HSC – Core Module 2: Acidic Environment

2. While we think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon nitrogen and sulfur. The concentrations of these acidic oxides have been increasing since the Industrial Revolution.

Identify oxides of non-metals that act as acids and describe the conditions under which they act as acids.

Oxides are chemical compounds formed when an element reacts with oxygen.

In school, a metal such as magnesium or a non-metal such as phosphorous can be burned in a gas jar of oxygen. The magnesium burns with a brilliant white flame to generate an oxide. In the case of magnesium a white powder is formed. This is magnesium oxide.

When red phosphorous is ignited in air or oxygen, it burns with a yellow flame to form a white smoke. Depending on the amount of oxygen present, oxides such as P4O6 (phosphorous (III) oxide) or P4O10 (phosphorous (V) oxide) will form.

The acid-base properties can be investigated using indicators. When magnesium is added to water containing universal indicator changes from green to blue-violet. This shows that magnesium oxide is a basic oxide. This reactions occurs because of the presence of water. The magnesium oxide dissolves in the water to generate magnesium hydroxide. The hydroxide ions that are released cause the colour change of the indicator.

When the white oxides of phosphorus are dissolved in water containing universal indicator, the mixture turns from green to red. Thus phosphorus oxides are acidic. Once again the presence of water is essential. The phosphorus oxides partly dissolve in the water to produce phosphoric acid. This acid releases hydrogen ions that cause the change in the colour of the indicator.

Some oxides, however, are not soluble in water, so indicators cannot be used to determine their acidity of basicity. In order to determine wether any oxide is acidic or basic we need to investigate wether it reacts with strong acids such as HCl or bases such as NaOH.

The reaction that occurs is classified as a neutralisation reaction. In the process of neutralisation the acidic and basic properties are destroyed and a new compounds called ‘salt’ is formed.

Neutralisation: The reaction between an acid and a base to produce a new compound that usually has neither acid nor basic properties.

Salt: The ionic compound formed when a base is neutralised by an acid.

Neutralisation:

Acid + Base Salt + Water

Basic Oxides

An oxide that dissolves and reacts with a strong acid is said to be a basic oxide. They are usually oxides of metals. They react with water to form a hydroxide ion. React with an acid to form a salt.

Example:

Copper (II) reacts and dissolves in hydrochloric acid and to form copper (II) chloride and water. Copper (II) chloride is the salt.

Acidic Oxides

An oxide that dissolves and reactions with a strong basic solution such as sodium hydroxide is said to be an acidic oxide. Usually oxides of non- metals, react with water to form an acid and react with a base to form a salt.

CO2

Another Example:

Sulfur trioxide reacts with sodium hydroxide solution to form sodium sulfate and water. Sodium sulfate is the salt.

Neutral Oxides

These do not react with acid or bases

REMEMBER:

CO (Carbon Monoxide)N2O (Dinitrogen Monoxide)NO (Nitrogen Oxide)

Amphoteric Oxides

In some cases, the oxide reacts with both strong acids and strong bases. Such oxides are classified as amphoteric.

Aluminium oxide is an example of an amphoteric oxide. It reacts with hydrochloric acid to produce aluminium chloride.

Al2O3: Basic Oxide

Aluminium oxide also reacts with sodium hydroxide solution to form sodium aluminate.

Al2O3: Acidic Oxide

Or

Generalisations

Common Acids and Names for Salts

Hydrohalic acids (such as HCl, HBr) when neutralised lead to the formation of salts known as halides (such as chlorides, and iodides).

Most of our common acids are what are called oxyacids: they have an oxygen attached to an element such as sulfur, nitrogen, phosphorous, chlorine or carbon. When the names of these acids end in ‘-ic’ the salt that is formed ends with ‘-ate’. When the names of these acids end in ‘ous’ the salts that they produce end in ‘-ite’Anions formed from oxyacids are called oxyanions.

Analyse the position of these non-metals in the Periodic Table and outline the relationship between position of elements in the periodic table and acidity/basicity of oxides.

The generalisations discussed allow us to demonstrate the trends in acidity and basicity of the oxides in the periodic table.

The metals of groups I and II all form basic oxides. The basicity of these oxides increases down each group. Thus barium oxide is a stronger base than magnesium oxide.

Example:

Most non-metals (other than the noble gases) form acidic oxides. The acidity of these oxides decreases down each group as the elements become more metallic in character.

In addition the non-metallic oxides with the highest oxidation states in the non-metal tend to be more acidic. Thus, sulfur trioxide (SO3) is more acidic than sulfur dioxide.

Eg:

On the reverse is a diagram of the periodic table showing the oxides formed and their respective nature.

In many cases an element may have a basic, acidic or amphoteric oxide due to the different oxidation states of the element.

Generally, the higher the oxidation state of the metal or semi-metal, the more amphoteric or acidic the oxide.

The effect of differing oxidation states on the acidity of oxides makes it difficult to make generalisations and to discuss trends in acidity and basicity in the oxides of the elements of the periodic table.

Periods 2 and 3 of the periodic table show interesting trends in the acid base properties of the oxides.

As we move from left to right the basicity of the oxides decreases until an amphoteric oxide is reached. Beryllium oxide and aluminium oxide are amphoteric oxides. The oxides then show increasingly acidic properties as we move along the period until group VII.

This trend is related to the decreasing metallic character of the elements across each period. The trend is also related to the type of crystal lattice that forms.

The basic oxides of period 3 are ionic compounds. Aluminium oxide lattice is also an ionic compound and is insoluble in water. The silicon dioxide lattice is a giant covalent network, and this is also very insoluble in water.

