Nitrogen vs Air in Car Tyres

The tread on my car tyres had worn down, so I popped into my local tyre retailer to buy 4 new tyres.

"We can inflate your new tyres with air, or, for an extra $5 per tyre we can fill them with nitrogen gas", the sales person told me, " Nitrogen gas doesn't react with tyre and rim material so your tyres will last longer, it will help maintain the pressure in your tyres so you won't need to check your tyre pressure as often, and it reduces the running temperature of the tyres so your tyres are less likely to explode."

Is this just marketing hype or are there good reasons for choosing to inflate your car tyres with nitrogen instead of air?

Read the March 2020 edition of AUS-e-NEWS to find out more.

Subscribe to AUS-e-NEWS, AUS-e-TUTE's free quarterly newsletter for chemistry teachers and students, at https://www.ausetute.com.au/ausenews.html

Gutful of Gas

People can eat interesting things.
In the 1960s in Australia, Leon Samson was eating razor blades to amuse audiences, he even started to eat, slowly, in bits, a car for a bet. A little later, France produced Michel Lotito who was also eating odd things to entertain us, including an airplane between 1978 and 1980.

A razor blade is made out of steel, a particular kind of steel known as razor blade steel. A 2.61 gram razor blade contains between about 13% chromium, 0.6% carbon, and  the rest is iron. Samson would chew up a razor blade and swallow it.

After leaving the mouth, the chewed-up bits of razor blade travel to the stomach. An empty stomach has a volume of about 75 mL but when we eat the volume of the stomach can expand out to about 1 L. Protein-digesting enzymes known as proteases are released into the stomach to help begin the break up of the proteins like you find in meat, fish, eggs and cheese. The optimum pH for these proteases is at about pH 2, so hydrochloric acid is also released into the stomach.

So the small bits of razor blade now find themselves surrounded by hydrochloric acid. Hydrochloric reacts with metals, like the iron in steel, to produce hydrogen gas. Now gases have an interesting property, they expand out to fill the available space. This suggests that eating razor blades might result in a feeling of being bloated. Thankfully, it appears that it takes about 24 hours for the complete reaction between a razor blade and hydrochloric acid in the stomach, plenty of time to remove the gas build-up via burping or, um, farting.

Hydrogen gas is commonly found in our intestines, along with other gases like carbon dioxide, oxygen and nitrogen. Some of these gases get there when we swallow air along with our food and drink, but they also come from chemical reactions inside our bodies. If our bodies are healthy and working well, all the usual gases will be present in the usual concentration, but if there is something wrong, if we are sick, the nature and composition of the gases will change.

Researchers at RMIT University in Melbourne, Australia, have developed and trialed an "ingestible electronic capsule" which is capable of sensing and measuring the gases in the gut like hydrogen, carbon dioxide and oxygen. These capsules offer a new, non-invasive way to monitor the health of our gut. The capsules can be collected after they have been excreted (apparently painlessly).

Reference: 
Kourosh Kalantar-Zadeh, Kyle J. Berean, Nam Ha, Adam F. Chrimes, Kai Xu, Danilla Grando, Jian Zhen Ou, Naresh Pillai, Jos L. Campbell, Robert Brkljača, Kirstin M. Taylor, Rebecca E. Burgell, Chu K. Yao, Stephanie A. Ward, Chris S. McSweeney, Jane G. Muir, Peter R. Gibson. A human pilot trial of ingestible electronic capsules capable of sensing different gases in the gut. Nature Electronics, 2018; 1 (1): 79 DOI: 10.1038/s41928-017-0004-x

Suggested Further Reading 
Experimental Design: http://www.ausetute.com.au/experimentd.html
Name and Formula of Binary Covalent Compounds: http://www.ausetute.com.au/namcform.html
Percentage Composition:  http://www.ausetute.com.au/percentc.html
Mass-moles Calculations: http://www.ausetute.com.au/massmole.html
Molar Gas Volume Calculations: http://www.ausetute.com.au/molarvol.html
Ideal Gas Law Calculations:  http://www.ausetute.com.au/idealgas.html
Metal + Non-Oxidising Acid Reaction: http://www.ausetute.com.au/metalhcl.html
Hydrogen Ion Concentration of Strong Acids: http://www.ausetute.com.au/hstronga.html 
Reaction Calculations: Mass and Moles http://www.ausetute.com.au/molreact.html 

