M8 ACI 3

IQ3 What are the implications for society of chemical synthesis and design?

8.3.1 evaluate the factors that need to be considered in designing a chemical synthesis process, including but not limited to:

a) availability of reagents

b) reaction conditions (ACSCH133)

c) yield and purity (ACSCH134)

d) industrial uses (eg pharmaceutical, cosmetics, cleaning products, fuels) (ACSCH131)

e) environmental, social and economic issues

In this final section of the course, there are a couple of important applications of Chemistry which are best looked at through a particular case study focusing on an Industrial process, eg the Haber Process. No specific processes are named, so you need to be able to discuss the key factors (listed below) in relation to a given process which you may be given during an exam.

In order to cover a few processes briefly, the next few pages will focus on a particular industrial process with one, or more, of the factors highlighted. You will need to add depth if you want to focus on one or more of these as part of your learning menu.

Chemical synthesis is the process of making a new chemical substance. We can synthesise ammonia from its constituent elements: hydrogen and nitrogen, using the Haber Process. In an industrial process, the twin focuses are scale and economy: to produce large quantities while keeping costs manageable. This may mean considering a balance between factors which impact on an equilibrium system, cost of materials, disposal of wastes and safety for workers.

For any industrial process, you should be able to discuss:

• Availability of reagents
• Reaction Conditions
• Yield and Purity
• Industrial Use
• Environmental Issues
• Social Issues
• Economic Issues

Investigate using two of the following options (brief outlines of a couple are given below):

• Solvay Process

• Haber Process

• Frasch/Contact Processes

• Saponification and production of synthetic detergents

• Ostwald Process

• Sodium Hydroxide Production

Create a table to compare the two processes in terms of:

• availability of reagents
• reaction conditions
• yield and purity
• industrial uses and
• environmental, social and economic issues

8.3.1 evaluate the factors that need to be considered when designing a chemical synthesis process, including but not limited to

a) availability of reagents

It is preferable for manufacturing plants to have easy access to the reagents used to produce the desired product. This can be an economic benefit, a time and availability benefit, or both. For such reasons, oil refineries are located very close to shore whereby the transportation distance of ships to transport oil extracted from machines in the middle of the ocean can be reduced. This saves both time and money for the oil supplying companies.

Another example is that coal-fired power stations are built near coal deposit sites so that coal that is obtained from the site can be quickly transported to the power stations. Again, this saves both time and money.

Case Study: The Solvay Process

Sodium carbonate is an industrially important compound. It is an alkaline substance like sodium hydroxide but it is cheaper to produce. The process used to produce sodium carbonate is known as the Solvay process. The raw materials used in the Solvay Process are brine (salt water) and limestone and the products are sodium carbonate and calcium chloride (waste product). Ammonia is a raw material used in the process, however it is recycled.

The Solvay Process involves four main steps:

1. Brine Purification

2. Sodium Hydrogen Carbonate Formation

3. Sodium Carbonate Formation

4. Ammonia Recovery

However, the overall process can be represented as follows:

2NaCl(aq) + CaCO3(s) → CaCl2(aq) + Na2CO3(s)

Brine purification refers to salt water, often sourced from the ocean. The other main raw material is calcium carbonate which is usually sourced from limestone. Ideally the location for an industrial Solvay plant should be close to a limestone quarry and/or brine supply, whether this is a solid mineral deposit or near the ocean. If either material is not close at hand, costs of transportation will significantly increase. It is also better, when there is a choice, to transport the solid limestone rather than the liquid brine.

2.1 evaluate the factors that need to be considered when designing a chemical synthesis process, including but not limited to

b) reaction conditions

Some revision of Module 5:

LE CHATELIER’S PRINCIPLE

• general rule that helps in predicting the direction in which an equilibrium reaction will move when a change in pressure, volume, concentration or temperature occurs
• “if an external stress is applied to a system in equilibrium, the system adjusts in such a way that the stress is partially offset as the system reaches a new equilibrium position”

PRESSURE

• changes in pressure do not affect the concentration of the reacting species in condensed phases (e.g. solids and liquids) because solids and liquids are incompressible (can't be compressed).
• concentrations of gases are greatly affected by changes in pressure
  • INCREASE in PRESSURE [is an INCREASE in CONCENTRATION of the gas species] favours the reaction that decreases the total number of moles of gases. In the example given, the increase in pressure favours the reverse reaction.
  • DECREASE in PRESSURE favours the reaction that increases the total number of moles of gases.

Suppose that the equilibrium system given below is in a cylinder fitted with a movable piston. The decrease in pressure favours the forward reaction.

