Acidic Oxides and the Atmosphere

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

Most oxides of non-metals, excluding the neutral oxides of carbon monoxide, nitrogen monoxide, and nitrous oxide (CO, NO and N_2O respectively), are acidic when in solution. Examples of such acidic oxides include carbon dioxide (CO_2) and nitrogen dioxide (NO_2).

Remember- Non-metal oxides are acidic in solution.

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 easiest way to remember the information below is to think of substances you know that are basic or acidic. We know that nitrogen dioxide is a component of acid rain, so must be acidic. Nitrogen is a non-metal, so generally non-metal oxides are acidic. We also know that lime, calcium oxide, is basic. Calcium is a metal, so generally metal oxides are basic. This will save you learning by rote.

Metal oxides on the left side of the Periodic Table are generally basic. Non-metal oxides on the right side of the Periodic Table are generally acidic. Oxides of the five lower metalloids, those elements bordering metals and non-metals, can be amphoteric depending on their oxidation states. An amphoteric substance shows both acidic and basic properties.

Remember- The left side of the periodic table generally consists of acidic oxides, while the right side is generally made of basic oxides. Thus metals form acidic oxides, and non-metals form basic oxides.

Define Le Chatelier’s principle

The symbol for equilibrium is ↽⇀ and simply means that, in a closed system, the rate of the forwards reaction is equal to the backwards reaction. This simple means that the reactants are

converting to products at the same rate that the products are converting back into the reactants. Whilst there appears to be no change on a macroscopic level, the system is continually changing on a microscopic level. This process, known as dynamic equilibrium, results in the concentration of the substances in the system remaining constant.

According to Le Chatelier’s principle, if a system at equilibrium is disturbed, then the system will adjust itself in order to minimise the disturbance. However, note that the effects of the disturbance are never fully removed. They are only minimised, or lessened to a degree.

Identify factors which can affect the equilibrium in a reversible reaction

Le Chatelier’s principle is one which plays a crucial role in the HSC Chemistry course. Thus, a sound understanding of it is important, and it may appear again in this subject depending upon what Option you do. For this reason, a treatment sounder than required for this dotpoint will be provided.

Several factors can affect the equilibrium in a reversible reaction. These disturbances to the system can be in the form of changes in concentration, pressure, volume, or temperature.


Imagine a system in equilibrium of four compounds, A, B, C, and D.

A + B ↽⇀ C + D

The simplest way of visualising changes in concentration is simply seeing Le Chatelier’s principle as working to minimise any changes made to the equilibrium. As more of A or B is added, then the system will try to minimise the change by converting more A and B into C and D. As such, the equilibrium shifts to the right.

Conversely, if more of C or D is added, increasing the concentration of the products, then the system will convert more C and D into A and B, shifting the equilibrium to the left.

Note that a system can only minimise a disturbance. It cannot completely undo it.


Imagine a system in equilibrium of four compounds, A, B, C, and D. Unlike the example used to illustrate changes in concentration, the four compounds in this example are gases, and the number of moles of A is two rather than one.

2A_{(g)} + B_{(g)} ↽⇀ C_{(g)} + D_{(g)}

Determining the affect of changes in the pressure of a system is simply an exercise in counting moles of gases. In the equilibrium above, there are three moles of gas on the left side, and 2 moles of gas on the right. Any increase in pressure will result in the system trying to relieve the pressure by ‘leveling’ the moles of gas within the system. As such, in the above system, an increase in pressure will lead to a shift in the equilibrium to the right. This occurs simply because the system is essentially counteracting the fact that three moles of gas are becoming two moles of gas.

Conversely, a decrease in pressure will shift the above equilibrium to the left in an attempt to increase pressure once again.

Changes in pressure affect only gases. Increasing the pressure in the following system will lead to equilibrium shifting to the right, as there are two moles of gas on the left side and only one on the right.

A_{(g)} + B_{(g)} ↽⇀ C_{(g)} + D_{(s)}


Any change in volume in a gaseous equilibrium is simply a change in pressure. As such, treat increases in volume as decreases in pressure, as there are more moles of gas in the fixed space, and treat decreases in volume as increases in pressure.