The remaining oxides all form covalent molecular lattices. These differences in bonding affect the behaviour of each oxide with water. The high solubility of sodium oxide allows the oxide ion to interact with the water and form hydroxide ions. Similarly, the high solubility of sulfur trioxide in water allows a chemical change to occur which leads to the formation of acidic hydrogen ions.

The two oxides that are very insoluble in water ( ) exhibit their amphoteric and weakly acidic properties respectively, only in the presence of strong acids and bases.

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen.

Our atmosphere is a sink for many pollutants. Some of these pollutants are oxides.

Carbon dioxide

Carbon dioxide in the atmosphere is part of the natural carbon cycle in which carbon moves through plants, animals, oceans and rivers, carbonate rocks and the atmosphere. However, in the past 150 years human

activity in the form of combustion of fossil fuels has led to an 30% increase in the concentration of carbon dioxide in the atmosphere – from about 280 ppm at the start of the industrial revolution to 360 ppm today.

Although it is not a pollutant, there is considerable concern because it is a major factor in the enhanced greenhouse effect.

Sulfur dioxide

Volcanoes, geysers and geothermal hot springs release many different gases into the atmosphere, including sulfur dioxide. The unpredictable amount of SO2 released by volcanic eruptions can produce a large amount of aerosols (very small solid or liquid particles suspended in air). Aerosol production can reduce the amount of solar radiation reaching the Earth’s surface.

Human activities that release sulfur dioxide into the atmosphere:- processing or burning fossil fuels- Extracting metals from sulfide ores

Coal generally contains 0.5-6% sulfur, mainly as metallic sulfides and sulfur in carbon compounds. It is not practical to extract sulfur from coal before burning, so in the combustion process most of the sulfur is converted into sulfur dioxide:

Although some power stations remove SO2 from the effluent gas, the alternative is to use low-sulfur coal, such as Australian. Hence there is great demand. Crude oil and natural gas may also contain a variety of sulfur compounds. Most of these are extracted at oil refineries and sold to H2SO4 manufacturers. However, there is also some release of SO2 by refineries. In addition small amounts of sulfur remains in petrol and so is released into the atmosphere as SO2 during combustion.

Many metals occur as sulfides, and the first step in extracting the metal is to roast the sulfide ore in air.

Carbon monoxide, carbon dioxide, and sulfur dioxide are also formed during the combustion of organic matter during bushfires.

Bacterial decomposition can also release sulfur dioxide into the environment. Lightning storms produce toxic nitrogen oxides as nitrogen and oxygen molecules react together.

The activities of human beings results in the release of oxides into the atmosphere. Industries generate a variety of oxides such as carbon monoxide, carbon dioxide, sulfur dioxide and the oxides of nitrogen.

Nitrogen Oxides

There are three common oxides of nitrogen: Nitrous Oxide (N2O), Nitric oxide (NO) and nitrogen dioxide (NO2)NO2 is acidic, the other two are neutral.

The major natural source of nitric oxide is lightning. This slowly reacts with oxygen to form nitrogen dioxide

Combustion, both in stationary sources (power station) and moving ones (vehicles) is the major source of nitric oxide and nitrogen dioxide. At the high temperatures in the combustion chambers oxygen and nitrogen react to form nitric oxide, which is then released and eventually forms nitrogen dioxide.

Nitrous oxide is also formes naturally by the action of certain bacteria on nitrogenous materials in the soils. Human activity has increased the release of nitrous oxide to the atmosphere through increased use of nitrogenous fertiliser which provides more raw material for the bacteria. This is of considerable concern because nitrous oxide contributes to the enhanced greenhouse effect.

The activities of human beings results in the release of oxides into the atmosphere. Industries generate a variety of oxides such as carbon monoxide, carbon dioxide, sulfur dioxide and the oxides of nitrogen.

Describe, using equations, examples of chemical reactions that release sulfur dioxide and chemical reactions that release oxides of nitrogen.

Three acidic oxides are minor constituents of the atmosphere:Carbon dioxide – CO2 – 360 ppmSulfur Dioxide – SO2 – 0.001 ppmNitrogen Dioxide – NO2 – 0.001 ppm

Oxides of Carbon

The natural cellular respiratory process of animals and plants produce carbon dioxide, which is ultimately released into the atmosphere. Green plants utilise the carbon dioxide for photosynthesis. Large amounts of carbon monoxide and carbon dioxide are also produced during bushfires.

The process of incomplete and combustion produces carbon monoxide, whereas carbon dioxide is the product of complete combustion. The combustion of fossil fuels such as petrol, kerosene and diesel oil releases large amounts of carbon oxides into the atmosphere. Coal-fired power stations also release vast amounts of carbon oxides into the atmosphere.

The increasing levels of carbon dioxide in the atmosphere are often linked to global warming. Carbon monoxide does not build up over time in the atmosphere as it rapidly removed by the actions of soil organisms or be oxidation to carbon dioxide.

Oxides of Sulfur

The oxides of sulfur are irritating, poisonous gases. These gases particularly affect people who suffer from respiratory problems such as asthma and emphysema.

About half of the sulfur dioxide that enters our atmosphere is derived from the oxidation of hydrogen sulfide (H2S) produced from bacterial decomposition. The remainder is produced by human activity.

Sulfur dioxide is produced by the combustion of fuels such as coal and diesel oil.

In Australia considerable quantities of coal are burnt to generate most of our electricity (about 75%). Fossil fuels usually contain small quantities of sulfur minerals (eg. FeS2) although the bituminous coal from Eastern NSW is lower in sulfide minerals than coal from other locations. These sulfide minerals in coal are oxidised during combustion, and sulfur dioxide is released.

