anthropogenic nitrogen oxides. 

• Atmospheric nitrogen can be thermally decomposed to its various oxides NO_x when subjected to high temperatures; one minor natural source is lightning. More significantly, these oxides are also formed when "clean" fuels such as natural gas (mainly methane, CH₄) is burnt in air.
• NO_x can also react with other atmospheric pollutants to produce photochemical smog.

Formation of acids

The gases noted above react with atmospheric oxygen, with each other, and with particulate materials to form sulfuric, nitric, and hydrochloric acids.

• Most of the H₂SO₄ comes from the photooxidation of SO₂ released from the burning of fossil fuels and from industrial operations such as smelting.
• Nitric acid similarly results from photooxidation of nitrogen oxides. Large quantities of these oxides are formed in high-temperature combustion processes in gasoline- and jet engines, and in the turbine engines of fossil-fueled electric power plants.
• Hydrochloric acid is formed naturally through the reaction of sea salt aerosols with atmospheric H₂SO₄ and HNO₃. The salt particles form when ocean waves break the surface and some of the water evaporates before the spray settles back to the surface. On a much smaller scale, fossil fuel combustion and waste incineration are believed to the principal human-caused sources of HCl.

Wet and dry deposition

There are two basic mechanisms by which acidic material is transported through the environment. The acid that dissolves in the water droplets that form clouds and precipitation, and is eventually incorporated into fog and snowflakes, undergoes wet deposition; this is ordinarily what is referred to as "acid rain". Because it is widely dispersed, often high in the atmosphere, it can travel very large distances from the original source.

About 40% of the acidic substances introduced into the atmosphere becomes adsorbed onto particulate matter, which can be soot, smoke, windblown dust, and salt particles formed naturally from sea spray. This spreads less widely (typically around 30 km from the source), and falls onto surfaces by "dry deposition". Once the acid-laden particles land on surfaces, ordinary rain, fog, and dew release the acidic components, often in considerably more concentrated form than occurs through dry deposition alone.

Incidence and effects of acid rain

Worldwide, acid rain is closely associated with regions where large quantities of coal are used for electric power generation. In North America, this area encompasses much of the northeastern part of the country.

↓ Click on an image to enlarge it ↓

Acid rain - North America [link]

Acid rain -Europe [link]

 

 

 

 

 

 

 

Atmospheric circulation can spread local pollution sources over sizeable regions.  For example, coal burning and other sources in the northeastern U.S. has increased the acidity of surface waters from the mid-Atlantic coast to eastern Canada.

Since the 1970s when the consequences of acid deposition became apparent, most advanced countries have taken steps to reduce atmospheric pollution. Coordinatge efforts to reduce trans-border pollution have been especially useful; a good example is this 2008 summary of progress under the US-Canada Air Quality Agreement.

 

Effects on soils

Because soils support the growth of plants and of the soil microorganisms that are essential agents in the recycling of dead plant materials, acid rain has a indirect but profound effect on soil health and plant growth.

• Most soils also contain clay and humic substances that bind and retain ions such as Ca²⁺, Mg²⁺ and K⁺ which are essential plant nutrients. Added H⁺ binds even more strongly to these substances, displacing the nutrient ions from the upper layers. The overall effect is to leach these ions to greater depths where they may be inaccessible to plants, or to wash them away into ground water.
• The SO₄^2– component of acid rain converts some nutrient cations into insoluble sulfates, reducing their availability to plants.
• Clays, which are complex aluminosilicates, gradually break down, releasing aluminum ions which normally form insoluble Al(OH)₃. Addition of H⁺ dissolves this hydroxide, raising the concentration of Al(H₂O)₆³⁺ to levels that can be toxic to plants.
• Many toxic heavy metals such as Pb, Cd and Cr are present in trace amounts in soils but are tied up in the form of insoluble salts. Acid rain can mobilize the ions of these metals, much as happens with aluminum.

Effects on plants

The effects on soils noted above affect plants most strongly. However, direct impingment of acidic rain and fog on leaves has other effects which can be especially serious when air pollutants such as SO₂ are present.

• Acid deposition on leaves and needles tends to weaken them by eroding away their protective waxy coatings.  This often results in the development of brown spots that interfere with photosynthesis.
• Forests in mountain regions are frequently bathed in clouds and fog which can be even more acidic than normal precipitation, exacerbating the above problem. In addition, the leaves and needles tend to accumulate the fog aerosol into larger droplets. These eventually fall to the forest floor, magnifying the effects on soils described in the previous section.
• Food crops tend to be less affected by acid rain where good farming practices such as the addition of fertilizers to replace depleted nutrients and the addition of limestone to raise the soil pH.

Effects on lakes and aquatic ecosystems

Lakes are subject not only to wet and dry acidic deposition, but also by the water they receive from streams and surface runoff. Any toxic elements released by the action of acid rain on soils and sediments can thus be conveyed to, and concentrated within lakes and the streams that empty them.

Lakes in poorly-buffered areas such as are found in alpine regions (western North America, Colorado and much of Switzerland) or on the Canadian Shield and in the Adirondacs and Appalachians) are highly sensitive to acid deposition.

 

 

Severely acidified lakes (such as the one depicted at the above right) can be so devoid of life, including algae, that the water appears to be perfectly clear and bright blue in color.

Aquatic organisms are generally adapted to "ordinary" pH conditions of 6-8, but vary greatly in their tolerance of low pH. As pH declines below 6, the diversity of aquatic animals, plants, and microorganisms diminishes. 