Suggested Study Questions

1.   Design an experiment to determine how long it would take for a razor blade to react completely with the hydrochloric acid in the stomach. Investigate ways to speed up, or, to slow down, this reaction.
2. Give the formula for each of the following:
   • hydrogen gas
   • nitrogen gas
   • oxygen gas
   • carbon dioxide gas
   • hydrochloric acid
3.   Determine the percentage composition of each of the following compounds
   • hydrogen chloride gas
   • carbon dioxide gas
4. From the information in the article, calculate the mass of each of the following elements found in a razor blade:
   • iron
   • chromium
   • carbon
5. Using the information above, calculate the moles of each of the following elements found in a razor blade:
   • iron
   • chromium
   • carbon
6. Calculate the moles of hydrogen gas that occupy the entire volume of a "full" stomach under the following conditions:
   • 0^oC and 100 kPa
   • 25^oC and 100 kPa
   • 37^oC and 100 kPa
7. Write a balanced chemical equation for the reaction between the hydrochloric acid in the stomach and the iron in a razor blade.
8. Calculate the concentration of acid released into the stomach using the information in the article. 
9. Use the balanced chemical equation to determine the volume of hydrogen gas produced when all the iron in a razor blade has reacted with hydrochloric acid.
10. Consider all the information in the article, and the calculations you have performed so far. Explain why it takes 24 hours for a razor blade to be completely digested in the stomach.

AlP Rat Poison

Dozens of mysterious sealed silver canisters containing aluminium phosphide have washed up on Australian beaches between 2012 and 2015. The Australian Maritime Safety Authority (AMSA) suspects all the canisters have come from the same ship which dumped or lost its cargo in the Pacific Ocean. Aluminium phosphide is used as a fumigant to poison rats on ships.

When solid aluminium phosphide, AlP, is exposed to water, it releases highly toxic phosphine gas, PH₃, which smells like rotting fish. The chemical reaction can be represented by the balanced chemical equation shown below:
AlP_(s) + 3H₂O_(l) → PH_3(g) + Al(OH)_3(aq)

This is a proton-transfer reaction in which water is acting as Brønsted-Lowry acid by donating a proton to phosphorus. Phosphorus is therefore acting as a Brønsted-Lowry base by accepting a proton from water. Aluminium phosphide will react with acids according to the following chemical
AlP_(s) + 3H⁺_(aq) → PH_3(g) + Al³⁺_(aq)

These reactions make aluminium phosphide a good choice for ridding a ship of rats.
Firstly, as a solid, AlP can easily be stored as pellets in air-tight, water-tight, containers until it is ready to be used. When required, the pellets can be scattered in the effected area . In the humid air aboard ship, the AlP will start reacting to produce toxic phosphine gas, that is, the area will be fumigated. But it is also possible to entice rats to eat AlP pellets mixed with food, in which case it will act as pesticide, because on entering the acidic stomach of the rat, it will produce the toxic phosphine.

Aluminium phosphide is  a very effective way to get ride of rats, so much so, that is widely used in agriculture to remove rats from grain silos.

References:
"Toxic canisters washing up on Australian beaches pose serious health risk"
 http://www.smh.com.au/environment/toxic-canisters-washing-up-on-australian-beaches-pose-serious-health-risk-20150917-gjp5se.html

"Controlling rabbits with aluminium phosphide tablets"
 http://agriculture.vic.gov.au/agriculture/farm-management/chemical-use/publications/chemical-industry-news/chemical-industry-news-no.-75-summer-autumn-2013

"Phosphine fumigation"
https://www.worksafe.qld.gov.au/injury-prevention-safety/hazardous-chemicals/specific-hazardous-chemicals/phosphine-fumigation

Further Reading
Definition of Acids and Bases
Proton-transfer Reactions
Mass-mole Calculations
Molar Volume of Gases

Suggested Study Questions:

1. The symbols of some elements are listed below. Name each element.
   • Al
   • P
   • H
   • O
   • K
   • He
   • At
2. Calculate the amount of aluminium phosphide in moles given the masses of AlP given below:
   • 10 g
   • 10 kg
   • 10 mg
   • 10 μg
3. Calculate the moles of phosphine gas produced when each mass of AlP below reacts with excess water in a ship's hull:
   • 10 g
   • 10 kg
   • 10 mg
   • 10 μg
4. Based on your answers to question 3 above, calculate the mass of phosphine produced for each mass of AlP used.
5. Your ship is sailing towards eastern Australia and has just crossed the Tropic of Capricorn. You have been asked to estimate the volume of phosphine gas that will be produced when you release AlP pellets into the ships hold. Which molar gas volume will you use; 22.71 L or 24.79 L ? Explain your answer.
6. Rats are currently infesting a small part of your ship, about 150 m³. How much solid AlP would be required to fumigate this area, but not leave any AlP residue left over?
7. The Cook has already tried to fumigate the pantry and is sure there is a silver canister around that still contains some AlP, it could be in the pile of empty canisters, or, it could be in the pile of full canisters. No-one wants to kill themselves by opening the canisters to find out, so can you suggest a method that could be used on board ship to determine how much AlP is present in each canister.
8. Explain why the reaction between aluminium phosphide and water is described as a proton-transfer reaction and not as a redox reaction.
9. Explain why, even though aluminium phosphide and phosphine are toxic, it is considered safe to use these to fumigate silos containing grain which will be eaten by humans.
10. The silver canisters that have washed up on Australian beaches have no labels, presumably these have come off while they were in the ocean. You have been asked to design new labels for the canisters. The labels must include suitable safety and handling information.

