Friday, March 2, 2012
In scientific terms, acid rain refers to any kind of precipitation, including mist, snow, fog, and of course, rain, that is more acidic than normal. Most rain is naturally a bit acidic, but acid rain contains an above average level of acid in it. Generally speaking, acid rain is caused by emissions of sulfur dioxide and nitrogen oxides that react with hydroxyl radicals and water vapor that exist in many industrial environments. When this combination exists, the acid rain may come down as either dry acid deposition or, when it is mixed with water, it is known as acid rain.
What is most acid rain composed of? Acid rain as it falls in the eastern part of North America and parts of Europe is composed mostly of sulfuric acid and nitric acid. How do these things make up acid rain? Acid rain generally occurs when the burning of fuels produces sulfur dioxide and nitrogen oxides. These different oxides get into our atmosphere because of both natural environmental activity as well as human activity. When these oxides reach the troposphere, they become oxidized by the hydroxyl radicals in the atmosphere that then break down the oxides into sulfuric and nitric acids. These acids will usually break down readily into water that is then brought down in the form of precipitation, or acid rain.
So is acid rain a by-product of global warming? It is not so simple. Many natural sources are also a part of acid rain. Many tons of sulfur is released into the earth's atmosphere each year from natural sources, including volcano eruptions, microbial processes, and sea sprays. Nitrogen oxides are also released into the earth's atmosphere in a natural manner, including from burning, lightning, the burning of biomass, and many microbial processes.
However, in a sense, acid rain is indeed a type of by-product of global warming because human activity often is responsible for some kinds of acid rain. It is estimated that human beings release up to 100 to 130 million tons of sulfur dioxide into the atmosphere. Human beings are also estimated to be responsible for roughly 60 to 70 million tons of the nitrogen oxides that are released into the earth's atmosphere each year. Most acid rain occurs in highly industrialized areas where these oxides are released into the earth's atmosphere on a regular basis. However, human activity has caused more oxides to be released into the earth's atmosphere in certain concentrated areas. Thus, human activity is definitely a strong factor in the occurrence of acid rain, especially in highly concentrated areas.
The effects of acid rain are becoming recognized as a growing problem, especially around highly industrial areas. Areas that have been highly industrialized for more than 100 years are considerably more susceptible to experiencing acid rain. However, all parts of the world are susceptible to some kind of acid rain. Acid rain is especially having an effect on many fragile ecosystems, including many of the earth's aquatic ecosystems. Acid rain can also have a devastating effect on forests.
The appearance of calcium carbonate is generally that of a white powder or stone. A characteristic quality of calcium carbonate is that it will fizz and release carbon dioxide upon contact with a strong acid, such as hydrochloric acid. After the carbon dioxide is released, the remainder is calcium oxide (CaO), commonly called quicklime.
When calcium carbonate comes into contact with water saturated with carbon dioxide, it forms a soluble compound, calcium bicarbonate. Underground, this often leads to the formation of caves. The reaction is as follows:
CaCO3 + CO2 + H2O → Ca(HCO3)2
Calcium carbonate becomes marble when highly compressed and heated deep underneath the Earth’s surface. In caves, when temporarily dissolved by the above chemical mechanism, calcium carbonate creates magnificent speleothems:” cave formations such as stalagmites, stalagmites, curtains, and dozens of others.
There are many mineral formations characteristic of calcium carbonate, but one of the most common forms is the scalenohedron, or “Dogtooth Spar,” for its resemblance to the canine tooth of a dog.
Calcium carbonate in the form of calcite has an interesting optical property: double refraction. This occurs when a ray of light enters the crystal and splits into distinct fast and slow beams. When an observer looks through the crystal, they see two images of everything behind it.
Calcium carbonate possesses other unusual properties, such as fluorescence and triboluminescence. Meaning, combined with a small amount of manganese and put under a UV light, calcium carbonate glows bright red. Under some conditions, the glow even persists when the UV light is removed. Triboluminescence, the property of demonstrating light when pieces of crystal are struck against each other, is more difficult to demonstrate, but it has been noted.
Photochemical smog is an atmospheric condition that produces severe eye irritation and poor visibility, to name just two of the effects. Three ingredients -- energy from a light source (ultraviolet), hydrocarbons, and nitrogen oxides -- are needed for photochemical smog to be formed. Two of those components are produced through the burning of fossil fuels, most notably automobiles. Photochemical smog is also sometimes known as "oxidizing smog", in that it has a high concentration of oxidizing agents. Ozone is a common oxidizing agent found in photochemical smog. Another type of smog, "reducing smog", has high concentrations of sulfur dioxide, which is a reducing agent. Its presence has historical significance in studies done in places like London, which used sulfur-containing coal as its main energy source.