Metal smelters that convert sulfide minerals into metals are also a major source of sulfur dioxide pollution. For example, the smelting of chalcopyrite (CuFeS2) during the production of copper results in the release of sulfur dioxide. Towns near these smelters can suffer from sulfur dioxide pollution. In regions such as the lower Hunter, about 95% of SO2 emissions are associated with power and smelting industries.

Sulfur trioxide is produced mainly by oxidation of sulfur dioxide in the atmosphere. Dioxygen and ozone are the oxidants involved in this oxidation process

Many metals occur as sulfides and the first step in extracting the metal is to roast the sulfide ore in air. This produces sulfur dioxide along with the metal or metal oxide. Eg: Zinc Sulfide:

Under certain conditions, bacteria may decompose organic matter to produce hydrogen sulfide (H2S). When H2S is oxidised, sulfur dioxide is formed, according to the equation:

Oxides of Nitrogen

Nitric oxide (NO) and nitrogen dioxide (NO2) are the most common atmospheric oxides of nitrogen found in urban air. They are collectively referred to as NOx.

The major natural source of nitric oxide is lightning.

In Sydney, about 86% of the total emissions of NOx comes from the engines of motor vehicles and other transport. The nitric oxide (a colourless gas) is formed when nitrogen and oxygen react at high temperatures. Nitric oxide is a neutral oxide.

NO2 is a brown gas and is produced by the oxidation of NO. The rate of NO2 formation depends on the concentration of NO in the air. Up to 10% of the total NOx in the air is NO2.

Nitrogen dioxide is an acidic oxide.

NOx is also derived from other sources including industries, electrical power production and oil refining. NO2 is of concern as it causes damage to the respiratory system o humans as well as irritating the eyes. Young children and the elderly are particularly susceptible. Un-flued gas and kerosene heaters also contribute to NO 2 pollution of the air inside houses. An un-flued gas heater is one that burns gas, but has no flue or chimney to take combustion products outside; instead, the products are released directly into the room.

Health Effects

Sulfur dioxide irritates the respiratory system and causes breathing difficulties even at low concentrations (1 ppm). People suffering from asthma and emphysema are particularly susceptible. The effects of sulfur dioxide are magnified if particulates are present also.

Nitrogen dioxide irritates the respiratory tract and causes breathing discomfort at concentrations above about 3 to 5 ppm and at higher concentrations does extensive tissue damage. Concentrations above 3 ppm have rarely been reached even in heavily polluted cities.

The main problem with NO2 is that it leads to the formation of ozone in what is called photochemical smog. This is a form of air pollution in which sunlight acts upon nitrogen dioxide in the presence of hydrocarbons and oxygen to form ozone which results in poor visibility due to small particles in the air. Ozone has harmful effects at concentrations as low as 0.1 ppm.

Environmental effects

The major environmental effects of emissions of oxides of sulfur and nitrogen to the atmosphere are those coming from the acid rain that these oxides eventually lead to.

Releases of sulfur dioxide to the atmosphere in industrialized cities up until the middle of the twentieth century and in mining and smelting areas until much more recently produced air that often had an unpleasant odour and which was detrimental to health, particularly for aged people and those with respiratory weaknesses.

Releases of oxides of nitrogen leads to formation of photochemical smog which is both visually unattractive and a health hazard.

Assess the evidence that indicates increases in atmospheric concentration of oxides of sulfur and nitrogen.

Monitoring SO2 and NO2 levels

During the industrial revolution of the 19th century, large amounts of coal were burnt to provide power for factories and their machines. Vast quantities of carbon dioxide and sulfur dioxide poured into the air. Iron smelters generated large volumes of sulfur dioxide as they produced the growing quantities of steel required for industry. The atmosphere of large industrialised cities in Europe and the USA became highly polluted with acidic gases.

The increased use of motor vehicles in the 20 th century (particularly after 1945) increased oil consumption. Emissions of sulfur dioxide doubled in the 25 year period following Word War II. Adding to this pollution burden atmosphere was the increasing production of nitrogen oxides in internal combustion engines.

Following numerous deaths (about 4000) in London due to heavy acidic smogs, pollution controls began to be introduced to clean up the air of these large cities. High levels of photochemical smog (produced by the action of sunlight on air containing moisture, ozone, hydrocarbons and NOx) in cities such as Los Angeles and Tokyo in the 1960s accelerated the push for emission controls on motor vehicles. In the 1970s the development of more sensitive gas analysis technologies allowed chemists to monitor the global increase in sulfur dioxide emissions due to the expansion of industries in Asia.

In recent years the air quality has improved in most westernised countries. In Europe, the sulfur dioxide emissions dropped by about 45% in the 1990’s. In the same period, nitrogen dioxide emissions dropped by about 20%. However, due to increasing population and usage of motor vehicles, the levels of population have stabilised rather than continuing to decrease. The rapid industrialisation of Asia (particularly China) has led to huge increases in sulfur dioxide emissions. It is predicted that sulfur dioxide emissions in Asia will triple in the 20 year period from 1990.

The Environmental Protection Agency of NSW (EPA) monitors the levels of pollutant gases in the atmosphere in many regions across NSW.

The 2003 Report reveals the following information:

- The peak 1-hour measurements of sulfur dioxide in Sydney air are now less than 25% of the National Environmental Protection Council (NEPC) standard of 0.20 ppm.

- NO2 concentrations have remained fairly stable over the last decade, with concentrations rarely above the NPEC 1-hour peak standard of 0.12 ppm. In Sydney, typical levels have been much lower than 0.08 ppm over the last decade. This compares favourably to the 1980’s when the standard was exceeded over many days in the winter.

