October 25 – Sheer Nonsense

Today’s factismal: The first nylon stockings went on sale in 1939.

Back in 1939, women had a big problem: they wanted to wear silk stockings but they couldn’t afford them. The price of a typical pair of silk stockings had risen by more than 50% in the past year alone, thanks to rising demand and embargoes on foreign goods. And even if she could afford the $0.69 ($11.26 in today’s money) that a pair of stockings cost, a woman was likely to see her investment ruined the first time that she wore them. Fortunately, chemistry was about to come to the rescue.

Artificial silk had been known since 1855 when nitrocellulose (aka guncotton or “oops! I blew your legs off”) was turned into fine, extremely flammable threads that became known as “mother-in-law’s silk”. The process was further refined into the creation of rayon from sawdust in the early 1920s, but the threads were coarse and irregular. So scientists searched for an alternative and finally found it in 1935. The nylon silk that they produced was first used to make bristles for toothbrushes; once the process had been refined enough to create long fibers, they started to manufacture stockings, parachute cloth, and other fabric goods.

A war poster encouraging recycling silk and other scarce goods (Image courtesy Truman Library)

A war poster encouraging recycling silk and other scarce goods
(Image courtesy Truman Library)

Their discovery came just in time as many of the traditional sources for rope (hemp from Indonesia), tires (rubber from Indonesia and Thailand), silk fabric (silk from China) and other materials had been embargoed due to concerns about the war that had begun. Thanks to their work, the US was able to substitute synthetic materials for the natural goods; today, many of those synthetic materials are not only still used but often preferred due to their superior quality and strength. If you’d like to learn more about the chemistry behind nylon and other synthetic fabrics, then head on over to Chemspider:
http://www.chemspider.com/

October 23 – We Are All Starstuff

Today’s factismal: There are about as many atoms in 5 1/2 ounces of oxygen as there are stars in the universe.

If you had been a chemist in the 1800s, you would have had a real problem. You knew for a fact that oxygen plus carbon would make water(H2O), but you would be able to say how much oxygen or how much hydrogen was needed to leave nothing but water in the reaction chamber. Sometimes you’d have oxygen left over and sometimes you’d have carbon left over and you’d always have a big mess. It was uncertainties like this that kept chemistry from being an exact science.

The reason that chemistry was an uncertain science was because the number of oxygen atoms in a pound of oxygen is different than the number of hydrogen atoms in a pound of hydrogen. (This is why Mark Whatney blew up the lab in The Martian.) Because chemistry takes place on the atomic scale, you couldn’t just add two pounds of hydrogen to one pound of oxygen and get nothing but water; you had to find some way of scaling the weight (or, more appropriately, the masses) of each chemical so that you’d be adding the right number of atoms. Fortunately, a scientist by the name of Avogardo pointed the way.

Avogardo (or “Avocado” as he is known to all freshman chemistry students) had the bright idea in 1811 that the volume of space taken up by a gas at a given pressure and temperature might be related to the number of atoms in that gas; based on that, he and other scientists were able to derive the relative atomic weights of the elements. It took the chemists nearly a century, but by 1909, we had a periodic table that listed the atomic weight of each element. That allowed us to know exactly how much of each to add in order to get reactions that worked perfectly every time.

There are a mole of stars in the universe (Image courtesy NASA)

There are twenty moles of stars in the universe
(Image courtesy NASA)

Avogardo and the chemists who came after him called the standard amount of stuff a mole (short for “molecular volume”). And, because it was Avogardo’s bright idea that made it all possible, the number of atoms (or molecules) in a mole is known as Avogardo’s number. And it is a mighty large number – there are 6.02 x 10^23 atoms of oxygen in 16 grams (one mole) of oxygen. To give you an idea of how many atoms that is, just go outside tonight and take a look at the night sky. If you were to count every star in every galaxy in the universe, there would be about 10^24 stars. So there are as many atoms of oxygen in ten moles of oxygen as there are stars in the mole of the universe!

Chemists celebrating Mole Day (Image courtesy ACS)

Chemists celebrating Mole Day
(Image courtesy ACS)

In honor of Avogardo’s discovery, today is Mole Day (because it is 10/23 – get it?). So take part in a mole day celebration somewhere. Go eat a mole cake and drink some mole juice. And then make a un-moley mess, just so you can appreciate why chemists were so happy to become an exact science!
http://www.acs.org/content/acs/en/education/students/highschool/chemistryclubs/activities/mole-day.html

September 23 – Hole Lotta Trouble

Today’s factismal: The ozone hole stretched to cover a city for the first time fifteen  years ago.

One of the great successes in pollution control was the 1992 international treaty banning the use of chlorofluorocarbons (CFCs) due to their effect on the ozone layer. Following the signing of the treaty, nations were required to change their refrigerators and hairsprays so that they didn’t use CFCs; the only exceptions were for national security. So with the pollution stopped, the problem was solved, right?