A schematic is again useful in understanding the chemistry of smog production in the troposphere:
This graphic is relatively self-explanatory. It begins at the bottom with the production of NO and reactive hydrocarbons by fossil fuel burning (such as an automobile). On the left side, the NO reacts with tropospheric ozone or a hydrocarbon radical (RO2 ) to produce NO2 (a radical is a molecule fragment that has an unpaired electron). This absorbs solar energy (represented by the letters hv) to create NO (which propagates the system) and atomic oxygen. Atomic oxygen reacts to form tropospheric ozone, which feeds back into the NOx system (the "x" here refers to the number of oxygens, and serves as a general notational term for the nitrogen oxides). Atomic oxygen can also react with hydroxyl radicals, OH, and ozone to form the reactive hydrocarbon radicals utilized in the NOx system. These radicals also react to form other components of smog, such as PAN (peroxyacetyl nitrate) and aldehydes (RC=OH, where R is some hydrocarbon chain).
The graphic below shows the essential workings of the NOx system, with the interactions between NO and NO2 on the left and the production/washout of HNO3 on the right:
We can also look at the formation of photochemical smog from a kinetic perspective. This chart shows the nine key equations of smog production, and the rate constant that affects the speed (or rate) at which the reaction takes place:
In this graphic, we see the involvement of the NOx system and the production of ozone. Here is a narrative version of the graphic above:
Reaction 1: NO2, reacts with light energy, hv , to form NO and a singlet oxygen atom. The rate of this reaction depends on how much light energy there is -- a sunny day versus a cloudy day!
Reaction 2: singlet oxygen reacts with the oxygen molecule (what you breathe) in the presence of a catalyst "M" to form ozone, O3. The catalyst M remains unchanged (which is the definition of a catalyst!). The rate of this reaction depends on the temperature in the atmosphere.
Reaction 3: Ozone reacts with NO to produce more NO2 and O2. These products feed back into Reactions 1 and 2, thus ensuring a steady production of ozone! The rate of this reaction also depends on the temperature in the atmosphere.
Reaction 4: Ozone is degraded (broken down) by light energy, forming a charged form of singlet oxygen, O(1D), and more molecular oxygen. Notice that this reaction proceeds at a much slower rate than the first reaction (about one-fourth of one percent as slow!)
Reaction 5: the charged oxygen reacts with a catalyst to return to its normal state of singlet oxygen (which is, by the way, poisonous to breathe!)
Reaction 6: some of the charged oxygen reacts with water in the atmosphere to form a hydroxyl radical, OH . Radicals are fragments of molecules that have at least one unpaired electron, and are highly reactive. The hydroxyl radical, for example, is responsible for the majority of the chemical reactions that happen in the atmosphere during the day. Other radicals take control at night-time when there is no energy from the sun.
Reaction 7: carbon monoxide in the atmosphere, produced by fossil-fuel burning such as automobiles, reacts strongly with hydroxyl radicals to form carbon dioxide and HO2 radicals.
Reaction 8: the HO2 radicals formed in Reaction 7 react with the extra NO in the atmosphere to form more NO2 and more OH radicals. The rate of this reaction is dependent on the temperature in the atmosphere.
Reaction 9: the hydroxyl radicals react with NO2 to form nitric acid, HNO3, which will eventually be one of the culprits in the formation of acid rain, but that's (thankfully!) another model!
In summary: the development of photochemical smog is dependent upon solar radiation, source emissions of hydrocarbons and nitrogen oxides, and atmospheric stability (for enhanced concentrations). Early in the morning, commuter traffic releases NO and hydrocarbons. At the same time, NO2 may decrease through because the sunlight can break it down to NO and O. The O is then free to react with O2 to form O3. Shortly thereafter, oxidized hydrocarbons react with NO to increase NO2 by midmorning. This reaction causes NO to decrease and O3 to build up, producing a midday peak in O3 and minimum in NO. As the smog ripens, visibility may be reduced due to light scattering by aerosols. Primarily due to the dependence on commuter traffic between surburbs and cities, there are presently more than 40 urban areas in violation of the US ambient air quality standard for ozone.
Pure water is neutral and has a pH of 7.
Natural rain water is slightly acidic mainly because of dissolved CO2 which produces carbonic acid or H2CO3
H2O(l) + CO2(g) <==> H2CO3(aq)
The pH of unpolluted rainwater ranges from pH 5 to 6.
Acid rain is rain water with a pH of less than 5.
In some parts of the Northern Hemisphere the pH of the rain water has been as low as 2!
Acid rain is caused by caused by industrial pollutants.
The main industrial gases responsible are SO2 and NOx (a mixture of NO and NO2).
Major sources of industrial sulfur dioxide.
SO2(g) comes from mining smelters and the burning of coal.
i) The roasting of minerals releases SO2(g) from
Metal sulfide + oxygen ==> Metal oxide + SO2(g)
ii) Electrical power stations that burn coal produce sulfur dioxide from the sulfur impurities in the coal.
S(s) + O2(g) ==> SO2(g)
The SO2(g) combines with water to produce sulfurous acid.