 

• Some plant species such as sphagnum mosses and certain filamentous algaes that thrive at very low pH can proliferate in acidified lakes, producing thick mats that seal off oxygen and thus inhibit the decay of litter on the lake floor.
• Acid deposited onto winter snow is released during spring melting and can cause rapid drops of pH in poorly-buffered lakes that receive waters from these sources. This "spring acid shock" can seriously affect viability of organisms such as fish, amphibians, and insects that lay their eggs in the water which hatch in the spring. The hatchlings are often unable to adapt to the rapid change, and end up deformed or killed.

• ↓ Click on image to enlarge

  

  Fish damaged by acid deposition [link]

  Even those aquatic species able to survive at in low-pH waters can be indirectly affected if their food supply is pH-limited.  For example, mayflies and some other insects, which, which are important food sources for frogs, cannot survive below pH 5.5.
• Fish are severely affected by the aluminum that is released from the action of acid on sediments; it causes a coating of mucus to form on the gills, impeding absorption of oxygen.  This has led to the extinction of some species from affected lakes.
• Low pH increases the solubility of calcium salts, making it more difficult for bone to form in developing embryos of fish and amphibians.
• Low pH can also make the eggs of aquatic organism thicker and more difficult for embryos to penetrate, thus delaying hatching. When the embryos continue to grow within the confined space of the egg, their spines can be deformed, interfering with their viability when they finally hatch.

 

Effects on buildings and statues

Acid deposition most strongly affects heritage buildings made of limestone and similar carbonate-containing stone materials.

The principal chemical process involves the reaction of sulfuric acid with calcium carbonate:  CaCO₃(s) + H₂SO₄ → CaSO₄(s) + CO₂.

The acid first erodes and breaks up the surface of the stone. As the hydrated calcium sulfate (gypsum) forms, it picks up iron and other components of the stone and forms an unsightly black coating. Some of this gradually blisters off, exposing yet more stone suface. The gypsum crystals can sometimes grow into the stone, further undermining the surface.

Very old structures such as the Taj Mahal, Notre Dame, the Colosseum and Westminster Abbey have all been affected.

Statues and monuments, including those made from marble, are also susceptible to erosion and damage from acid deposition.

Modern buildings are less affected, although acid rain can erode some painted surfaces.

Modern buildings are less affected, although acid rain can erode some painted surfaces.

 

Some related videos:

• Overview of acid rain chemistry (RThornley, 10½ min) ***
• How acid rain affects limestone buildings (Cwinski, 5 min)
• Acid rain lab demo (FlinnScientific, 10 min)

The carbonate system, consisting as it does of a soluble gas CO₂, soluble ions HCO₃^– and CO₃^2–, and sparingly soluble carbonate salts, spans all four realms of nature: the atmosphere, hydrosphere (mainly the oceans), the lithosphere, and the biosphere. And owing to the acid-base equilibria that govern the transformations between these carbonate species, carbon is readily transported between these geochemical reservoirs. This chapter presents an overview of carbon geochemistry.

Distribution of carbon on Earth

Carbon is the fourth most abundant element in the universe. Within the Earth's crust it ranks 15th, mostly in the form of carbonates in limestones and dolomites. Kerogens, which are fossilized plant-derived carbon mostly in the form of oil shale, constitute another major repository of terrestrial carbon.

 

Source	moles C × 1018	relative to atmosphere
carbonate rock	1530	28,500
fossil carbon	572	10,600
land - organic carbon	.065	1.22
ocean	3.2	61.8
atmospheric CO2	.0535	1.0

 

The geochemical carbon cycle

The carbonate system encompasses virtually all of the environmental compartments– the atmosphere, hydrosphere, biosphere, and major parts of the lithosphere. The complementary processes of photosynthesis and respiration drive a global cycle in which carbon passes slowly between the atmosphere and the lithosphere, and more rapidly between the atmosphere and the oceans. Thus the "carbon cycle" can be divided into "fast" and "slow" parts, operating roughly on annual and geological time scales.

 

 

↑ Image by David Bise, Pennsylvania State University; see here for a full discussion of this diagram.

Carbon dioxide in the atmosphere

CO₂ has probably always been present in the atmosphere in the relatively small absolute amounts now observed. Precambrian limestones, possibly formed by reactions with rock-forming silicates, e.g.

CaSiO₃ + CO₂ → CaCO₃ + SiO₂ (4-1)

 

Carbon in the oceans

Natural waters aquire carbon from sediments they are in contact with, and of course also from the atmosphere.  The pH is an important factor here; CO₂ and solid carbonates are more soluble at high pH, and pH also controls the distribution of carbon species, as is seen in the plot just above.

 

Carbon content and pH of typical natural waters

	rain	river/lake water	sea water
ppm of carbon	1.2	35	140
pH (unpolluted)	5.6	6.5 - 8.5	7.5 - 8.4

At the slightly alkaline pH of most bodies of water (of which the oceans constitute 97% of the earth's surface waters), bicarbonate is the principal dissolved carbon species. The quantity of organic carbon is fairly small.

Make sure you thoroughly understand the following essential concepts that have been presented above.

• Describe the major features of the global carbon cycle, and outline the role of acid-base processes in both the slow and fast parts of the cycle.
• Describe the principal acid-base reactions in aerobic respiration.
• Define acid rain, and explain its origins.  Outline its major effects on the environment.