Nitrous Oxide Ban

On Saturday 29^th August 2015, Brisbane's Couriermail reported that, "Supermarkets have removed nitrous oxide canisters from their shelves because of the alarming trend of people using the gas to get high".
Nitrous oxide, N₂O, commonly known as "laughing gas", is found as an additive in food (E942). It has been used as an aerosol propellant in whipped cream canisters and cooking sprays, and has also been used as an inert gas to displace oxygen from food packages like potato chips to prevent the food from spoiling. The nitrous oxide canisters being removed from supermarket shelves refer to the small canisters used to re-charge re-usable whipped cream dispensers. Nitrous oxide is soluble in fats and oils so it is pumped into the fatty cream inside the dispenser where it dissolves. When the cream leaves the dispenser, the gas pressure inside the cream is greater than atmospheric pressure, so the nitrous oxide gas leaves the cream quickly, creating a foam. This "whipped cream" has about 4 times the volume of the original cream. However, if the whipped cream is left standing for some minutes, the gas pressure equilibrates with atmospheric pressure and the cream reverts to the "unwhipped" state.
Nitrous oxide has been the preferred choice for whipped cream dispensers because it does not react with the food and prevents it from spoiling. If air were used, the oxygen in the air would accelerate the rate of oxidation causing the cream to spoil. Carbon dioxide gas is not used because it would dissolve in the water present in the cream which would lower the pH and the cream would curdle.

In the laboratory, nitrous oxide can be produced by the thermal decomposition of ammonium nitrate (which is explosive!):

NH₄NO₃(s) → 2H₂O(g) + N₂O(g)

It can also be produced by heating a mixture of sodium nitrate and ammonium sulfate:

2NaNO₃ + (NH₄)₂SO₄ → Na₂SO₄ + 2N₂O + 4H₂O

Nitrous oxide has been used in dentistry as an anaesthetic for more than a century.
The chemistry of how nitrous oxide reacts in the body to produce a "high" is not well understood, but research continues.

Further Reading
Density
Types of Chemical Reactions
Molecular Formula
Name and Formula of Covalent Compounds
Fats and Oils 
Solubility and Le Chatelier's Principle
Mass-mole Calculations
Mass and Moles in Chemical reactions 
Molar Volume of Gases 
Safety in the Laboratory

Suggestion Study Questions:

1. Give the molecular formula for each of the following compounds: 
   • nitric oxide
   • nitrogen dioxide
   • nitrous oxide
   • dintrogen tetroxide
2.  Name each of the following compounds:
   • CO₂
   • CO
   • SO₂
   • SO₃
   • PCl₃
   • PCl₅
3.  Draw possible Lewis (electron dot) structures for each of the following compounds:
   • nitric oxide
   • nitrogen dioxide
   • nitrous oxide
   • dintrogen tetroxide
4.  "..whipped-cream has four times the volume of the original cream", describe what has happened to the density of the cream as a result of the "whipping" process.
5. When I order an "iced-coffee" at a Cafe, they usually top it with some whipped-cream from a canister . Why does the whipped-cream float on top of the ice-coffee drink?
6. If I don't the spoon the whipped-cream off the top of my "iced-coffee" I can stir it into the drink quite easily without it floating to the surface again. Explain why this happens.
7. Explain what is meant by the term "fats and oils" when used by a Chemist and give examples.
8. Most aerosol cans contain a warning to the effect that you should not heat the can. Explain what would happen if you heated a dispenser of whipped-cream containing nitrous oxide. 
9. Imagine placing 10 grams of ammonium nitrate in a 500 mL sealed steel can, which you then (very irresponsively!) throw onto a bonfire.
• Calculate the moles of ammonium nitrate in the can
• Calculate the moles of nitrous oxide gas produced 
• Calculate the volume of this number of moles of gas at 25^oC
• Predict what will happen to the sealed steel can in the bonfire. 
10. Your company, "Ammo-nite", sells ammonium nitrate to farmers for use as a fertiliser. You have been asked to re-design the 100 g package so that it includes relevant safety precautions.

