# Outline three uses of sulfuric acid in industry

This dotpoint is largely introductory, so don’t feel that you need to go into a great deal of detail for three separate uses- the keyword is ‘outline’.

• Dehydration: Sulfuric acid is a powerful dehydrating It is used to increase yield and act as a catalyst during esterification, and is also used in the dehydration of ethanol into ethylene and water.
• Cleaning metals: Sulfuric acid is used to clean pieces of metal, a task that is necessary before a piece of metal can be galvanised or coated in
• Production of fetilisers: Sulfuric acid plays an important role in the production of a few different In a coke oven, sulfuric acid can extract ammonia from a mixture of gases by reacting with the gas. This can later be used to produce ammonium sulfate fertiliser, In addition, sulfuric acid can convert insoluble calcium phosphate into a soluble mixture which can then be used in phosphate fertilisers.

Remember- Sulfuric acid can be used in the dehydration of ethanol and alkanols & alkanoic acids during esterification. Sulfuric acid can also be used to clean metals, as well as produce ammonia sulfate and superphosphate fertilisers.

# Describe and explain the exothermic nature of sulfuric acid ionisation

The ionisation of sulfuric acid is a highly exothermic reaction, meaning that a great deal if heat is released upon ionisation. This is because concentrated sulfuric acid is virtually 100% molecular, meaning there are very little ions. Upon ionisation, the breaking of the covalent bonds present in the acid releases a vast amount of energy, thus leading to an exothermic reaction.

# Identify and describe safety precautions that must be taken when using and diluting concentrated sulfuric acid

Safety is a point that has really been pushed in the new HSC. As such, make sure you are prepared to answer safety questions on any topic, in any module. Fortunately, safety procedures are fairly common sense, as the following list will prove.

• Wear full protective lab gear when conducting this experiment. This includes a lab coat and Appropriate footwear and gloves are also recommended.
• Have a supply of water, such as a tap, nearby should you need to flush a part of the body with running
• Always add acid to water, never the other way
• Store and use sulfuric acid in volumes smaller than 1
• Ready a base in the case that a spill must be Sodium carbonate is one such example.
• Avoid conducting experiments at eye level, or near the

# Use available evidence to relate the properties of sulfuric acid to safety pre- cautions necessary for its transport and storage

As mentioned previously, concentrated sulfuric acid is virtually 100% molecular. For this very reason steel containers can be used to transport concentrated sulfuric acid, as there are no ions present to attack the metal.

If dilute sulfuric acid is being transported however, glass is the preferred medium of transport, followed by plastic.

In general, regardless of the method of transport, sulfuric acid should be split into many small containers, minimising the danger on the off-chance that a container breaks or there is a contamination of concentrated sulfuric acid with water.

In addition, excess amounts of a substance to neutralise the spill is highly recommended. One example is sodium carbonate, which even when used in excess is not likely to create a situation of danger.

# Describe the processes used to extract sulfur from mineral deposits, identifying the properties of sulfur which allow its extraction and analysing potential environmental issues that may be associated with its extraction

Note that this dotpoint is essentially three points in one. To make sure you cover all possible examinable materials, as well as to break this dotpoint into more manageable portions, this guide will treat each point separately under different subheadings.

<h2>Processes used to extract sulfur from mineral deposits

One of the more common methods of extracting sulfur from mineral deposits is known as the Frasch Process. Using this process, one large pipe is placed into a sulfur deposit. Within this pipe are three pipes, all placed concentrically (All pipes form circle with equal centres, but differing diameters). Water superheated to 160◦C is funnelled down the outermost pipe, melting the sulfur in the deposit, which combines with the water to form an emulsion. Compressed air is then sent down the innermost pipe, pushing the emulsion up to the surface through the middle pipe.

Upon cooling, 99.5% pure sulfur is obtained as it separates from the water due to its insolubility.

## Properties of sulfur which allow for its extraction

Two properties of sulfur allow for its extraction through the Frasch Process. The first is its relatively low melting point of 113◦C, which allows the sulfur in the deposits to dissolve in the water and be pushed up to the top. The second property is the insolubility of sulfur, which allows for the easy extraction at the surface.

## Environmental issues associated with the Frasch Process

The sulfur brought to the surface can be readily oxidised or reduced to form sulfur dioxide or hydrogen sulfide respectively. Both of these pollutants can have severe consequences even when in small concentrations, and as such the water used to flush the sulfur from the deposit must be reused, or recycled before it can be disposed of.

In addition, should the sulfur dioxide find its way into the atmosphere, large concentrations will result in the formation of acid rain, which can have serious repercussions upon both manmade infrastructure as well as flora and fauna, particularly marine ecosystems.

Remember- Sulfur can be extracted from mineral deposits through the Frasch Process. Superheated water is flushed down the outermost pipe of a set of concentric pipes, and compressed air sent down the innermost pipe sends the emulsion to the surface, where the sulfur can be gathered due to its insolubility. The process must be carefully monitored, as sulfur dioxide and hydrogen sulfide are serious pollutants, with the former having the potential to lead to acid rain.

# Outline the steps and conditions necessary for the industrial production of $H_2SO_4$ from its raw materials (inluding ’Describe the reaction conditions necessary for the production of $SO_2$ and $SO_3$’ and ’Apply the relationship between rates of reaction and equilibrium conditions to the production of $SO_2$ and $SO_3$’)

These three dotpoints are better approached as one, as they all lead to the production of sulfuric acid. Make sure you break down the production process into three parts, as you’ll find that the contact process (As it is called) makes a lot more sense when you can picture it as three big steps. Also keep in mind equilibrium considerations (Le Chatelier’s Principle and rates of reaction).