The effect of Le Chatelier’s principle with changes in temperature can often be confusing. However, simply thinking of heat as either a product or reactant greatly simplifies any problems, as shown in the equilibrium below, where the reaction is endothermic (Absorbs heat in order for the reaction to occur) rather than exothermic (Releases heat).

A + B + Heat ↽⇀ C + D

In the above endothermic equilibrium, an increase in temperature will result in the system working to reduce the temperature by shifting the equilibrium to the right, converting A and B into C and D in order to reduce temperature.

Conversely, a decrease in temperature will shift the equilibrium to the left, converting C and D into A and B in order to produce more heat.

In the case of an exothermic reaction, the equation will be of the form

A + B ↽⇀ C + D + Heat

As shown above, treating heat energy as an actual item in the equilibrium is a much simpler method of thinking of a problem. Simply determine whether a reaction is exothermic forwards, i.e. the heat is placed on the right, or endothermic forwards, i.e. the heat is placed on the left.

Remember- Changes in concentration, pressure, volume and temperature will all disturb a system in equilibrium.

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

This dotpoint is really just an application of Le Chatelier’s principle using the solubility of carbon dioxide as an example. As such, set yourself in the habit for such questions by starting with the equation, and then working through changes in concentration, pressure, volume and temperature.

Note that when you have soft drink in a glass or open bottle, you can see bubbles rising in it. This is because the carbon dioxide gas is constantly escaping, thereby constantly favouring the backwards reaction in an attempt to minimise the disturbance to the system. In comparison, a closed bottle of soft drink has no bubbles unless you shake it, because it is in equilibrium.

CO_{2(g)} + H_2 O_{(l)} ↽⇀ H_2CO_{3(aq)} + Heat

Using the above equilibrium as a practical example of Le Chatelier’s principle:

  • An increase in the concentration of CO_{2(g)} will shift the equilibrium to the right, converting carbon dioxide and water into carbonic acid in order to reduce the concentration of carbon dioxide.
  • An increase in pressure will shift the equilibrium to the right, converting carbon dioxide and water into carbonic acid in order to reduce the
  • An increase in the volume of CO_{2(g)} will shift the equilibrium to the right, converting carbon dioxide and water into carbonic acid in order to reduce the volume of carbon Thus

the system will attempt to counteract this change by favouring the backwards reaction.

  • An increase in temperature will shift the equilibrium to the left, converting carbonic acid into carbon dioxide and water in order to reduce the
  • An increase in temperature will shift the equilibrium to the left, converting carbonic acid into carbon dioxide and water in order to reduce the

Remember- Le Chatelier’s principle will ensure that equilibrium is reached once again. However, this new point of equilibrium will not be same as the original point of equilibrium, as the impact was only minimised, not completely reversed. This is the reason why opened soft drinks will go ‘flat’ irreversibly.


Calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0◦C and 100kPa or 25◦C and 100kPa

These calculations are made simpler by the fact the HSC data sheet will provide you with the relevant values. The only thing remaining for you to do is to calculate the number of moles and multiply this by the relevant constant. As a note, be wary of earlier papers which used different constants.

According to the hypothesis developed by Avogadro, equal volumes of gases contain equal moles when at the same temperature and pressure.

At 0◦C and 100kPA, 1 mole of gas is equal to 22.71L. At 25◦C and 100kPA, 1 mole of gas is equal to 24.79L.

Example: Excess hydrochloric acid is added to 5g Cu. What is the volume of gas emitted at 0◦C and 100kPA?

Step 1: Write out the equation. In this case:

Cu_{(s)} + 2 HCl_{(aq)}H_{2(g)} + CuCl_{2(s)}

Step 2: Find the required moles. 5g of copper = \frac {5}{63.55} moles = Number of moles of hydrogen gas produced

Step 3: Use the constants provided. At 0^oC and 100 kPA, 1 mole of gas equal to 22.71L. {5}{63.55} x 22.71 = 1.79L of H_2 gas emitted.

