Chemistry Notes

HSC Core Topic 3

Chemical Monitoring and Management

  1. Much of the work of chemists involves monitoring the reactants and products of reactions and managing reaction conditions
  • Outline the role of a chemist in a named industry or enterprise, identifying the branch of chemistry undertaken by the chemist and explaining a chemical principle that the chemist uses

Case study example (from CCHSC)

  • A plant chemist in chemical manufacturing company that makes ethylene (ethene) from ethane, and then polymerises it or sells it for further processing
  • Branch of chemistry: analytical chemistry
  • Monitors:
  • Output (ethylene, propylene) quality (impurities)
  • Waste water for environmental standards (pH, suspensions, sulfate, grease, hydrocarbons)
  • Working with cracking (ethane -> ethene) to adjust operating conditions for optimal product yields
  • Correct operation of equipment (periodic recalibration)
  • Mostly uses gas chromatography:
  • Mixture vaporised into stream of helium flowing through tube over stationary phase (solid or liquid)
  • Gas-solid chromatography: different components of the mixture accumulate on the solid, and thus pass through the column at different rates
  • Gas-liquid chromatography: components dissolve into liquid at different extents. The more soluble a substance is, the slower it moves through the column.
  • A device at the end of the column detects each substance as it exits and thus can measure its concentration quantitatively when compared to standards
  • Describe an example of a chemical reaction such as combustion, where reactants form different products under different conditions and thus would need monitoring

Reaction condition effects

Incomplete combustion occurs if there is not enough oxygen gas available for reaction. Rather than producing CO2, it produces poisonous CO and solid C (soot). It also does not release as much energy as wanted from the complete combustion reaction. Therefore it is necessary for the reaction to be monitored in order to ensure that there is an excess of oxygen so that complete combustion occurs:
Complete combustion:C8H16(l) + 14O2(g) --> 8H2O(l) + 8CO2(g)
Incomplete combustion:C8H16(l) + 6O2(g) --> 8H2O(l) + 4CO(g) + 4C(s)

Another example is the reaction of ethene with oxygen. In a plentiful supply of oxygen at a high temperature, ethene reacts to form carbon dioxide (normal combustion).

C2H4(g) + 3O2(g) - 2CO2(g) + 2H2O(g)

However at lower temperatures and with lesser amounts of oxygen and with suitable catalysts, quite different products are formed. It is important to monitor reaction conditions to ensure that the yield of the desired product is maximised.

  • Gather, process and present information from secondary sources about the work of practising scientists identifying: the variety of chemical occupations, and a specific chemical occupation for more detailed study

Environmental chemist: are employed by the Environmental Protection Authority (EPA) as well as by mining companies, industries and local government to collect, analyse, and assess environmental data. Water resource authorities and air quality management authorities employ environmental chemists who monitor water and air samples for pollutants.

Metallurgical chemist: are scientists who specialise in the properties, applications and development of metals and alloys in out technological society, They give advice on the extraction of metals from ores and ways in which they could be combined with other materials such as polymers or ceramics.

Biochemist: studies the chemical structure and functions of molecules (eg. carbohydrates, fats, proteins, nucleic acids) in living things. Through their research new medical, industrial and agricultural products are developed.

Polymer chemist: investigate the properties of large polymeric molecules. They manipulate their structure to alter their properties in order to produce new and useful plastic products and materials.

Industrial chemist: study the structure and chemical reactions of materials that can be used in industry. Their research and development programs lead to the production of a wide variety of commercial products ranging from petrochemicals, detergents, and plastics to semiconductors. Industrial chemists may be specialists in analytical chemistry, organic chemistry or inorganic chemistry. An industrial chemist works in a team. Research ideas are developed into experimental procedures on a small scale. Consideration in normally given to factors such as reaction speed, use of catalysts and the position of the reaction equilibrium. With the aid of chemical engineers, the procedures are scaled up to the industrial production levels. Chemical monitoring procedures, quality control testing and environmental monitoring procedures are designed. An industrial chemist works closely with environmental officers and marketing and management personnel. Planning and organisational skills as well as analytical and laboratory skills are vital for in industrial chemist.