H2O(l) + SO2(g) ==> H2SO3(g)
Note: Sulfur dioxide is not readily oxidized to sulfur trioxide in dry clean air. Water droplets and dust particles however, catalyse the reaction between O2 and SO2 in the air producing sulfur trioixde, SO3.This dissolves in water and produces sulfuric acid which is a much stronger acid. This can cause considerable damage to buildings, vegetation and fish populations by destroying fish eggs.
SO2(g) + ½O2(g) ==> SO3(g)
H2O(l) + SO3(g) ==> H2SO4(aq)
Source of nitrogen oxides
Sources of NOx are more widespread. Nitrogen is a diatomic molecule and is fairly inert because its triple bond. However, at temperatures over 1300°C, nitrogen combines with oxygen to form nitrogen monoxide.
N2(g) + O2(g) 2NO(g)
These high temperatures can be achieved by
i) the internal combustion engine (human activity)
ii) lightning in the atmosphere (natural source)
The nitrogen monoxide slowly combines with oxygen to form soluble nitrogen dioxide gas.
2NO(g) + O2(g) => 2NO2(g)
Nitrogen dioxide readily dissolves in water producing a mixture of nitric and nitrous acids.
2NO2(g) + H2O(l) ==> HNO3(aq) + HNO2(g)
Acidic rain is mainly caused by atmospheric pollutants of sulfur dioxide and nitrogen oxides.
The chemical formula of acidic rain is dependent upon the type of acids present. Acidic rain is a complex mixture of nitrous, nitric, sulfurous and sulfuric acids which all combine to lower the pH.
THE PRODUCTION OF ACID RAIN
Pure water is neutral with a pH of 7. Many gases are soluble in water. Some of these gases, most notably CO2 nitrogen oxides (NOx) and sulfur oxides (SOx), form acids when they dissolve in water. Rain water falling through a atmosphere containing these gases will absorb the gases and become more acidic. This is what we call acid rain. The nitrogen and sulfur oxides from car exhaust and industry are the most serious causes of acid rain, but it is dangerous to generate those gases in the lab. Carbon dioxide forms carbonic acid in water, which also makes water slightly acidic. It is easy, safe and fun to generate CO2 in the lab and demonstrate its effect when bubbled though water, as a simulation of acid rain formation. In this reaction:
vinegar + baking soda ² ¨ CO2 + sodium acetate + H2O
HOAc + NaHCO3² ¨ CO2 + NaOAc + H2O
In this particular experiment, we would like to show how certain gases dissolved in water can make the water more acidic.
- To have the students lean about chemical reactions with water
- To show the students how this relates to the environmental world
- To help the students see the reaction which will further their understanding of the topic
- 500 ml bottle, or any clean soda bottle 16 oz or smaller
- Baking soda
- Vinegar (White vinegar is a 5% solution of acetic acid.)
- 25 ml graduated cylinder
- Two l5O ml beakers
- Bromothymol blue indicator solution*
*This is easily made from a small amount ( spatula tip or an amount the size of a small peppercorn) of the bromothymol blue sodium salt in 10 ml of water). This should be more than enough for a class of 35 students working in units of two or three.
- Pour 50 ml of vinegar into the bottle
- Put about 40 ml of water into each beaker
- Add a few drops ( it may need as much as half a medicine dropper ) of indicator to the water in the beakers. It should turn the water blue or blue-green.
- Using the funnel put two heaping teaspoons of baking soda into the balloon
- Carefully place the end of the balloon over the top of the bottle, taking care to prevent the baking soda from falling in
- Once the balloon is in place, lift it up to allow the baking soda to fall into the bottle
- The carbon dioxide formed in the reaction between vinegar and baking soda should blow up the balloon
- Pinch or twist the balloon to save the gas, and while holding the gas in the balloon, take the balloon off the bottle
- Twist the end of balloon around one end of a straw. Then place the other end of the straw in the water in one of the beakers
- SLOWLY release the pinch or twist that is holding the gas in the balloon, allowing it to bubble into the water
The indicator should turn yellow, indicating that the solution has become more acidic. You can compare this to the color of the water of the control beaker.
The carbonic acid formed in water by the dissolution of carbon dioxide is not a strong acid and only makes the pH of the water around pH = 5. In fact, most water from the faucet is initially pH = 5-6 because of the CO2 which is dissolved in it. The nitrogen and sulfur oxides from car and industrial emissions form much stronger acids when they dissolve in water, making rain with a pH as low 3, an acid level very close to vinegar.
The students could go to the library or look on the internet and look for information on acid rain. Each student could then create a booklet on the new information that they have collected, including what they could do to help the problem.
There are several other labs in this booklet that have other experiments dealing with the effects of acid rain on the environment. These include "Effects of Acid Rain on Marble Statues", "Must it Rust? The Reaction Between Iron and Oxygen", and "A Green Penny?"
There are also labs in the booklet dealing with the water cycle and this experiment on acid rain could be used in conjunction with 'The Water Cycle" to show how the acid rain is effecting the environment.