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 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

October 25 – The Big O-No!

Today’s factismal: Lake Okeechobee is only nine feet deep on average!

When a lake is named “big water”, that’s what you expect – big water. And in one sense, Lake Okeechobee delivers. It is 36 miles long and 25 miles wide. But, despite being big enough to give a bath to all of Houston with room to spare, the lake is just nine feet deep on average and only 12 feet deep at the deepest part! And, as you might guess, that combination of wide and shallow makes for some interesting water chemistry.

The "Big Water" from space (Image courtesy NASA)

The “Big Water” from space
(Image courtesy NASA)

The lake gets its water from several rivers and runoff from the surrounding farms. As a result, it is both rich in sediment and in nutrients making it rich in growth, which is why bluegill bass and other sport fish are common there. But it is also rich in arsenic and pesticides (a legacy from early farms in the area), which is why eating the fish you catch there is a no-no. There are several efforts underway to clean up the lake and restore it to its pristine condition. If you’d like to take part, then why not join the Florida Lakewatch program?
http://lakewatch.ifas.ufl.edu/

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.

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 – Its 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-2012-web-update.html