  1. Chemical processes in industry require monitoring and management to maximise production and ensure quality control
  • Identify and describe the industrial uses of ammonia

Ammonia ranks second to sulfuric acid in terms of quantity provided worldwide per year. It is used to make:

  • Fertilisers (sulfate of ammonia, ammonium nitrate, urea)
  • Fibres and plastics (rayon, acrylics, nylon)
  • Nitric acid, which in turn is used to make fertiliser (ammonium nitrate), dyes, fibres and plastics, and explosives such as ammonium nitrate, TNT (trinitrotoluene) and nitro-glycerine (in dynamite)
  • Household cleaners
  • Detergents (non-ionic ones).
  • Gather and process information from secondary sources to describe the conditions under which Haber developed the industrial synthesis of ammonia and evaluate the significance at that time in world history

By the beginning of the twentieth century there was a growing need for an industrial method of synthesising ammonia. Increasing demands for nitrogenous fertiliser to grow food for increasing world populations were placing strains on the supply of naturally-occurring Chile saltpetre (sodium nitrate), the main ‘artificial’ fertiliser at that time. In addition the growing militancy of Germany was promoting calls for more explosives (generally made from nitric acid which in turn was mainly produced from saltpetre).

The German, Fritz Haber, in 1908 first developed a method of synthesising ammonia from its elements, through it was not until 1914 that Carl Bosch successfully converted it into an industrial process. This synthesis contributed significantly to the German war effort in World War I in that it insulated German agriculture from any harm caused by interrupted supplies of saltpetre from South America and it facilitated the production of nitric acid and hence of explosives.

  • Identify that ammonia can be synthesised from its component gases, nitrogen and hydrogen
  • Describe the synthesis of ammonia as a reversible reaction that will reach equilibrium
  • Identify the reaction of hydrogen with nitrogen as exothermic
  • Explain why the rate of reaction is increased by higher temperatures
  • Explain why the yield of product in the Haber process is reduced at higher temperatures using Le Chatelier’s principle
  • Explain why the Haber process is based on a delicate balancing act involving reaction energy, reaction rate and equilibrium
  • Explain that the use of a catalyst will lower the reaction temperature required and identify the catalyst(s) used in the Haber process
  • Analyse the impact of increased pressure on the system involved in the Haber process

The synthesis of ammonia uses the simple exothermic reaction:

N2(g) + 3H2(g) -> 2NH3(g)

H = -92 kJ/mol

This is an equilibrium reaction which at ordinary pressures and temperatures lies well to the left.

Equilibrium Considerations

Le Chatelier’s principle shows us how to maximise the conversion of nitrogen and hydrogen to ammonia.

  1. If the pressure on a reaction system is increased, the equilibrium moves in the direction which tends to reduce pressure; that is, the direction which corresponds to a decrease in the number of moles of gas, since a decrease in the number of moles of gas in a container of fixed volume leads to a decrease in pressure.

If the pressure on equilibrium mixture of N2 and H2 is increased, the reaction moves to the right: some N2 and H2 react to form NH3. This happens because that direction corresponds to a decrease in the number of moles of gas:

4 moles of gas - 2 moles of gas

  1. If the temperature is lowered the equilibrium will move in the direction which tends to increase temperature (release heat). It is exothermic, so if temperature is lowered, it will move towards the right (form more ammonia).

The percentage conversion of nitrogen to ammonia varies with pressure (at constant temperature) and temperature (at constant pressure). On equilibrium considerations alone the reaction should be conducted at high pressure and low temperature.

Rate Considerations

However another consideration is how long will it take for the reaction to reach equilibrium: that is we need to consider the rate of the reaction. As for most reactions, the rate of the reaction decreases as temperature decreases. If we lower the temperature in order to move the equilibrium towards the right, we make the reaction very slow and so it takes a very long time to reach equilibrium.

One way to increase the rate of reaction is to find a suitable catalyst. Iron is a good catalyst. While this catalyst does speed up the reaction, the rate is still too slow at room temperature to be practical. Remember that while a catalyst speeds up a reaction, it does not affect the position of equilibrium. This is because it speeds up both the forward and reverse reactions.

Hence we have these situations: a low temperature produces a high yield (say 90% conversion of hydrogen to ammonia), but a very long time (weeks to months) is required to reach equilibrium, even with a catalyst. A high temperature causes equilibrium to be reach more quickly (in a few minutes) but the equilibrium yield is extremely low (say 0.1%).