Data from the NSW Department of Environment and Conservation shows that air quality in Sydney with respect to NO2 and SO2 levels is high compared to NPEC standards.The air in Sydney however compares favourably with that of much larger cities and overseas. The large industrialised cities of China, the USA and Europe have much higher levels of atmospheric NO 2 and SO2. In most major centres in Europe and the USA the emission controls that have been instituted by government environmental agencies have improved air quality since the 1950’s.

In order to reduce emissions of sulfur dioxide and nitrogen oxides there are a number of options available. These include:

- Reducing the amount of coal being burnt – this can be achieved it people used appliance that are more energy efficient; as individuals we can help by reducing our own personal use of electricity.

- Power companies burning low-sulfur coal rather than high-sulfur coal – they can also switch to natural gas, which produces very little sulfur dioxide on combustion.

- Reducing our reliance on fossil fuels – developing other forms of power production such as hydroelectricity and solar power.

- Collecting the sulfur dioxide produced by smelting metal sulfides and using it to make sulfuric acid.

- Reducing acidic emissions from smoke stacks by a process called ‘scrubbing’ – the acidic gases are passed through a slurry of a base such as calcium oxide; the sulfur dioxide reacts with the calcium oxide to form calcium sulfate.

- Ensuring that exhausts from motor vehicles pass through a catalytic converter – nitrogen oxides are then converted back to nitrogen gas.

Evidence:

There is extensive evidence for an increase of over 25% in atmospheric carbon dioxide levels over the last two hundred years. The evidence comes from quantitative analysis of trapped air bubbles in Antarctic ice and measurement of carbon isotopes in old trees, grass seeds in museum collections and calcium carbonate in coral.

Holes are drilled in Antarctic ice where these gases formed air bubbles. New ice formed is up the top, so the layers of ice further down are older, whereas the layers further up are newer. By analysing the amount of oxides of sulfur and nitrogen at different levels, trends are observed which indicate an increase in the atmospheric concentration of the oxides.

It has been observed that the damage to buildings, forests, and aquatic organisms is increasing, and this is thought to be due to acid rain, which is formed by the sulfur and nitrogenous oxides reacting with water in the atmosphere. Hence the increase in acid rain implies in an increase in levels of these oxides in the atmosphere. Also, there has been an increase in the acidity of rivers and lakes. This can be seen through the damage to the Black Forest in Germany due to acid rain and the hundreds of lakes in Canada with no living things in them.

Higher atmospheric concentrations of sulfur and nitrogen oxides have been detected in industrial areas as compared to non-industrial areas. Note that Sydney’s air does not suffer the extreme levels of pollution found in large industrialized centres of Europe and the USA. This is part due to low sulfur content in Australian coal.

Finding evidence for increases in atmospheric sulfur oxides and nitrogen oxides is more difficult for the following reasons:

- Whereas atmospheric CO2 concentrations are about 360 parts per million (ppm), the levels for SO2 and NOx are only about 0.001 ppm in populated parts of the Earth. Thus it is much more difficult to detect. In more recent times Infrared spectroscopy has actually allowed us to measure and confirm that these oxides are continuing to increase on a annual rate.

- The chemical instruments able to measure very low concentrations, like those for SO 2, have only been commercially available since the 1970s. Hence results obtained before this period may have not been accurate.

- CO2 changes to carbonate ions when it dissolves in water and most carbonates are insoluble. Seashells and coral are made up of carbonates that came from atmospheric CO2. Isotope ratio measurements using mass spectrometers on shells and corals of different ages give clues as to past atmospheric CO2 concentrations. On the other hand, SO2 eventually forms sulfate ions and NO2 forms nitrate ions. Most sulfates and all nitrates are water-soluble. Soluble sulfates and nitrates circulate in the hydrosphere and biosphere and are chemically changed while insoluble carbonates tend to stay in inert forms such as shells or coral. Sulfur dioxide and nitrogen dioxide are readily soluble in water, as mentioned, which means they can circulate in the biosphere and hydrosphere and thus also become chemically changed. This results in any atmospheric measurements of their presence inaccurate and their validity questionable.

Although accurate measures of atmospheric oxides of sulfur and nitrogen are at present non-existent, there is sufficient indirect evidence to conclude that significant increases in the atmospheric concentrations of oxides of sulfur and nitrogen have indeed taken place, especially after industrialization.

Furthermore, measurements of carbon isotope contents of old trees, grass seeds collected over hundreds of years and the calcium carbonate in coral, all point to a steady increase in atmospheric carbon dioxide since the industrial revolution.

Specfic Example:

During a black out in Canada and the US beginning on 14 august 2003 the entire electrical grid was shut down. During the time that there was no burning of coal in power stations, the atmosphere became much clearer and cleaner as a result. Which was evidence due to the reduction of the brown layer in the sky (known as photochemical smog). Hence this means that levels are increasing due to the ever increasing amount of fuel being used and thus release of emissions.

Explain the formation and effects of acid rain

Rain or precipitation, which is more acidic than ‘normal’ rain is referred to as acid rain regardless of whether is has natural or man-made causes.

When all gases are removed from pure water, its pH is 7.0 (neutral) at 25 C. However, natural water contains dissolved gases including carbon dioxide, which makes the water weakly acidic (due to the presence of carbonic acid). This has a typical pH of 6.0 – 6.5.

When the atmosphere is polluted with acidic oxides such as sulfur dioxide and nitrogen dioxide, rainwater can become quite acidic (pH 4.0 – 5.0) due to the high solubility of these gases in water. Rain that has such a low pH is called acid rain. Acid rain with a pH of 3.6 has been recorded in many severely polluted industrialised areas in Europe and the USA.

The term acid rain is usually defined as rain that has a pH lower than 5.

Acid precipitation, had been observed in the seventeenth century and its causes related to industrial emissions.

