Chemistry: Acidic:2

2. While we usually 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 which act as acids and describe the conditions under which they act as acids
Acidic solutions are produced by many gaseous non-metal oxides when they dissolve in and react in water. These are called acidic oxides, eg. carbon dioxide, sulfur dioxide and nitrogen dioxide. They are the oxides of the non-metal elements except for CO, NO and N2O which are neutral).
To detect that a non-metal oxide gas is acidic with indicator paper, the paper must be moist. Moisture enables the gas to dissolve and form the aid that produces hydrogen ions. Reaction of a hydrogen ion with and indicator causes the colour change.

Analyse the position of these non-metals in the Periodic table and outline the relationship between position of the elements in the Periodic table and acidity/basicity of oxides
The oxides of the elements of the third period of the periodic table range from basic to acidic in character from left to right along the period. That is, as the element becomes more and more non-metallic in character, its oxide becomes more and more acidic. This is because that metal oxides are mostly basic. The more metallic an element is, the more basic its oxide. The five elements close to the borderline between metals and non-metals (Be, Al, An, Sn, Pb) are amphoteric, that is they show both acidic and basic properties. CO and NO are neutral oxides.
All the metals in Group 1 form basic oxides, however the strength of these oxides as bases are different. Going down the group, the strength increases. The stronger the metal oxide as a base, the more energy is released when it reacts with a given acid. Metal oxides tend to be basic, non-metals acidic and those in the middle tend to form amphoteric oxides. Therefore, the acidity of an oxide tends to increase across the Periodic table (l to r) and increases down the table as well. Group 8 does not form oxides.
Group 1
Group 2
Group 3
Group 4
Group 5
Group 6
Group 6
Li
Basic
Be
Amphoteric
B
Acidic
C
Acidic
N
Acidic
O
Neutral
F
Acidic
Na
Basic
Mg
Basic
Al
Amphoteric
Si
Acidic
P
Acidic
S
Acidic
Cl
Acidic
K
Basic
Ca
Basic
Ga
Amphoteric
Ge
Acidic/Amphoteric
As
Acidic/Amphoteric
Se
Acidic
Br
Acidic
I
Basic
Sr
Basic
In
Basic
Sn
Amphoteric
Sb
Acidic/Amphoteric
Te
Acidic/Amphoteric

Cs
Basic
Ba
Basic
Tl
Basic
Pb
Amphoteric
Bi
Acidic/Basic
Po
Acidic/Amphoteric


Define Le Chatelier’s principle
Le Chatelier’s principle qualitatively predicts the effects of disturbances in equilibrium reactions in general: If a change is imposed in a system at equilibrium, the position of equilibrium will shift in a way (direction) that tends to reduce/minimize that change.

Identify factors which can affect the equilibrium in a reversible reaction
- concentration: increased concentration = increased number of collisions
- temperature: increased temperature = shifts endothermic equilibrium to the products side and an exothermic equilibrium to the reactants side
- gas pressure: is a measure of the effective concentration of the particles. Increased pressure = increased concentration

Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle
In a closed system (such as the bloodstream or a bottle), if the concentration of CO2 is increased, then the CO2 goes into solution as H2CO3. If total pressure is increased, some CO2 dissolves to counteract the increased pressure. Because the forward reaction is exothermic (creation of bonds) and the reverse reaction is endothermic (bond breaking), if the temperature increases, the concentration of CO2 increases.
The solubility of carbon dioxide gas in water can be fully described using four equilibrium equations:
CO2(g) CO2(aq) + heat (ΔH=-ve)
H2O(l) + CO2(aq) H2CO3(aq)
H2CO3(aq) H+(aq) + HCO3-(aq)
HCO3-(aq) H+(aq) + CO32-(aq) This reaction does not occur a lot though, as HCO3 is amphoteric.
An equilibrium shift to the left releases carbon dioxide gas. An equilibrium shift to the right dissolves carbon dioxide gas. Le Chatelier's principle predicts that:
- addition of acid (increased concentration of H+) shifts equilibrium to the left
- addition of base (reacts with and reduces concentration of H+) shifts equilibrium to the right
- addition of a soluble carbonate (increased concentration of CO32-) shifts equilibrium to the left.