Compromise

Compromise conditions are therefore used. A moderate temperature produces a moderate yield moderately quickly. Typical conditions for the industrial process, called the Haber process are:

  • A temperature of about 700 K (or about 400C) and
  • A total pressure of about 250 atmospheres.

With a reactant mixture having H2 and N2 in the ratio of 3 : 1, these conditions give an equilibrium conversion to ammonia of about 40%. The catalyst is magnetite, Fe3O4, with its surface layer reduced to free iron.

(Re: pic CCHC 193). Reactants pass through the catalyst reactor, then the mixture is cooled to condense out the ammonia formed: this can be drained off as required. Unreacted gases are fed back into the catalyst chamber along with incoming fresh reactants. None of the reactant is wasted. We essentially drive the reaction to a completion by condensing out the reaction product.

A stoichiometric mixture of hydrogen and nitrogen is used because, as ammonia is formed and condensed out, left-over reactants can be recycled through the process (with some fresh reactant mixture added) without any build-up of one reactant over the other.

An important factor in designing an industrial process is energy management. In the Haber process we would like to use the head released to heat up, at least partially, the incoming reactants and so minimise our energy costs: this has the added advantage that it stops the catalyst overheating and so losing activity. This is the reason for the incoming reactants flowing over the outside of the catalyst chamber entering it.

The Source of Reactants

Nitrogen can be obtained from the atmosphere, so hydrogen is the ‘difficult’ or expensive reactant to obtain.

In the laboratory we make hydrogen gas by reacting zinc of magnesium with hydrochloric acid. This process is too expensive for industrial use. Alternatively hydrogen can be made by the electrolysis of water, but except in special situations this is also very expensive. Industrially, hydrogen is generally produced by reacting methane (natural gas) or some other hydrocarbon with steam in the presence of a nickel catalyst at the temperature of about 750C:

CH4(g) + H2O(g)-> CO(g) + 3H2(g)

H = +206 k/J mol

Carbon monoxide poisons the iron catalyst in the Haber process and so must be removed. This is done with other catalytic reaction:

CO(g) + H2O(g)-> CO2(g) + H2(g)

Which has the added advantage of producing more hydrogen. The catalyst used is either Fe3O4 at 500C or Cu at 250C.

A complication for the Haber process is that we want a mixture of hydrogen and nitrogen that does not contain any oxygen (which can react explosively with hydrogen under the conditions used). Methane us also used to remove oxygen from air (in effect just normal combustion to CO2 and steam). By adjusting the quantities of methane, steam and air used, this combination of reactions can be made to product a 3 : 1 mixture of hydrogen and nitrogen. The only unwanted gas in the mixture is carbon dioxide. This is removed by reaction with a base, if oxygen has to be excluded and if the reactant mixture has to be recycled through the reactor, the use of a 3 : 1 mixture is the most efficient way of making ammonia.

Re. picture CCHSC p195

  • Explain why monitoring of the reaction vessel used in the Haber process is crucial and discuss the monitoring required

Because many different conditions must be maintained for efficient and safe operation of the Haber process, monitoring is essential. First temperature and total pressure must be monitored to keep them in the range for optimum conversion of reactants to products: in addition excessive temperature can damage the catalyst. Then it is essential to monitor the composition of the incoming gas stream: we need to ensure that the ratio of H2 to N2 is kept at 3 : 1 (to avoid a build-up of one reactant), that oxygen is absent (to avoid risk of explosion) and that concentrations of carbon monoxide and sulfur-containing species are sufficiently low to prevent poisoning of catalyst. With a well-maintained plant the catalyst can last up to eight years. Any build-up of unreactive gases such as argon and methane needs to be watched also because that can lower the efficiency of the conversion.

In Summary:

Ammonia can be synthesised from nitrogen and hydrogen (equilibrium).

N2(g) + 3H2(g) <--> 2NH3(g) ΔH = –92 kJ/mol

The rate of reaction would be increased by higher temperature, since particles move faster, causing more collisions and reactions.

But since the reaction is exothermic, an increase in temperature will favour the back reaction and thus reduce yield of NH3.

Since some changes that would usually be made to the set-up to increase reaction rate and decrease activation energy shift equilibrium left, there needs to be a delicate balancing of these factors to compromise between yield and reaction rate.