Henry's Law and the Solubility of Gases

What is the relationship between the solubility of a gas in a solvent, and the pressure of the gas?
What happens to the solubility of gases as the temperature changes?
If you are asking these questions, you need AUS-e-TUTE's new Henry's Law resources!

AUS-e-TUTE Members will find the new resources listed under Gas Laws in the Members Only Test Centre. Where relevant to your syllabus, links to these resources have also been added to your Syllabus Study Guide.

If you are not an AUS-e-TUTE member, there is a "free-to-view" tutorial available for evaluation purposes at http://www.ausetute.com.au/henryslaw.html

Then, when you are ready to start improving your chemistry results, you can join AUS-e-TUTE at
http://www.ausetute.com.au/register.html

Chlorine as a Chemical Weapon

World War I, saw the birth of the ANZAC legend*, and also the widespread use of "chemical warfare".

The first gas used by the German military as a killing agent was chlorine gas, also known as bertholite at this time.

Chlorine is a powerful irritant, which can damage the eyes, nose, throat, and lungs. Prolonged exposure to high concentrations, 1,000ppm, can cause death by asphyxiation.
Chlorine gas reacts with water in the tissues of the body to produce hydrochloric acid:

2Cl_2(g) + 2H2O_(l) → 4HCl_(aq) + O_2(g)

The chlorine gas was released from cylinders facing the enemy trenches in a favourable wind. The grey-green cloud of chlorine gas would then drift across the enemy positions.
The density of chlorine gas at 0^oC and 101.3kPa (1 atm) is 3.2g/L, while the density of air at the same temperature and pressure is 1.2754 g/L. Because chlorine gas is more dense than air, it would be more concentrated at the bottom of the trench, and less concentrated at the top. Those who suffered the worst effects were often the wounded lying on the ground or on stretchers.

Initially, German troops were issued with gauze pads filled with cotton, and bottles of bicarbonate solution. When the chlorine gas was to be released, the soldiers would dampen the gauze pad with the bicarbonate solution and breathe through it. The bicarbonate would neutralize the hydrochloric acid produced. If sodium bicarbonate solution were used, the reaction would be:

HCl_(aq) + NaHCO_3(aq) → NaCl_(aq) + CO_2(g) + H₂O_(l)

As other gases were being developed for use as chemical weapons, the need for better protection became important. One of the earliest devices was a hood with eyepieces. The hoods could be impregnated with sodium hyposulphite (sodium thiosulfate). Sodium thiosulfate reacts with dilute acids to produce sulfur, sulfur dioxide and water:

Na₂S₂O₃ + 2HCl → 2NaCl + S + SO₂ + H₂O

The gas mask was developed later. It was composed of an impervious mask and a box respirator or canistor. Air came through the canister which contained charcoal and granules of soda-lime, a mixture of sodium hydroxide and calcium hydroxide.

*ANZAC (Australian and New Zealand Army Corps) Day is commemorated each year on 25th April, marking the anniversary of the first major military action fought by Australian and New Zealand forces during the First World War. ANZAC Day ceremonies, which are held in towns and cities all over Australia and New Zealand, typically include an introduction, a hymn, a prayer, an address, the laying of wreaths, a recitation, the Last Post, a period of silence, and either the Rouse or the Reveille, and the national anthem. Anzac Day has evolved to acknowledge the sacrifice and service of subsequent wars.

Further Reading
Temperature Conversions
Density Calculations
Elements and Compounds
Writing Ionic Formula
Balancing Chemical Equations
Molecular Mass
Definitions of a Mole
Ideal Gas Law
Acid-Base Titration Calculations

Study Questions

1. Draw a table listing each element and each compound mentioned in the article above.
2. In April 1915, the German Army is said to have stockpiled 168 tons of chlorine which was contained in 5,370 cylinders. on average:
   • how many kilograms of chlorine gas was contained in each cylinder?
   • how many moles of chlorine gas was contained in each cylinder?
   • what volume would this moles of gas occupy at 25^oC and 101.3 kPa (1 atm)?
     