When oxides of sulfur and nitrogen dioxide dissolve in water, they produce solutions of various acids. Sulfur dioxide forms weak sulphurous acid (H2SO3) whereas sulfur trioxide produces strong sulfuric acid (H2SO4). Sulfurous acid can be catalytically oxidised to produce sulfuric acid.

Nitrogen dioxide produces weak nitrous acid (HNO2) and strong nitric acid (HNO3) when it dissolves in water. In the presence of water and oxygen the nitrous acid is catalytically oxidised to nitric acid.

Dry Deposition

The process of ‘dry deposition’ is also included under the umbrella of acid rain. In areas where the weather is dry, the acid chemicals may become incorporated into dust or smoke and fall to the ground through dry deposition, sticking to the ground, buildings, homes, cars. In this process the oxides of sulfur and nitrogen may be absorbed by or directly onto surfaces. These oxides may then be converted into acids when the surface becomes moist through contact with rain, fog, mist or dew.

Example: Due to the large surface are of foliage in forests, they are especially effective at absorbing ‘dry’ atmospheric gases and particles. When rain occurs, or there is significant moisture in the air, the dry oxides are converted into acids which destroy the plant. Sulfur dioxide is a prominent example of this, whereby hardwood forests of Tennessee (in the US) have been destroyed due to dry deposition.

Wet Deposition

Wet deposition refers to acidic rain, fog, and snow. If the acid chemicals in the air are blown into areas where the weather is wet, the acids can fall to the ground in the form of rain, snow, fog, or mist. As this acidic water flows over and through the ground, it affects a variety of plants and animals

Effects of Acid Rain:

When you visit the ancient cities of Europe you may observe that many marble statues and building facades are eroded. Acid rain is partly responsible for this problem. When the calcium carbonate of the marble is attacked by the sulfuric acid in acid rain, the surface of the marble is converted into insoluble calcium sulfate.

Acidic rain can also attack metallic structures composed of iron and steel. The iron is oxidised by the hydrogen ions in the acid and becomes chemically weathered.

Acid rain has had a devastating affect on many northern hemisphere forests. The famous Black Forest in Germany has bee significantly damaged by acid rain. Pine needles lose their waxy coating and turn brown and the trees become denuded of foliage. In the 1980’s up to 50% of European forests had been affected by acid rain. The sugar marples of Quebec have also been damaged by acid rain that as has formed due to sulfur dioxide emissions from the industrialised north-east of the USA.

Acid rain also affects the soil, and the acidified soils inhibit the growth of plant seedlings. This can reduce crop production.

Basic minerals in the soil (such as dolomite and limestone) are attacked and dissolved by acidic water. Many types of sandstone have grains that are cemented together with calcite (calcium carbonate). The acid rain will dissolve this cement and so cause significant chemical weathering and erosion.

Mineral nutrients such as potassium, calcium and magnesium that are required for plant growth can be removed when acid rain soaks into the ground. Some insoluble minerals can also be dissolved by the acidified water and cause a release of toxic levels of metal ions. High levels of aluminium ions (formed when clay minerals are attacked by acid) in the soil interfere with normal mineral uptake by plant roots.

Soil chemistry can be changed making it difficult for plant roots to take up essential elements such as calcium and potassium.

Nitrogen fixing bacteria that are vital to a healthy soil also severely affected by increasing acidification. This also includes the fact that micro-organisms in the soil, which are essential for recycling material are killed

Changes in soil Ph can cause reduced productivity and in some cases the plant will not grow at all.

Smog and acid rain can combine to be a human health hazard.

A 1997 study showed that around the world acid rain between 4 and 5 pH was most common in areas where there are fewer controls on sulfur dioxide emissions from industries. When acid rain drains into lakes, it can significantly acidify them.

Lakes have become acidified due to the sulfur dioxide emissions carried from the industrialised zones of Eastern Europe. As the lake water becomes more acidic the populations of aquatic organisms become stressed. The presence of acidic hydrogen ions interferes with the carbon dioxide / carbonate equilibrium in the water. The amount of dissolved carbon dioxide in the water drops as carbonate ions are removed. This leads to stress on photosynthetic organisms. Many aquatic invertebrates cannot reproduce in an acidic environment. Most fish eggs will not survive in water if the pH drops below 5.5. Below a pH of 5 adult fish will die as they cannot extract enough calcium from the water to maintain their skeletons.

Studies in the 1970’s showed that acidified lakes also had high levels of toxic heavy metals. The increased acidity of the lakes leads to a chemical leaching of heavy metals from bedrock and soil into the lakes. Thus the death of many fish can also be attributed to heavy metal poisoning. The drop in fish numbers affects all parts of the food chain. In some locations, small acidified lakes have been restored to normal by large scale ‘liming’. This involved adding large amounts of hydrated calcium oxide or sodium carbonate to the lake to raise the pH back to 7.

It has been estimated that by the mid twenty first century, the world’s oceans will have become significantly acidic due to the absorption of the acidic greenhouse gas carbon dioxide.

There is also concern regarding the effects of acid rain. The food chain may be affected through biological magnification of toxic metals and other substances dissolved by acid rain. Acidic drinking water can dissolve metals such as copper and lead often present in pipes in plumbing. Sulfur and sulfur dioxide are known to be harmful if inhaled, especially by people with respiratory problems.

Discuss Issues associated with programs to reduce acid rain in industrialised cities

Industrialisation has been associated with the use of energy for industry, mining and transport. In most countries, this energy is derived from the combustion of fossil fuels. Their combustion produces produce acidic gases (CO2, SO2, NO2) that lead to the formation of acid rain, an environmental hazard. Acid rain causes chemical erosion and disturbs food chains.