A catalyst in the reaction lowers activation energy and thus less heat is required for the reaction to occur. Iron (as surface on magnetite, Fe3O4) is the heterogenous catalyst used (N2 & H2 break apart and react on its surface).

Increased pressure shifts equilibrium right, in order to gain fewer overall moles of gas, therefore increasing yield in the Haber process.

The Haber process is thus performed with:

  • approximate temperature 700K, pressure 250atm
  • catalyst Fe3O4 with Fe surface
  • the product constantly being condensed and removed (shifting equilibrium right)
  • left-over reactants are recycled
  • incoming reactants heated up by previous reactions by flowing over reaction chamber (it also stops overheating the catalyst)

Monitoring is required:

  • Temperature and pressure monitored to keep in optimum range (and not to damage the catalyst)
  • H2:N2 :: 3:1 (to avoid build up of one reactant)
  • Oxygen not present (explosion)
  • Other contaminants not present (CO or species with S can poison catalyst)
  • Avoid build-up of unreactive gases (reduces efficiency of reaction)
  • Gather and process information from secondary sources to perform calculations to demonstrate the effect of volume, temperature and concentration changes on product formation in the Haber process

Quantitative analysis of equilibrium:

for reaction aA + bB <--> cC + dD,
This is known as the equilibrium constant, for a particular reaction at particular temperature. When concentrations of involved species change, the other concentrations adjust in order for K to remain constant.
A higher number means more to the right in equilibrium.

Hence subbing in numbers for volume, temperature and concentration on the Ammonia equilibrium will give us different K values.

  1. Manufactured products including food, drugs and household chemicals, are analysed to determine or ensure their chemical composition and/or energy content
  • Describe chemical tests to identify the following: anions: phosphate, sulfate, carbonate, chloride, cations: barium, calcium, lead, copper and iron
  • Perform a first-hand investigation to carry out a range of tests, including flame tests, or use information from secondary sources to identify the following ions: phosphate, sulfate, carbonate, chloride, barium, calcium, lead, copper and iron

Tests for cations

Cation / Chemical tests / Flame test colour
Pb2+ / Cl– precipitate (white)
I– precipitate (yellow)
Ba2+ / SO42– precipitate (white) / Pale green
Ca2+ / SO42– precipitate (white)
F– precipitate (white) / Brick-red
Cu2+
Blue / OH– precipitate (blue)
Precipitate with NH3dissolves / Blue-green
Fe2+
Pale green / OH– precipitate
MnO4– decolourises
Fe3+
Yellow / OH– precipitate (brown)
SCN– deep red

Pb2+ and Ca2+ in small concentrations may not precipitate; if no precipitate is found add KI, to distinguish.

Tests for anions

Anion / Tests (performed with nitrates for multiple)
CO3– / PH test 8 – 11
Add H+ Bubbles form
SO42– / Ba2+ precipitates (in acid)
Cl– / Ag+ precipitate
PO43– / Add NH3 and Ba2+ precipitates
  • Deduce the ions present in a sample from the results of these tests

From the above tests, we can deduce the ions present in a sample.

For multiple ions, tests need to be performed in an order, making sure that salts that will not react are used (ie. NO3– tests with anions). When precipitates form, the reaction needs to be driven to completion and then the precipitate filtered (or preferably centrifuged off) and the filtrate tested.

  • Describe the use of atomic absorption spectroscopy (AAS), in detecting concentrations of metal ions in solutions and assess its impact on scientific understanding of the effects of trace elements

Atomic absorption spectroscopy(AAS) is a method to determine concentrations of cations to very fine accuracy(below 1ppm). When light of an element’s emission spectrum is shone on it, it is absorbed. Each element has a unique set of wavelengths whose energy it will absorb.

The substance is placed in a flame, and lamps shine possible emission spectra onto the substance (now melted into atoms). The resulting intensity of light is measured and calibrated against the intensity without the substance (and with known substances), and thus the concentrations of the ions in the substance can be determined. The major disadvantage is the cost of equipment and need for different light sources for each spectrum tested. The electrons of the exposed atom absorb the relevant energy. Since nearly 100% of the atoms absorb, AAS is extremely sensitive. It is used to find metals in the environment (eg. Mercury in fish), micronutrients in soils, contaminants in foods, trace elements in organisms. By using AAS, it was discovered that organisms require some elements in very minute amounts (1-100ppm), and thus allowed research into their essentiality & effects within organisms.