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

Natural
Industrial
SO2
- Geothermal hot springs and volcanoes
- Decomposition and combustion of organic (bush fires)
- Smelting of sulfide ores, eg. Cu, Zn, Pb, Ni
- Processing/ burning fossil fuels (coal contains 0.5% sulfur) but Aust. has low sulfur coal so it is high demand
N2O
- Action of certain bacteria on nitrogenous material in soils
- Increased use of nitrogenous fertilizers
NO
- Lightning (at high localized temperature generated by lightning, oxygen and nitrogen combine)
- Combustion in power stations and cars (at high temperatures in combustion chambers oxygen and nitrogen from air combine)
NO2

- Combustion (as above)

Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen
In a metal sulfide smelter, the ore is heated in air and converts to a metal oxide, releasing sulfur dioxide.
2ZnS(s) + 3O2(g) 2ZnO(s) + 2SO2(g)
Cu2S(l) + O2(g) 2Cu(l) + SO2(g)
When coal, petroleum or natural gas are burnt, sulfur in sulfur compounds is converted to sulfur dioxide.
S + O2 SO2
4FeS2(s) + 11O2(g) 2Fe2O3(s) + 8SO2(g)
Lightning strikes cause reaction between the two most common gases in the atmosphere.
N2(g) + O2(g) 2NO(g)
High temperature combustion reactions in furnaces and internal combustion engines produce significant amounts of NO [called nitrogen monoxide, nitric oxide or nitrogen(II) oxide] above 1300oC.
N2 + O2 2NO
Colourless, neutral nitrogen monoxide reacts with oxygen in the air to form brown, acidic nitrogen dioxide [nitrogen(IV) oxide].
2NO(g) + O2(g) 2NO2(g)

Assess the evidence which indicates increases in atmospheric concentration of sulfur and nitrogen
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.
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.
- The chemical instruments able to measure very low concentrations, like those for SO2, have only been commercially available since the 1970s.
- 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.
Evidence:
- Increase in respiratory problems could be linked to increased air pollution
- Altered fish breeding, because of change in pH, could be attributed to acid rain
- Yellowing of leaves and leaf drop, when SO2 enters stomata of leaves
- Increase in number of forests affected by acid rain (also from comparison of ice core samples)
- Several smog incidents in heavily populated cities, related to the pollution of air
- Increase in number of buildings experiencing damage from acid rain

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°K and 100kPa or 25°K and 100kPa
The equation linking the mass of a substance and its volume is given by:
n = V , where V = the volume of a gas at a specified temperature and pressure,
Vm Vm = the molar volume at the same conditions of temperature and pressure,
n = the number of moles in this volume
Temperature
Pressure
(kPa)
Molar volume
(L mol-1)
0 K or 273°C
100
22.71
25 K or 298°C
100
24.79
The mass (of a gas) is readily converted to moles. Given the conditions, the volumes of gas can then be determined.

Explain the formation and effects of acid rain
Acid rain is rain that has a hydrogen ion concentration higher than about 10-5mol/L (pH lower than 5). Distilled water in the atmosphere has a pH of 5.5 to 6, due to the absorption of CO2 from the atmosphere. If the pH is below 5, an acidic substance, such as SO2 or NO2, has dissolved in the water.
2SO2(g) + O2(g) 2SO3(g)
Both sulfur dioxide and sulfur trioxide react with water when it rains to produce acids.
SO2(g) + H2O(l) H2SO3(aq)
SO3(g) + H2O(l) H2SO4(aq) or
SO3(g) + H2O(l) H+(aq) + HSO4-(aq)
Nitrogen dioxide can react with oxygen to form nitrous acid HNO2 (see above), and nitric acid HNO3.
2NO2(g) + H2O(l) HNO2(aq) + HNO3(aq) or
2NO2(g) + H2O(l) HNO2(aq) +H+(aq) + NO-3(aq)
Effects:
- Soil pH can drop, making it difficult for plants to absorb sufficient calcium or potassium. A decrease in soil fertility will lead to crop failure and deaths of forests
- Soil chemistry can change, leading to the death of important micro-organisms and release of normally insoluble aluminium and mercury into soil water
- Protective waxes can be lost from leaves, causing leaf damage
- Buildings made of carbonates, such as concrete, mortar, limestone and marble can be gradually dissolved away
- Aquatic animals can die as pH drops below 5 in rivers and lakes
- Smog and acid rain can combine to form killer fog, as it happened in 1952 in London, when many homes burnt sulfur dioxide-releasing coal. Over 4000 lives were lost
- Also, the cost of treating drinking water may affect society