Currently there is little choice for industry other than to continue to use fossil fuel as energy sources. Solar energy and other alternative energies cannot currently provide the level of energy for industrialised societies. Consequently, the only achievable programs are those in which acidic gases are removed from the combustion gases before they are vented into the environment.

Much research is currently under way to reduce the emissions of sulfur and nitrogen oxides into the atmosphere. For sulfur oxides this includes reducing the sulfur content of fuels and removing sulfur oxides from atmospheric emissions using a scrubber. A scrubber removes sulfur dioxide by passing the gaseous emissions through a saturated solution of magnesium hydroxide, Mg(OH)2 or calcium carbonate, CaCO3. For example, sulfur dioxide reacts with magnesium hydroxide as shown in the following equation:

The magnesium sulfite produced must be disposed of as landfill or converted by another chemical process back to magnesium hydroxide.

Nitric oxide emissions from motor vehicles are reduced by the use of catalytic converters, which catalyse the decomposition of NO back to N2 and O2.

Define Le Chatelier’s Principle

Not all reactions proceed to completion. In many reactions, the final reaction mixture consists of both products and reactants.

Chemical Equilibria are dynamic and not static. All systems in chemical equilibria exhibit the following characteristics.

- The system is closed. This means that atoms molecules or ions cannot enter or leave the system. Thus the reaction of calcium carbonate with HCl in a beaker is not a closed system as the CO 2 escapes. Water in a capped bottle will not continue to evaporate, because and equilibrium eventually becomes established between the liquid water and the water vapour.

- The observable or measurable properties of the system are constant. These properties are often called macroscopic properties and they include colour, electrical conductivity, concentration and gas pressure.

- The concentrations of reactants and products are constant once equilibrium has been achieved.

- The rate of the forward reaction equals the rate of the reverse reaction, hence the dynamic system.

Le Chatelier’s Principle

If a system is at equilibrium and a change is made that disturbs the equilibrium, then the responds in such a way as to counteract the change and eventually a new equilibrium is established.

Identify factors which can affect the equilibrium in a reversible reaction

The factors that may cause a disturbance to a chemical system in equilibrium are:

- Temperature- Concentration- Total Gas pressure- pH (only applies to some cases such as carbon dioxide + water)

Temperature

Reactions can be classified as endothermic and exothermic on the basis of whether heat is absorbed or released in the forward reaction. Using Le Chatelier’s principle, the following predictions can be made:

- For endothermic equilibria, an increase in temperature causes the equilibrium to shift to favour the products.

- For exothermic equilibria, an increase in temperature causes the equilibrium to favour the reactants.

For an Exothermic Reaction:

Heat is a product of the forward reaction. If the tube is cooled, heat is removed. This disturbs the equilibrium and so a change occurs to oppose the disturbance. The equilibrium moves towards the right to make more heat; consequently more products are formed.

If the tube is heated, the system responds to remove some of the added heat by shifting the equilibrium to the left. Thus more reactants are formed.

Concentration:

Increasing or decreasing the concentration of reactants or products can shift the position of a chemical equilibrium. Using Le Chatelier’s principle, the following predictions can be made:

- Increasing the concentration of reactants or decreasing the concentration of products shifts the equilibrium to favour the products.

- Decreasing the concentration of reactants or increasing the concentration of products shifts the equilibrium to favour the reactants.

Example + Explanation:

When more reactants are added, there is a significantly higher concentration of reactants within the beaker. Thus the system responds to counteract the change. In doing so, the equilibrium shifts to the right to use up some of the additional reactants. Thus more products are formed.

When the solution is diluted with products, the disturbance causes the equilibrium to shift to counteract the change. Thus the equilibrium lifts to the left to use up some of the added products. Thus more reactants are formed.

Gas Partial

pressure:

In a mixture of gases, the total gas pressure is the sum of the individual pressures of each component gas. This individual gas pressure is called the partial pressure of the gas. Partial pressure is therefore a measurement of the concentration of as within the total gas mixture. Using Le Chatelier’s principle the following predictions can be made:

- Increasing the partial pressure of the reactants gas shifts the equilibrium to favour the products.

- Increasing the partial pressure of a product gas shifts the equilibrium to favour the reactants.

For example:

When there is an injection of reactant gas, the partial pressure has disturbed equilibrium; to counteract the change, the equilibrium has shifted to the right to reduce the amount of reactant gas in the mixture.

Total gas pressure:

A simple way to investigate the effect of changing the total gas pressure on a gaseous equilibrium is to use a gas syringe as a reaction vessel. The has mixture to be examined is placed in the syringe. The volume of has can be increased or decreased by moving the plunger. It has been proven that the pressure of gas (at a constant temperature) is inversely proportional to its volume. Thus if the volume of gas is decreased by compression then its pressure will increase.

Example:

Say we had a syringe of some brown gas, and instantly we doubled the volume (the gas partial pressure is halved). The syringe will lighten initially due to he doubling of the volume and the spreading out of gaseous particles. This change has disturbed equilibrium and halved the partial pressures. The system responds by shifting to the left to counteract the change. The shift to left increases the number of particles in the syringe and as these will be brown, the brown colour intensifies.

Consider the equilibrium system:

If the volume of the system is halved by doubling the external pressure, the total concentration of particles in the system is increased.

According to Le Chatelier’s principle, if the total concentration of particles is increased, the system will adjust to re-establish equilibrium in such a way as to partially decrease the concentration of gaseous particles in the system.

In re-establishing equilibrium, this means that the reverse reaction is favoured. For every two molecules of NO2 that react, only one molecule of N2O4 is produced, therefore leading to a decrease in the total number of particles.

This partially counteracts the increased concentration of particles produced by halving the volume of the equilibrium system.