CH4(g) + H2O(g) ⇌ CO2(g) + 3H2(g) (endothermic)

VOLUME

• Changing the volume and maintaining the same pressure does not alter the equilibrium established by the system.
• Decreasing the volume increases the pressure (and therefore concentration), while increasing the volume decreases the pressure (and therefore concentration).
• It is possible to change the pressure of the system without changing its volume. The total pressure of the system can be increased by adding an inert gas.
• Addition of inert gas at constant volume increases the total gas pressure and decreases the mole fractions of the gaseous species. However, the partial pressure of each gas does not change. As such, the presence of an inert gas does not affect the equilibrium condition.

CONCENTRATION

• Changing the concentration of either reactants or products can disturb the equilibrium condition.
• An INCREASE in the concentration of the reactants shifts the equilibrium to favor formation of products. On the other hand, INCREASING the concentration of the products shifts the equilibrium to the left.

Suppose that the equilibrium system given below is in a cylinder fitted with a movable piston. Increasing the amount of hydrogen gas favors equilibrium to form CH4 and H2O. Decreasing the concentration of CH4 favors equilibrium to the left, thereby producing more CH4 and H2O.

CH4(g) + H2O(g) ⇌ CO2(g) + 3H2(g) (endothermic)

TEMPERATURE

• Changes in concentration, pressure and volume may alter the equilibrium position but do not change the value of the equilibrium constant.
• However, change in temperature can alter the equilibrium constant (more on that later).

CH4(g) + H2O(g) ⇌ CO2(g) + 3H2(g) (endothermic)

• The forward reaction is ENDOTHERMIC (the system ABSORBS HEAT):

heat + CH4(g) + H2O(g) ⇌ CO2(g) + 3H2(g)

• The reverse reaction is EXOTHERMIC (the system RELEASES HEAT):

CO2(g) + 3H2(g) ⇌ CH4(g) + H2O(g) + heat

• At equilibrium at a certain temperature, the heat change is ZERO because there is no net reaction.
• An INCREASE in temperature favours ENDOTHERMIC reaction whereas a DECREASE in temperature favours the EXOTHERMIC reaction.

b) reaction conditions (contd.)

Case Study: The Haber Process

Ammonia is an important industrial chemical. It is used as a fertiliser, a component in many alkaline cleansers, eg window and floor cleaners, and as a refrigerant gas. Ammonia is used in the manufacture of

• fertilisers - ammonium sulfate, ammonium nitrate, ammonium hydrogen phosphate, and urea
• nitric acid - used in the manufacture of ammonium nitrate fertiliser, dyes, fibres, plastics and explosives, such as ammonium nitrate, trinitrotoluene (TNT) and nitroglycerine
• cyanides -used in the manufacture of synthetic polymers, such as nylon and acrylics, and to extract gold from ore bodies

Ammonia can be synthesised from its constituent elements, according to the following equation:

N2(g) + 3H2(g)  2NH3(g) ΔH = -92 kJ

The synthesis of ammonia from nitrogen gas and hydrogen gas does not go to completion. It is a reversible reaction and an equilibrium will be established where the rate of reactants forming product equals the rate of product decomposing into reactants.

This reaction is known as the Haber Process. As it is an equilibrium reaction, varying the reaction conditions can significantly affect the behaviour of the reactants and products of this reaction.

Effect of Temperature

Higher temperatures increase the kinetic energy of the reactant molecules. This results in a greater proportion of the molecules reaching energies greater than the activation energy, hence causing a greater number of molecules to react.

Thus, higher temperatures increase the rate of the reaction.

The synthesis of ammonia is an exothermic reaction, releasing heat as a product. This means the reverse reaction is endothermic.

Increasing temperature is equivalent to adding energy to the system. This is a change in the equilibrium conditions. According to Le Chatelier’s principle, the system will respond in such a way as to counteract the shift. In this case it will shift so as to use the extra energy causing the equilibrium to shift in the reverse direction, favouring the endothermic reaction. This opposes the change in energy by absorbing heat and lowering the temperature.

Higher temperatures in the reaction vessel result in a reduced yield of ammonia in the Haber process.

The Haber process is a commercial process. As a result there are costs associated with the manufacture of ammonia and hence it is important to have a high reaction rate and a high yield. Both of these are temperature dependent.

As temperature increases, the rate of reaction increases, however the yield decreases as a result of Le Chatelier’s Principle. At lower temperatures, rate of reaction decreases, however yield increases.