2009 Ozone Hole (Image courtesy NASA)

2009 Ozone Hole
(Image courtesy NASA)

Wrong. The problem with pollution is that it doesn’t stop doing harm just because you’ve stopped putting more trash into the atmosphere. You still have to deal with all of the junk that was put into the atmosphere before you stopped. Some environmentalists call this the “teenager’s room problem”: sure, your kid has gone to college and left his room empty – but you still have the ten years of empty soda cans, candy bar wrappers, and dirty laundry piled in the corners that need to be cleaned out before it can be turned into a sewing room. And that’s where we are with CFCs in the atmosphere. We’ve stopped adding them but we still have to wait for the ones in the air to break down and go away. And, until they do, we will have problems.

This year's ozone hole (Image courtesy MACC)

This year’s ozone hole
(Image courtesy MACC)

In 2000 we saw one example of the sort of problem we’ll have; the ozone hole grew to cover an area three times the size of the continental United States. It got so large that it covered all of Antarctica and part of South America, including the city of Punta Arenas. For two days, the residents were exposed to more UV radiation than normal. Though they haven’t reported much in the way of side effects that is because UV damage is a long-term problem (e.g., skin cancer, glaucoma) caused by a short-term exposure. Fortunately, that was the largest that the ozone hole has ever gotten; since then it has shrunken considerably.

Of course, a hole in the ozone layer isn’t the only problem we’ve got. If you’d like to help monitor air quality, then why not join NASA’s Citizens and Remote Sensing Observation Network Air Quality project?
http://terra.nasa.gov/citizen-science

October 24 – Fine As Silk

Today’s factismal: The first nylon stockings went on sale in 1939.

Back in 1939, women had a big problem: they wanted to wear silk stockings but they couldn’t afford them. The price of a typical pair of silk stockings had risen by more than 50% in the past year alone, thanks to rising demand and embargoes on foreign goods. And even if she could afford the $0.69 ($11.26 in today’s money) that a pair of stockings cost, a woman was likely to see her investment ruined the first time that she wore them. Fortunately, chemistry was about to come to the rescue.

Artificial silk had been known since 1855 when nitrocellulose (aka guncotton or “oops! I blew your legs off”) was turned into fine, extremely flammable threads that became known as “mother-in-law’s silk”. The process was further refined into the creation of rayon from sawdust in the early 1920s, but the threads were coarse and irregular. So scientists searched for an alternative and finally found it in 1935. The nylon silk that they produced was first used to make bristles for toothbrushes; once the process had been refined enough to create long fibers, they started to manufacture stockings, parachute cloth, and other fabric goods.

A war poster encouraging recycling silk and other scarce goods (Image courtesy Truman Library)

A war poster encouraging recycling silk and other scarce goods
(Image courtesy Truman Library)

Their discovery came just in time as many of the traditional sources for rope (hemp from Indonesia), tires (rubber from Indonesia and Thailand), silk fabric (silk from China) and other materials had been embargoed due to concerns about the war that had begun. Thanks to their work, the US was able to substitute synthetic materials for the natural goods; today, many of those synthetic materials are not only still used but often preferred due to their superior quality and strength. If you’d like to learn more about the chemistry behind nylon and other synthetic fabrics, then head on over to Chemspider:
http://www.chemspider.com/

October 23 – It Is Full Of Stars

Today’s factismal: There are about as many atoms in 16 grams of oxygen as there are stars in the universe.

If you had been a chemist in the 1800s, you would have had a real problem. You knew for a fact that oxygen plus carbon would make water(H2O), but you would be able to say how much oxygen or how much hydrogen was needed to leave nothing but water in the reaction chamber. Sometimes you’d have oxygen left over and sometimes you’d have carbon left over and you’d always have a big mess. It was uncertainties like this that kept chemistry from being an exact science.

The reason that chemistry was an uncertain science was because the number of oxygen atoms in a pound of oxygen is different than the number of hydrogen atoms in a pound of hydrogen. Because chemistry takes place on the atomic scale, you couldn’t just add two pounds of hydrogen to one pound of oxygen and get nothing but water; you had to find some way of scaling the weight (or, more appropriately, the masses) of each chemical so that you’d be adding the right number of atoms. Fortunately, a scientist by the name of Avogardo pointed the way.

Avogardo (or “Avocado” as he is known to all freshman chemistry students) had the bright idea in 1811 that the volume of space taken up by a gas at a given pressure and temperature might be related to the number of atoms in that gas; based on that, he and other scientists were able to derive the relative atomic weights of the elements. It took the chemists nearly a century, but by 1909, we had a periodic table that listed the atomic weight of each element. That allowed us to know exactly how much of each to add in order to get reactions that worked perfectly every time.

There are a mole of stars in the universe (Image courtesy NASA)

There are a mole of stars in the universe
(Image courtesy NASA)

Avogardo and the chemists who came after him called the standard amount of stuff a mole (short for “molecular volume”). And, because it was Avogardo’s bright idea that made it all possible, the number of atoms (or molecules) in a mole is known as Avogardo’s number. And it is a mighty large number – there are 6.02 x 10^23 atoms of oxygen in 16 grams (one mole) of oxygen. To give you an idea of how many atoms that is, just go outside tonight and take a look at the night sky. If you were to count every star in every galaxy in the universe, there would be about 10^23 stars. So there are as many atoms of oxygen in a mole of oxygen as there are stars in the mole of the universe!