When equilibrium is re-established, there will be more N2O4 and less NO2 in the system than before the change was made.

However, the concentrations of both are greater than at the initial equilibrium.

If the volume of the equilibrium system had been increased, the immediate effect would have been to decrease the total concentration of particles.

The system would re-establish equilibrium by favouring the reaction producing an increase in the number of gaseous particles.

From the equation it can be seen that this would favour the production of NO2 and lead to a decrease in the amount of N2O4.

Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s Principle

Carbon dioxide is an acidic oxide. It dissolves water to produce a solution of carbonic acid (H 2CO3). Carbonic acid is in equilibrium with hydrogen ions and hydrogen carbonate ions as shown below:

This dissolution reaction is exothermic.

Carbonated soft drinks

Soda water and other carbonated soft drinks are manufactured by dissolving carbon dioxide in water under pressure. The water is supersaturated in carbon dioxide. The system will remain this way only if the bottle is sealed so as to maintain the high carbon dioxide gas pressure in the gas space under the cap. The increase in carbon dioxide gas pressure has significant effects of the gaseous equilibria.

Equilibrium (1) is shifted to the right to reduce the gas pressure. Thus more carbon dioxide is dissolved. This in turn affects equilibrium (2), which is also pushed to the right to make more carbonic acid. The increase in

carbonic acid shifts equilibrium (3) to the right as well, hence the acidity increases as more hydrogen ions are formed. That is why soda water tastes sour. It has a pH of around 4.

When the cap of a soft drink is unscrewed, there is a rapid effervescence observed. The soft drink eventually goes ‘flat’ and tastes less sour due to the reversal of these three equilibria. The escape of gaseous carbon dioxide has caused equilibrium (1) to shift to the left in an open system. The other two equilibria in turn shift to the left.

Equilibrium under various conditions:

1) Changing the pH

Variation: Drops of lemon juice are injected into soda waterObservation: The soda water degassesExplanation: The addition of the lemon juice increased the hydrogen ion concentration of the soda water. This caused the carbonic acid equilibrium to shift to the left to use up some of the added acid.

The increase in carbonic acid levels causes the following two equilibria to reverse and produce more gaseous carbon dioxide.

The opposite can occur if we make the solution more alkaline. If OH- is added into the solution, then the carbonic acid reacts with it:

This effectively removes the product from reaction ( ). The equilibrium then moves to try to counteract this, so more CO2 dissolves to form more H2CO3 which in turn reacts with OH- and so on.

2) Changing the temperature

Variation: The soda water has warmed (in a gas syringe)Observation: The soda water degasses.Explanation: The dissolution of carbon dioxide in water is exothermic. Heat is therefore a product of the reaction.

If heat is added, the equilibria will shift to the left to counteract the change. This causes the soda water to de-gas.

This is also observable due to the decreased solubility of carbon dioxide in water, as the temperature of the water increases.

3) Addition of Salt (NaCl)

4) Shaking

Shaking a bottle of drink when the cap is removed speeds up the release of carbon dioxide and so the drink froths up. Equilibrium is reached more rapidly.

This is due to the kinetic energy gained by the particles in the liquid. As the amount of concentration of carbon dioxide within the soft drink is greater than that of the surrounding air, when the particles are given excess kinetic energy, it results in more collisions which shifts all three equilibriums to the left, hence carbon dioxide is lost more rapidly. I.e. degassing occurs.

The reaction between carbon dioxide and water is of great practical significance. It is involved in the removal of carbon dioxide from our bodies, the transport of carbon dioxide in photosynthesis, the removal of carbon dioxide from the atmosphere and the preparation of aerated drinks.

The reaction between carbon dioxide and water will never go to completion. Hence equilibrium is achieved.

The CO2, HCO3 equilibrium and aerated drinks

Aerated (fizzy) drinks contain carbon dioxide. To make such drinks, the solution is exposed to a high pressure of carbon dioxide. This causes a significant concentration of CO2 to dissolve; the bottle is then sealed. The gas space above the solution still contains a high pressure of carbon dioxide. When the bottle is opened the high pressure of CO2 escapes. The reactant concentration (pressure) has been greatly lowered and so the equilibrium will move to the left, whereby aqueous CO2 is converted gaseous form. If the unopened bottle is exposed to the atmosphere, then the CO2 pressure above the solution falls to normal atmospheric CO2 levels, and the concentration in solution falls dramatically – hence the drink has gone flat.

Note:

It is dangerous to put bottles of aerated drinks in a freezer, because as the temperature falls, pure ice freezes out of the drink. This means that all the CO2 originally dissolved in the drink is forces as a gas into the small gas space. This builds up an excessive pressure and the bottle may explode. In other words, carbon dioxide has very low solubility in solid ice.

5) Effect of pressure

In a closed system such as a bottle of soft drink, CO2(g) is in equilibrium with CO2(aq), which is in equilibrium with H2CO3.

To increase the solubility of carbon dioxide, the gas above the solution in the bottle is enriched in carbon dioxide at higher than normal pressure.

This increased pressure (concentration) of carbon dioxide increases the solubility of the gas and hence increases the acidity of the solution.

When a bottle of soft drink is opened, the pressure of carbon dioxide gas decreases and the equilibrium adjusts by favouring the reaction direction that increases the pressure of carbon dioxide.

As a result, bubbles of carbon dioxide gas can be seen coming out of solution. The soft drink becomes ‘flat’ as gaseous carbon dioxide escapes.

analyse information from secondary sources to summarize the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment.

Oxides of Sulfur

About half of the sulfur dioxide that enters our atmosphere is derived from the oxidation of hydrogen sulfide (H2S) produced from bacterial decomposition.