As a result, a moderate temperature is required which balances these competing demands of reaction rate and yield. This provides the maximum yield of ammonia as possible in the shortest time.

Temperatures for this reaction are typically in the range of 400°C to 500°C.

Effect of Pressure

Gas pressures are also important to a system in equilibrium. In the Haber process, four moles of reactant gases produce two moles of ammonia gas. As a result, the forward reaction results in a decrease in the pressure of the system.

An increase in pressure causes the equilibrium to shift in the forward direction, to counteract the change by reducing the pressure, again according to Le Chatelier’s principle.

Higher pressures result in an increased yield of ammonia in the Haber process.

Typical pressures used in the Haber process range from 15 to 35 MPa (15 000 to 35 000 kPa).

Effect of a Catalyst

Using a catalyst does not change the position of equilibrium, but it lowers the activation energy of the reactants. This means the reaction temperature does not need to be as high for the reaction to proceed.

One catalyst used in the Haber process is magnetite (you don't need to know this - Fe3O4: very complicated lattice Fe²⁺Fe³⁺₂O_4. The chemical IUPAC name is iron(II,III) oxide and the common chemical name is ferrous-ferric oxide.)

The catalyst is ground finely ensuring it has a large surface area (around 50 m2/g), and the magnetite is reduced to elemental iron, Fe. The large surface area allows gaseous molecules to rapidly absorb and react.

The use of the catalyst allows the reaction rate to remain high despite the reaction occurring at a moderate temperature.

All these factors make it critical to monitor the reaction vessel during the Haber Process to ensure there is a safe balance of reaction conditions to ensure maximum yield at minimal cost.

3.1 evaluate the factors that need to be considered when designing a chemical synthesis process, including but not limited to

c) yield and purity

Yield

Yield = the amount produced of an industrial product, usually expressed as a percentage of the theoretical amount that could be produced if 100% of reactants formed 100% of products, eg 93% yield means 7% of reactants did not react to form products.

There are two different terminologies for yield – the theoretical yield and the actual yield.

Theoretical yield is basically the moles of product (e.g. chemical) that is expected to form (using mole ratio from BUMS) if all reactants supplied were converted to products.

Actual yield is the moles of product that did form when the reaction was carried out.

We calculate the theoretical yield prior to the reaction, and we measure the actual yield after the product has been obtained.

When we compare the actual yield with the theoretical yield, we can use a percentage yield.

Percentage yield = (actual yield / theoretical yield) x 100 = ? %

Case Study: Esterification

The esterification reaction is reversible, reactants are in equilibrium with products.

RCOOH + R′OH ⇌ RCOOR′ + H₂O

When the reaction reaches equilibrium there are still large amounts of the reactants left in the mixture, resulting in a poor yield of the ester.

The yield of ester can be improved by increasing the concentration of one of the reactants (either the alcohol or the carboxylic acid).

By Le Chatelier's Principle an excess of one reactant will drive the reaction to the right, increasing the production of ester, and therefore increasing the yield of ester.

In our experiment we are using an excess of carboxylic aicd (acetic acid which is also known as ethanoic acid).

From the balanced chemical equation, you will also appreciate that the presence of water in the reaction mixture will drive the equilibrium to the left, favouring the formation of reactants rather than ester.

So, another way to improve yield is to remove the water as it forms.

Concentrated sulfuric acid can be used for this purpose because it reacts rapidly with water to form a hydrated form of sulfuric acid, effectively removing the water from the reaction mixture.

The esterification reaction is quite slow.

Heating the reaction mixture will speed up the rate of reaction.

The experimental technique used is known as "heating under reflux". As the volatile components in the reaction mixture vapourise, the hot vapours rise up into the condenser. Cool water running around inside the water-jacket around the condenser removes the heat from the hot vapours, cooling them down and allowing them to condense on the inside of the condenser. This cooled liquid runs back down the inside of the condenser and is returned to the reaction mixture.

Heating the reaction mixture under reflux prevents the loss of volatile (easily evaporated at normal temperatures) reactants and products.

Concentrated sulfuric acid is used as a catalyst to speed up the rate at which the ester is formed.

Purity

Why could there be a difference between the theoretical yield and actual yield - one factor is purity.

(We touched on the concept of purity when looking at the need to produce a primary standard solution for titration. If the titrant is not pure, the calculation of the concentration of the analyte will be inaccurate.)

Purity might be affected because the total mass of the reactant may consist of other atoms (impurity) that are not used to produce the product, but some other by-products. Therefore, our actual yield will be lower than the theoretical yield.