In making these calculations, always remember that n = \frac {g}{M} , (where n is equal to the number of moles, g is equal to weight, and M is equal to the molecular weight of respective element). This equation is crucial to virtually all stoichiometric calculations.

Identify data, plan and perform a first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 25◦C and 100kPa

If you are unsure how the volume of gas is calculated, review the method outlined in the previous dot point. Also don’t forget that the volume of gas released is given in litres, not milliliters.


  1. Record the weight of a bottle of soft
  1. Shake the bottle slightly, opening it carefully to release the gas without spilling the contents. Repeat this step until shaking ceases to have a noticeable
  2. Remove the cap, and immerse the bottle in a warm bath for several
  1. Remove the bottle from the bath and leave it overnight in a place away from sources of evaporation.
  2. Replace the cap and reweigh the bottle and record its
  1. Calculate the volume of carbon dioxide gas

Expected Results:

This experiment is simply a practical application of the previous dot point. The formula n = \frac {g}{M} will thus be used. Divide the change in weight (in grams) by the molecular mass of carbon dioxide (M= 44.01), and multiply by the constant (At 25^oC and 100kPA, the relevant constant is 24.79).

For example, if the change in mass was 10g, the number of moles was thus 0.23, and the volume of carbon dioxide gas released must have been 5.63L.

You must be able to work backwards with this formula as well as forwards. Should you be given the volume of gas released, you should know the change in gas.


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

You will notice this dotpoint turning towards acid rain by first introducing the oxides from which it is formed. Be sure to include all the oxides of nitrogen. These include: Nitrogen monoxide (Also known as nitric oxide), nitrogen dioxide, and nitrous oxide. If you have trouble remembering these, focus on the ’monoxide’ and ’dioxide’, as these make it clear there are one and two oxygen atoms in the respective molecules.

Natural Sources

Two-thirds of all sulfur dioxide (SO_2) are produced naturally by geothermal hot springs and volcanoes.

Nitrogen monoxide (NO) is produced by lightning, as the high localised temperatures created by lightning are sufficient to convert oxygen and nitrogen in the atmosphere to form nitrogen monoxide. Nitrogen monoxide then reacts slowly with oxygen to form nitrogen dioxide (NO_2):

2 NO_{(g)} + O_{2 (g)}2 NO_{2 (g)}

Nitrous oxide (N_2O) is created naturally by certain bacteria in nitrogen-rich soils.

Industrial Sources

Industrial sources of sulfur dioxide include the combustion of fossil fuels and extraction and refinement of metals from sulfide ores, where SO_2 is often released during the smelting of the ores in order to remove sulfur impurities from the metal. This is because there are traces of sulfur within the ores

which are released upon extraction and smelting.

Combustion within power stations and automobiles is an industrial source of nitrogen monoxide and nitrogen dioxide, referred to as NO_x. Just as with lightning, the high temperatures allow the conversion of oxygen and nitrogen into nitrogen monoxide, which combines with oxygen to form nitrogen dioxide.

Remember- Sulfur dioxide occurs naturally due to volcanoes and industrially during the smelting of sulfur-contaminated metal ores. Nitrogen monoxide is formed from atmospheric reactions due to lightning, and nitrogen dioxide forms as nitrogen monoxide in turn reacts with the atmosphere. They are also formed industrially through combustion in automobiles. Nitrous oxide forms from bacteria.


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

This dotpoint naturally follows from the previous one, providing the equations for the processes described.

When various fossil fuels are combusted, sulfur within the compound can combine with oxygen to produce sulfur dioxide.

S_{(s)} + O_{2 (g)}SO_{2 (g)}

The smelting of iron pyrite produces sulfur dioxide.

4 FeS_{2 (s)} + 11 O_{2 (g)}2 Fe_2O_{3 (s)} + 8 SO_{2 (g)}

High localised temperatures produced in conditions such as lightning or combustion chambers result in the reaction of nitrogen and oxygen to form nitrogen monoxide.

N_{2 (g)} + O_{2 (g)}2 NO_{(g)}

Nitrogen monoxide slowly combines with oxygen to form nitrogen dioxide.