The remainder is produced by human activity. The industrial origins of sulfur dioxide are emitted from power plants, refineries, some chemical plants, pulp mills, extracting metals from sulfide ores and marine vessels. It also is emitted when coal and oil is burnt or processed.

Sulfur dioxide is produced by the combustion of fuels such as coal and diesel oil.

In Australia considerable quantities of coal are burnt to generate most of our electricity (about 75%). Fossil fuels usually contain small quantities of sulfur minerals (eg. FeS2) although the bituminous coal from Eastern NSW is lower in sulfide minerals than coal from other locations. These sulfide minerals in coal are oxidised during combustion, and sulfur dioxide is released.

Metal smelters that convert sulfide minerals into metals are also a major source of sulfur dioxide pollution. For example, the smelting of chalcopyrite (CuFeS2) during the production of copper results in the release of sulfur dioxide. Towns near these smelters can suffer from sulfur dioxide pollution. In regions such as the lower Hunter, about 95% of SO2 emissions are associated with power and smelting industries.

Industrial Origins of sulfur dioxide

Many metals occur as sulfides and the first step in extracting the metal is to roast the sulfide ore in air. This produces sulfur dioxide along with the metal or metal oxide. Eg: Zinc Sulfide:

Oxides of Nitrogen

Nitric oxide (NO) and nitrogen dioxide (NO2) are the most common atmospheric oxides of nitrogen found in urban air. They are collectively referred to as NOx.

The major natural source of nitric oxide is lightning.

In Sydney, about 86% of the total emissions of NOx comes from the engines of motor vehicles and other transport. The nitric oxide (a colourless gas) is formed when nitrogen and oxygen react at high temperatures. Nitric oxide is a neutral oxide.

NO2 is a brown gas and is produced by the oxidation of NO. The rate of NO2 formation depends on the concentration of NO in the air. Up to 10% of the total NOx in the air is NO2.

Nitrogen dioxide is an acidic oxide.

NOx is also derived from other sources including industries, electrical power production and oil refining. NO2 is of concern as it causes damage to the respiratory system o humans as well as irritating the eyes. Young children and the elderly are particularly susceptible. Un-flued gas and kerosene heaters also contribute to NO 2 pollution of the air inside houses. An un-flued gas heater is one that burns gas, but has no flue or chimney to take combustion products outside; instead, the products are released directly into the room.

Reasons for Concern:

The oxides of sulfur are irritating, poisonous gases. These gases particularly affect people who suffer from respiratory problems such as asthma and emphysema.

Nitrogen oxides have been increasing in abundance in the atmosphere. These gases are significant aspects of the greenhouse affect, thus it is in the best interest of the world that these emissions be cut down.

Health Effects AND environmental effects as mentioned previously

Effect due to acid rain as mentioned previously.

Social Cost: Due to the effects of acid rain, odours and photochemical smog, the tourist industry in a particular area is reduced dramatically due to the destruction of iconic structures (man-made or natural). This has a negative impact on the economy, thus plans are in place to reduce the amount of sulfur dioxide present in the atmosphere. For example due to the Acid Rain Programs present throughout many states in America and other countries, the amount of SO2 release is monitored in order to prevent significant environmental damage.

Economic Cost: The actual cost in repair and maintenance of structures, soils, and people that have been affected by SO2 emissions (either through acid rain or direct contact) is massive. This puts financial strain on political parties, thus it is in their best interest to avoid these costs by reducing the amount of SO2 produced.

Overall, there is significant concern of the release of SO2, and there should be due to the negative consequences it has.

Class Notes on Equilibrium:

Le Chatelier’s Principle:

If a system is in equilibrium and a change is made to disturb the equilibrium, the system responds in a way to counteract that disturbance and hence establish a new equilibrium.

The factors that may cause a disturbance to a chemical equilibrium are:- Concentration (or gas partial pressure)- Pressure- Temperature

Catalysts do not alter the position of the system when at equilibrium, but will reduce the time it takes to get to equilibrium.

Changes in Concentration

- Solids and pure liquids- Concentration does not change- [solids and pure liquids] =1

Concentration can be increased by:- Adding more reactants/products- Increasing pressure (gas only)- Decreasing volume (gas only)

Concentration can be decreased by:- removing reactants/products- adding more water (aqueous solutions)- Decreasing pressure (gases only)- Increasing volume (gases only)- Adding a substance which favours the reaction that uses it up- Removing a substance which favours the reaction which produces it.

Add: Reactant – Shifts position of equilibrium to the right; favours the productsProduct – Shifts position of equilibrium to the left; favours the reactants.

Remove: Reactant – Shifts position of equilibrium to the left; favours the reactants Product – Shifts position of equilibrium to the right; favours the products.

Relating this to collision theory

As the concentration of a reactant is increased, the frequency of successful collisions of the reactants would increase also, allowing for an increase in the forward reaction and the generation of the product.

Changes in pressure:

Only applies to equilibrium reactions which involve one or more gases.

Pressure can be increased by:- Adding more of the gas involved- Decreasing the volume of the system – ie. Compressing the gas

Pressure can be decreased by:- Removing some of the gas involved- Increasing volume

The effect of change in pressure has on an equilibrium system is determined by the number of gaseous molecules on each side of the equation.

Increasing pressure: Shifts equilibrium to the reaction side with the least gaseous moleculesDecreasing pressure: Favours the side with the most gaseous moleculesChanges in Pressure

If A + B -> C is exothermic then, C -> A + B is endothermic.

Note: Delta H refers to the forward reaction.

Increasing temperature: Favours the endothermic reaction

Decreasing temperature: Favours the exothermic reaction

Graphically, a temperature change is shown by a gradual change in the concentration of substances.