Another factor may be the reaction conditions: eg f the temperature of the reaction occurs at a lower temperature in reality than expected, the rate of reaction will be lowered and the actual yield of the product per hour may be lower than the theoretical yield.

It is important in many industries to measure the purity of their product. The importance of this is very broad, some reasons for the measurement of the chemical composition and purity of a chemical can be:

• Government regulations require accurate reporting of ingredients present inside a product (Nutrition and Ingredients Label) eg, people with an allergy to a specific ingredient can rely on such nutritional labelling before purchase.
• The purity of a drug is particularly important in the pharmaceutical industry as it involves interaction with living cells and can result in development or worsening of a disease.
• Marketing purposes (e.g. 100% pure honey) to deliver true and appealing information to capture consumers’ attention
• Sydney Water monitors and tests the purity of water to ensure that it meets the regulatory requirements for its water to be qualified as potable (drinkable) water when it reaches the consumers’ pipelines.
• Petrol composition and purity are tested to ensure that petrol quality meets standard requirements so that car engines are not damaged.
• The petroleum and natural gas obtained from petroleum extraction plants will be test for its chemical composition and purity as it will dictate the wholesale price.

Case Study: Esterification

The ester synthesised by refluxing is impure.

The round bottom flask contains a mixture of organic and inorganic substances:

• ester (desired product)
• alcohol, carboxylic acid (excess reactants now present as an impurity)
• water (undesired product of the reaction)
• sulfuric acid (used to catalyse the reaction)

These substances are all miscible with water, except the ester, which is the fraction needed to be isolated.

This isolation (separation) is achieved by taking advantage of the ester's low solubility in water, as well as its lower density, which means that it will float on top of a layer of water.

In the laboratory, a separating funnel is used to isolate the less dense ester layer from the aqueous layer. In industry, techniques such as liquid–liquid extraction and membrane separation are employed.

3.1 evaluate the factors that need to be considered when designing a chemical synthesis process, including but not limited to

d) industrial uses (eg pharmaceutical, cosmetics, cleaning products, fuels)

Case Study: Sulfuric acid

Sulfuric acid is one of the most important industrial chemicals, with many applications:

*Used by fertiliser industry (~70% total production) to make ammonium sulfate and superphosphate

*Removes oxides and grease from steel and iron before galvanising or electroplating

*Dehydration agent i.e. removes water liberated during manufacture of explosives, dyes,detergents, polymers, esters, electrolysis of sodium chloride solution

*Electrolyte in vehicle batteries

*In the production of nitroglycerine for explosives and as a vaso-dilator (treatment of heart conditions)

Case Study: Saponification

The main differences between the industrial and school processes are:

• Industrial processes occur at much higher temperatures
• Industrial processes often continue over days or weeks to reach completion
• Fats and oils are less volatile than esters and hence do not need to be refluxed
• Excess sodium hydroxide is neutralised and water distilled industrially, improving the purity of the final product
• Industrial ‘salting out’ process is difficult to replicate in a school laboratory

Many Uses

It is critical to plan and design the physical reaction pathway in which reactant and products are inserted and obtained in the manufacturing plant.

By doing so, any waste or by-products formed as a result of the reaction can be used as reactants or drivers for the same or a different reaction. This will lower the overall cost of production of the chemical.

Case Study: Steel Manufacture

The blast furnace gas (composed of mainly nitrogen and carbon dioxide) are produced as a result of smelting of iron ores. Iron is a component of steel. The heat derived from such the hot gas is used as source of energy to power up turbines and generate electricity used in the industry.

Furthermore, coke ovens used to heat coal to produce carbon that is as an ingredient in steel production. This process releases ammonia gas as a by-product which can be collected and be used to react with sulfuric acid to form ammonium sulfate, used as a fertiliser in the agricultural industry, another source of revenue for the steel manufacturers.

In the ammonia manufacturing plant designed to produce ammonia as a cleaning product, any unreacted hydrogen and nitrogen gas is recycled and fed back into the compression chamber for a further reaction to occur. This therefore reduces the amount of hydrogen and nitrogen lost and lowers the cost of ammonia. The process of recycling unreacted hydrogen and nitrogen gases will also drive or shift the equilibrium position to the right to form more ammonia when the recycled nitrogen and nitrogen gas is re-added into the reaction vessel.

• NOTE: Ammonia that is produced can also be removed from the reaction vessel to help drive the reaction forward (i.e to the right) to form more ammonia.