2 NO_{(g)} + O_{2 (g)}2 NO_{2 (g)}

Explain the formation and effects of acid rain

Approach a question like this by splitting it up into formation (atmospheric pollutants, solubility of the acidic oxides) and the effects (environmental and human). If you can mentally organise it this way, you should not only be able to structure a response, but recall the information quite easily as well.

As described before, both sulfur dioxide and nitrogen dioxide are soluble. When the concentrations

of the emissions of these gases reach high levels, then rain can become fairly acidic, depending on the amount of gas dissolved. When hydrogen ion concentrations reach around $latex 10^{−5} mol/L, i.e. pH of 5, the rain is described as acid rain.

Acid rain has had a large impact on the environment, increasing the acidity in some lakes to the point where marine life can no longer inhabit the waters for extended periods of time. In addition, forests on an international level have been ravaged by the effects of acid rain, as the surface of leaves upon which the leaves are dependant upon drawing in the water necessary for photosynthesis are destroyed.

The effects of acid rain are not limited to the environment, as man-made structures have been similarly affected. In particular, limestone and marble statues and buildings have had considerable damage dealt to them, as acid corrodes carbonates with considerable ease.

Remember- The reason acid rain is formed is because acidic oxides such as sulfur dioxide and nitrogen dioxide are soluble, reducing the pH of rain considerably at large concentrations. Acid rain impacts the environment and man-made structures alike.

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

Don’t be afraid to mention the fact that we have relatively little data to work with in calculating any changes in atmospheric concentration. Most information has only been gathered within the last few decades, with any further information being relatively inaccurate, only illustrating basic trends from sources such as trapped gases in the ice caps.

Being soluble in nature, SO_2 and NO_2 levels have not increased dramatically on a global scale. This is because the water cycle effectively ‘cleans’ the atmosphere of these oxides on a regular basis. Despite this, the average annual concentration of these two gases is still many times over the concentration of clean air, with numerous days appearing each year where emissions completely go past ‘safe levels’.

In contrast to the relatively stable levels of SO_2 and NO_2, N_2O emissions in Australia have been reported to rise 130% since the 1990s. This has largely been attributed to the use of fertilisers.

Despite such figures, any measurements obtained lack any real figures to compare with, as accurate methods of measurement have only been developed within the last few decades.

Remember- Nitrous oxide concentrations have increased significantly in recent decades, whereas sulfur dioxide and nitrogen dioxide levels have remained relatively constant.


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

As the industrial origins of sulfur dioxide and the various oxides of nitrogen have been covered in prior dotpoints, only the reasons for concern about their release will be covered in this dotpoint.

There exists concern about the release of sulfur dioxide due to health and environmental reasons. Sulfur dioxide aggravates existing lung conditions, and can trigger asthma attacks and bronchitis. In addition, the effects of acid rain previously mentioned are a real problem to marine life, forests, and man-made structures made of marble and limestone.

Similarly, nitrogen dioxide can be detrimental to asthmatics, and can irritate the airways. Even in moderate levels, long-term exposure can increase the chances of respiratory illness, and sensitise people to allergies. Nitrogen dioxide is an even more problematic source of acid rain, as the nitric acid formed is a stronger acid than sulfuric acid.

Photochemical smog produced by nitrogen oxides, sunlight and volatile organic compounds is another large problem, as it leaves ozone- a powerful lung irritant and dangerous in even small concentrations- at a level proximate to humans.

As for nitrous oxide, it has been estimated that N_2O has a global warming potential 300 times higher than carbon dioxide. When combined with the alarming growth reported before, of 130% since the 1990s, this has become a problem of great import.

Thus, the evidence is both plentiful and persuasive, making it abundantly clear that concerns regard- ing the release of sulfur dioxide and the oxides of nitrogen are in fact well-founded.

Remember- Sulfur dioxide is formed during the smelting of sulfur-contaminated metal ores, and nitrogen monoxide and nitrogen dioxide are formed industrially through combustion in automobiles. Environmental concerns include attacks on the respiratory system (lungs and airways) for people, acid rain, and global warming.