The current method used to produce sulfuric acid was discovered in the early 1800s in a process known as the contact process. Under this process, three main steps can be identified: The production of sulfur dioxide, sulfur trioxide, and sulfuric acid. You will note how these tie into one another as the process progresses.

## The production and preparation of sulfur dioxide

This step is simply a combustion reaction, where liquid sulfur is sprayed into an excess of dry air. These conditions are used to ensure a maximum output of sulfur dioxide while taking care to avoid producing sulfurous acid through the reaction of sulfur dioxide with water. Alternatively, water can be used to wash the gas, and then sulfuric acid is used to dry the gas once more. When produced, the sulfur dioxide reaches temperatures of $1000^oC. However, this must be cooled down to a temperature between$latex 400-450^oC\$ for the next stage of the process.

The catalytic conversion of sulfur dioxide and oxygen into sulfur trioxide

The sulfur dioxide previously produced is combined with oxygen as it is passed over a catalyst bed at approximately 450◦C, converting round 70% of the sulfur dioxide in the mixture into sulfur trioxide. The sulfur trioxide is then removed after the mixture has been cooled, after which the mixture is passed over yet another catalyst bed. This provides a yield of 99.7%, with the remaining sulfur dioxide at safe concentrations for release into the atmosphere.

The ideal conditions for this stage of the process are as follows:

• Low temperatures: The reaction is exothermic, and thus lower temperatures would promote a higher This is the reason why the temperature is dropped from 1000◦C to around 450◦C.
• Catalyst: Catalyst beds of vanadium pentoxide are used as the lower temperatures used to promote a higher yield also promote a slower rate of reaction. Without these catalysts, the reactions is not likely to ever reach a state of
• Pressure: Pressures slightly above atmospheric pressure are used as higher pressure promotes higher yield in this However, the effect is minimal, thus the pressure is not raised any higher as the costs and risks of doing so outweighs any gain in yield.
• Excess oxygen: A small excess of oxygen is also used to promote a higher yield of sulfur

The conversion of sulfur trioxide into sulfuric acid

Although sulfur trioxide can theoretically be added directly to water to form sulfuric acid, this is undesirable as the reaction is so exothermic that acid vapour is formed rather than a liquid.

Thus, to get around this problem the sulfur trioxide is added to existing sulfuric acid, producing $H_2S_2O_7$, a substance known as oleum. The oleum is then reacted with water, producing sulfuric acid.

Remember- Liquid sulfur is first sprayed into an excess of dry air to form sulfur dioxide, which is added to oxygen and passed over catalyst beds of vanadium pentoxide twice to form sulfur trioxide. The sulfur trioxide is then added to sulfuric acid to form oleum, which is reacted with water to produce more sulfuric acid.

In the production of sulfur trioxide, lower temperatures and high pressure promotes increased yield. However, the rate of reaction and safety and cost must be given consideration. As such, a com- promising 400-450C is used with a vanadium pentoxide catalyst and a pressure of a little over atmospheric pressure.

# Gather, process and present information from secondary sources to describe the steps and chemistry involved in the industrial production of $H_2SO_4$ and use available evidence to analyse the process to predict ways in which the output of sulfuric acid can be maximised

As a note, I will mention that any question where you are required to note how to maximise yield is essentially a practice in Le Chatelier’s Principle. As such, go through temperature and pressure. However, with these two, note that a low temperature reduces the rate of reaction, so for exothermic reactions a catalyst is likely to be required. In addition, higher pressure will not influence yield much if the imbalance in the moles of gas on the reactants side does not vary much from the moles of gas in the products side (In the case of the production of sulfur trioxide in the contact process, the ratio is only 3:2). As such, the cost and safety concerns of higher pressure might be given a larger weighting. Lastly, consider what products can be removed and what reactants can be used in excess so as to further push the equilibrium to the right. In the case of producing sulfur trioxide, sulfur trioxide can be removed and excess oxygen can be used.

# Perform first-hand investigations to observe the reactions of sulfuric acid acting as: An oxidising agent, and a dehydrating agent

For this dotpoint there is a very simple experiment which illustrates the use of sulfuric acid as both an oxidising agent and as a dehydrating agent.

Materials:

• One 250mL beaker
• Sucrose
• Concentrated sulfuric acid

Procedure:

1. Pour sucrose into an empty beaker, filling it up to one third of
1. Place the beaker in a fume
1. Pour 25mL of concentrated sulfuric acid into the beaker, and step

Expected Results:

It may take a few seconds for the beaker to show any change, but once it does the result will be rather quickly achieved. The sucrose should be replaced by a much larger carbon structure, which demonstrates sulfuric acid agent as a dehydrating agent:

Next, the carbon is oxidised, forming carbon dioxide, sulfur dioxide and water. These fumes are dangerous in quantity, so make sure the experiment is in fact conducted in a fume cupboard.

Describe, using examples, the reactions of sulfuric acid acting as: An oxidising agent, and a dehydrating agent

Sulfuric acid acts as an oxidising agent when a metal is added to sulfuric acid. In the example below, the copper metal is oxidised, losing two electrons and gaining a 2+ charge.

The use of sulfuric acid as a dehydrating agent is evident when examining its use during esterification, and during the dehydration of ethanol, the latter of which is shown below. The concentrated sulfuric acid is not shown as it is in fact a catalyst in the reaction.

Remember- Oxidisation Is Loss, Reduction Is Gain (OILRIG). Sulfuric acid acts as both an oxidising agent and as a dehydrating agent.