3.1 evaluate the factors that need to be considered when designing a chemical synthesis process, including but not limited to

e) environmental, social and economic issues

Case Study: Sulfuric Acid Production

Sulfuric acid is the strongest acid we have in the lab. We are already familiar with some of its uses, particularly its ability to catalyse a number of reactions, eg. conversion of ethanol to ethene and the process of esterification.

Sulfuric acid is a dehydrating agent. This makes it a good choice for reactions involving the removal of water.

75% of the sulfuric acid produced in Australia is used in the manufacture of superphosphate fertilisers. It is also used in the production of:

• Viscose rayon and other synthetic fibres
• Ethanol (hydrating ethene)
• Paint, plastic and paper pigments
• Detergents
• Explosives
• Drugs
• Dyes
• Lead-acid batteries
• Steel production
• Oil refining
• Ore refining

The industrial production of sulfuric acid involves three important steps.

1. S(s) + O2(g) → SO2(g) ΔH = -297 kJ
2. 2SO2(g) + O2(g) ⇌ 2SO3(g) ΔH = -198 kJ
3. SO3(g) + H2O(l) → H2SO4(l) ΔH = -133 kJ

Frasch Process

Liquid sulfur is sprayed into oxygen-enriched, dry air at normal atmospheric pressure. Water vapour is removed from the air using a dehydrating agent (eg sulfuric acid) This reaction is strongly exothermic, hence the products need to be cooled (eg from 1000oC to 400oC).

The first step involves the extraction of sulfur from mineral deposits. (Sulfur occurs in element form as well as in compounds of sulfates and sulfides). The process used to extract sulfur is called the Frasch process.

Water at high temperature and pressure (160oC) is forced into the sulfur deposit through a pipe. This melts the sulfur (MP = 113oC) forming an emulsion. Air is blown down another pipe forcing the sulfur/water emulsion up a third pipe. Once the mixture reaches the surface it cools and the sulfur solidifies. It can easily be separated from the water as it is insoluble. The low density of sulfur enables it to be blasted to the surface using compressed air. 50% of the sulfur used in the production of sulfuric acid is extracted using the Frasch process.

Sulfur is a relatively stable element which itself does not cause an environmental problem. However, sulfur is readily oxidised to sulfur dioxide or reduced to hydrogen sulfide. These are gases which pollute the atmosphere. It is important that neither oxidation nor reduction occurs during the Frasch process.

Water used in the Frasch Process may be contaminated with impurities and cannot be released into the environment. Instead it is recycled and reused.

Case Study: Detergents

Early detergents were made from benzene derivatives. These had poor biodegradability and produced excessive amounts of foam when washed into waterways.

To overcome problems with hard water, detergents may contain additional substances such as phosphate derivatives. These were designed to reduce water hardness and increase pH to optimum levels for the surfactant. Unfortunately phosphates are a significant contributing factor to eutrophication and the discharge of such detergents into our waterways can lead to algal blooms and create issues for flora and fauna of the waterways, as well as affecting the potability (drinkability) of the water.

Cationic synthetic detergents can act as antiseptics and biocides. This can have an impact in waterways if they are rinsed down the sink. Sewage treatment plants rely on bacteria to decompose the sewage and these bacteria may be killed by the biocidic action of these detergents if they reach the plant.

Many modern detergents and soaps are made from natural, organic materials and are biodegradable, reducing their impact on our waterways. They are readily broken down into harmless substances, such as carbon dioxide and water.

Contact Process

Step 2.

2SO2(g) + O2(g) ⇌ 2SO3(g) ΔH = -198 kJ

A catalyst (vanadium oxide) is also used in this step. It increases the reaction rate. The fact that the forward direction is exothermic places a counterbalance between increasing the reaction rate using temperature and favouring the endothermic formation of the reactants. To overcome this, the reaction occurs through contact with the catalyst at 550oC.

After about 70% of the SO2 reacts the remaining mixture is cooled to 400oC and passed over a second catalyst bed. The reaction at this temperature is slower but the yield is higher. To further reduce the concentration of sulfur dioxide a third catalyst bed may be used. The quantity of sulfur dioxide must be below 0.3% in order for it to be safely released into the environment.

Step 3. SO3(g) + H2O(l) → H2SO4(l) ΔH = -133 kJ/mol

Once the sulfur trioxide forms it can be dissolved in water droplets to form sulfuric acid. The mist that forms may be difficult to separate from the sulfur trioxide gas. An alternative is to pass the sulfur trioxide directly into concentrated sulfuric acid to form oleum (H2S2O7). Oleum can then be reacted with water to form sulfuric acid. This is a much safer step as the reaction between sulfuric acid and water is highly exothermic.