Chemists celebrating Mole Day (Image courtesy ACS)

Chemists celebrating Mole Day
(Image courtesy ACS)

In honor of Avogardo’s discovery, today is Mole Day (because it is 10/23 – get it?). So take part in a mole day celebration somewhere. Go eat a mole cake and drink some mole juice. And then make a un-moley mess, just so you can appreciate why chemists were so happy to become an exact science!
http://www.acs.org/content/acs/en/education/students/highschool/chemistryclubs/activities/mole-day.html

October 21 – Ka-Boom!

Today’s factismal: The world’s most famous chemist is known mostly for his charitable work.

Mining in the 1800s was a nerve-wracking job. Not only did you have to worry about bad air, cave-ins, and flooding, but the explosive of choice was almost as unstable as your boss. Known as nitroglycerin, it was easy and cheap to make but tricky and difficult to transport and use. It would go off if it got too hot or too cold, if it was jostled too much or not enough, or if it just didn’t like the way you looked at it. It frequently destroyed the factories where it was being made, and its habit of exploding while being moved led to laws against it being transported across state lines.

But in 1867, Alfred Nobel found a way to tame the beastly blast. By mixing the nitroglycerin with diatomaceous earth or sawdust, he was able to make it more stable and less dangerous. It could be easily stored and transported and could even be measured on the spot with very little chance of losing an arm. Needless to say, dynamite was an immediate hit and made Alfred Nobel very, very rich indeed. But every silver lining has a cloud, and dynamite had a big one.

Because nitroglycerine was so unstable, no sane Army would use it. But because dynamite was so stable, it immediately became the basis for new and more powerful weapons. Nobel knew this and it didn’t particularly bother him (his family fortune was founded in arms manufacturing), but it did upset a lot of other people. And when Alfred’s brother Ludvig died, he got an idea of just how much it bothered other folks. A French newspaper thought that it was Alfred that died, and took the opportunity to write one of the most scathing obituaries ever seen. The nicest thing that they called him was a “merchant of death”. Alfred was mortified.

He decided to redeem his family name. And, since science had gotten him into this predicament, he decided that science would get him out of it. He established the Nobel Prize, which was given out every year for the most important work in physics, chemistry, literature, and (in a deliberately ironic twist) peace. (Later groups would add an economics prize.) The Nobel Prize has become the gold standard of work and worth in the sciences and continues to this very day. Evey year on his birthday, the prizes are awarded in the name and memory of the most famous chemist ever to live.

If you’d like to learn more about this year’s winners in chemistry, then head over to:
http://www.scientificamerican.com/podcast/episode.cfm?id=the-2013-nobel-prize-in-chemistry-k-13-10-09

August 26 – Diamond Bright

Today’s factismal: It was once thought that things burned by releasing phlogiston.

If I were to ask you what water is made up of, odds are you’d tell me “H2O”. But did you ever stop to wonder how we know that? The answer is “Thanks to Antoine Lavoisier, who would be 270 years old today”. Known as “the father of modern chemistry”, Antoine had a quick mind and an inquisitive spirit that was willing to do things that no-one else dared. When he was born, the standard explanation for why things burn was that there was a “spirit” in them known as phlogiston that was generated heat and light as it was released; what was left over was known as calx. But Antoine wasn’t satisfied with that explanation. Why should metal gain weight when they lost phlogiston?

Antoine Lavoisier, the father of modern chemistry (Image courtesy Library of Congress)

Antoine Lavoisier, the father of modern chemistry
(Image courtesy Library of Congress)

Using a set of closed flasks, Antoine was able to show that gasses such as oxygen and nitrogen had different weights; before his work, everyone had thought that they were weightless or had the same weight. Even better, he was able to show that combining the different elements created new materials that had weights which could be found from the amount of each that was used. By changing chemistry from a purely qualitative science (“this plus this gives that”) to a quantitative one (” two parts hydrogen plus one part oxygen gives one part water”), he started chemists on the road to controlling the reactions and creating materials such as plastics, fertilizers, and light-weight alloys.

One of Antoine’s most famous experiments was also one of his most audacious. He wanted to prove that diamonds were made up of nothing but carbon. So he placed a large diamond into a flask and filled the flask with oxygen before sealing it. He then used a magnifying glass to set fire to the diamond. (Don’t try this at home!) After the diamond had finished burning, he was able to show that the resulting gas was the same that was produced when carbon was burned.

Sadly, Antoine’s contributions to science made him both famous and infamous. Because he lived in France during the time of the Revolution and because he came from an aristocratic family, he was soon tried on trumped-up charges and executed. Though the state formally pardoned him a year later, it nevertheless put an end to a career that changed the world.

If you’d like to honor Antoine’s memory, then why not go do a chemistry experiment today? The ACS has a website chock-full of stinky, smelly, chemistry fun:
http://www.acs.org/content/acs/en/education/whatischemistry/adventures-in-chemistry.html