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 25K and 100kPa

1. Carefully weigh a sample of soda water (the sample should be between 100g and 200g).
2. Gently warm the soda water for 10 minutes. Do not boil the soda water.
3. Reweigh the sample of soda water.
Or:
1. Carefully weigh a sample of soda water (the sample should be between 100g and 200g).
2. To the sample of soda water add 15g of salt. Add the salt slowly to ensure that the soda water does not degas too quickly. (If the soda water foams over the top of the container and water is lost, the experiment will have to be repeated).
3. Allow the soda water to degas for 5 to 10 minutes
4. Reweigh the sample of soda water
Soda water and other effervescent soft drinks contain carbon dioxide gas under pressure. As pressurised carbon dioxide is pumped into the water in the factory the equilibria in equations (1) and (2) below shift to the right:
(1) CO2 (g) CO2 (aq)
(2) CO2 (aq) + H2O(l) H+(aq) + HCO3-(aq)
Hydrogen ions (H+) are formed which give the soft drink its sharp, acidic taste (pH 3.8). The equilibrium is maintained as long as the bottle is sealed. When the cap is removed, effervescence is observed. This can be explained by the loss of carbon dioxide in an open system. Both equilibria shift to the left as carbon dioxide gas is lost to the surroundings. The soft drink loses its acidity and goes flat.
Results:
Mass of Soda Water: 505.29g (incl. bottle)
Mass of water after degassing: 504.79g
Hence mass of CO2 released: 0.5g
Calculations:
Number of moles: 0.5/44.01= 0.01136
NB: 1 mole of gas occupies 24.79 L at 25oC and 100 kPa
0.01136 = Volume/24.79, therefore Volume of CO2 = 0.282 L
Average mass = 1.3875g which is equivalent to 0.78 L
Table of results:
With Salt
Heating
Mass lost (g)
Volume CO2
Mass lost (g)
Volume of CO2
1.04
0.589 L
0.32
0.18 L
7.26
4.09 L
0.18
0.10 L
0.42
0.24 L
0.5
0.282 L
1.35
0.74 L
0.03
0.282 L
(thanks divz for all the stuff and results)


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
Industrial origins:
- great increase in SO2 after the Industrial Revolution in 1800’s mainly from burning coal and extracting metals
- serious pollution from oxides of nitrogen developed in the C20th as electricity and use of motor cars expanded. High temperature combustion releases NO and NO2
Concerns:
- will dissolve in water droplets to form acid and acid rain, (and all the problems that brings) see above
- SO2 irritates the respiratory system and causes breathing difficulties at concentrations 1ppm (asthma and emphysema sufferers particularly susceptible), in humans
- The damaging effects of sulfur in animals is mostly from brain damage, as it causes the hypothalamus to malfunction and damages the nervous system.
- Plants are very sensitive to SO2. Low concentrations retard the production of chlorophyll , and high concentrations result in te formation of sulfuric acid and plant death
- Effects are magnified if particles are present
- NO2 irritates respiratory tract and causes breathing discomfort at concentrations above 3-5 ppm, at higher concentrations extensive tissue damage can be caused
- Leads to the formation of ozone, ie. Photochemical smog at concentrations of 0.1ppm