3. Using your calculations in part 2, calculate the density of chlorine gas at 25^oC and 101.3 kPa (1atm) in g/L.
4. Compare the density of chlorine gas calculate in part 3, to the density of chlorine gas given in the article above. Account for the difference in the two density values.
5. Convert 1,000ppm chlorine gas to a concentration in:
   • mg/L
   • g/L
   • mol/L
6. What mass of HCl could be produced from 1L of 1,000ppm chlorine gas?
7. What mass of sodium bicarbonate would be necessary to neutralize the amount of hydrochloric acid produced in question 6 above?
8. What mass of sodium thiosulfate would be necessary to neutralize the amount of hydrochloric acid produced in question 6?
9. Write a balanced chemical equations for:
   • the reaction between hydrochloric acid and calcium hydroxide
   • the reaction between hydrochloric acid and sodium hydroxide
10. Explain why a gas mask containing soda-lime might be preferable to one containing sodium thiosulfate.
    
    
    
    

Colouring Plastics

At a temperature of 30.1^oC and a pressure of about 73.8 bar, carbon dioxide goes into a super critical state that gives the gas solvent-like properties. In this state it can be introduced into polymers, acting as a "carrier" in which dyes, additives, medical compounds and other compounds can be dissolved:

• liquid carbon dioxide is pumped into a high-pressure container with the plastic components to be impregnated
• temperature and pressure are increased steadily until carbon dioxide reaches its supercritical state
• pressure increased to 170bar and the powdered pigment dissolves completely
• the dissolved pigment diffuses with the carbon dioxide gas into the plastic
• high-pressure container is opened, the carbon dioxide gas escapes but the pigment remains attached to the polymer

Tests have shown that nanoparticles of antibacterial agents can be impregnated into polycarbonate and that E-coli bacteria placed on the surface will be killed. Tests conducted with silica and with the anti-inflammatory active pharmaceutical ingredient flurbiprofen were also successful.

The process is suitable for use with amorphous or partially crystalline polymers such as nylon and polycarbonate, but cannot be applied to crystalline polymers.

Reference
Fraunhofer-Gesellschaft (2011, January 4). Impregnating plastics with carbon dioxide. ScienceDaily. Retrieved January 10, 2011, from http://www.sciencedaily.com­ /releases/2011/01/110103110206.htm

Further Reading
Temperature Conversions
Chemical and Physical Changes
Lewis Structures
Molecule Polarity
Intermolecular Forces
Ideal Gas Law

Study Questions:

1. Convert the information provided in the dot points in the article into a flow chart.
   
2. Give the molecular formula for carbon dioxide.
3. Draw a Lewis Structure (electron dot diagram) for carbon dioxide.
4. Is carbon dioxide a polar or non-polar molecule? Explain your answer
5. What type of solvents would be the best to use to dissolve carbon dioxide?
6. What type of solutes do you expect to dissolve in carbon dioxide? Explain your answer.
7. Given that 1bar is equivalent to 100kPa, convert the following pressures in bars to pressures in atmospheres:
   
   • 73.8bar
   • 170bar
8. At 25^oC and 1 atmosphere pressure, 1 mole of gas has a volume to 24.47L. Assuming ideal gas behaviour, what volume would this gas occupy at 30.1^oC and 73.8bar?
9. Calculate the volume of gas above at the same temperature but at a pressure of 170bar.
10. Calculate the volume per molecule of carbon dioxide gas in question in 9.
11. Do you think it is reasonable to assume ideal gas behaviour at this temperature and pressure? Explain your answer.
12. What do you think the term "supercritical state" as used in this article might mean?
    
    

Champagne Bubbles

Tiny bubbles are the essence of fine champagnes and sparkling wines.
The bubbles, formed during the release of large amounts of dissolved carbon dioxide, help transfer the taste, aroma, and mouth-feel of champagne. Scientists have thought that the act of pouring a glass of champagne could have a big impact on gas levels in champagne and its quality.
Scientists in France have studied carbon dioxide loss in champagne using two different pouring methods:

• pouring champagne straight down the middle of a glass
• pouring champagne down the side of an angled glass

They found that pouring champagne down the side of an angled glass preserved up to twice as much carbon dioxide in champagne compared to pouring it straight down the middle.
They also showed that cooler temperatures help reduce carbon dioxide loss.

Reference:
Liger-Belair et al. On the Losses of Dissolved CO2 during Champagne Serving. Journal of Agricultural and Food Chemistry, 2010; 58 (15): 8768 DOI: 10.1021/jf101239w

Study Questions

1. What is the formula for carbon dioxide?
2. Is carbon dioxide a polar or non-polar molecule?
3. What is the structural formula for ethanol?
4. Is ethanol a polar or non-polar molecule?
5. Would you expect carbon dioxide to dissolve in ethanol? Explain your answer.
6. Describe an experiment you could conduct to test the hypothesis that cooler temperatures reduce carbon dioxide loss in champagne.