Using cool carbon dioxide videos to draw attention to surface tension and talk about what causes it and how it relates to “day-to-day” labb-y things like liquid measuring and crystallography? Guilty as charged! From meniscuses to hanging drops, to rain drops, and dry ice bubble pops, it all has to do with SURFACE TENSION. So I think it’s worth many a mention!
Surface tension holds a liquid’s surface together, giving it a “shape” & it’s caused by (collectively) strong intERmolecular (between molecule) forces – partial charges “gluing” the molecules together. When these forces are between like molecules (e.g. water to water) we call them cohesion & it can lead to things like water drop shapes. When the forces are between unlike molecules (e.g. water to glass) we call it adhesion, and it can lead to things like drops sticking to glass instead of just rolling off.
What causes it? Like most things in biochemistry, it comes down to the fact that opposite charges (even partial ones) attract, and they’ll do whatever they can to be near their loved ones, even if it means excluding other molecules in the process…
Water has a really high surface tension because water’s really “sticky” but if you look at a water molecule’s structure you don’t see any charges. So what’s going on? Water’s neutral overall, but if you look at charge distribution within an individual water molecule, that’s a whole ‘nother story…
The atoms WITHIN molecules are held together by (individually) strong covalent bonds, which involve sharing of pairs electrons (the negatively-charged subatomic particles that whizz around an atoms’ positively-charged-proton-containing atomic nucleus). Those bonds are hard to break. But if there is an uneven number of protons vs. electrons or even just an uneven distribution of shared electrons, you can end up with fully or partially charged molecules or molecule parts and, since opposites attract, they get attracted to oppositely-charged things, giving you intERmolecular bonds. Unlike the covalent bonds, these are more easily unstickable. So, for example, the 2 strands of DNA in a double helix can “unzip” from each other without unchaining the letters in each strand.
But not all molecules like to stick to all other molecules, or even to copies of themselves. They play favorites. Water molecules are highly polar – they have uneven charge distribution. This is because the oxygen is highly electronegative (electron-hogging), so it pulls (negatively-charged) electrons (e⁻) away from the hydrogens, so the O becomes partly ➖ & the hydrogens partly ➕. And now the “opposites attract” thing kicks in, so water molecules orient themselves so that the H’s of water molecule face the O’s of others.
Then it becomes kinda like trying to figure out the seating arrangement at a big wedding where everyone is friends with almost everyone else and they want to sit near as many of these friends as possible. But there are also some people who “no-one likes” that thus, “by default” get grouped together in order to maximize the friends sitting together. When it comes to water, the molecular version of these “outcasts” are hydrophobic molecules like fats and oils, which are nonpolar & don’t have charges – even partial ones – to offer, so they get excluded from water networks in the “hydrophobic effect.”
In a water drop, water molecules in the drop’s INTERIOR are surrounded by other water molecules on ALL SIDES, so they have lots of potential binding partners & are attracted in all directions. And because they’re in *all* directions, these attractive forces cancel out so there’s no NET force.
BUT water molecules on drop’s surface have air on one side instead of water, so they have fewer water binding partners. As a result, they form stronger bonds with the ones they do have. And since they’re not being pulled by water in the air-facing direction, they DO have a net force. This force is at a right angle (perpendicular) to the surface & it pulls the water towards the center of the sphere.
We use this to our advantage in hanging-drop crystallography, where we try to get proteins to crystallize within drops of liquid “hanging” from glass slides (basically you set up drops of liquid containing your protein and precipitants on a glass slide and then you flip the slide over and use it as a “roof” for a well containing more precipitants and then you let water evaporate out of the drop, concentrating the protein in the drop and hopefully getting it to crystallize since there’s less water to hang out with.
And speaking of “hanging,” the reason our drops are able to “hang” from the roof and not drop is the high surface tension of water. When the attractive forces at the surface of the drop are stronger than gravity trying to pull the liquid down, the drop can stay suspended even if we flip our slide over (unless you accidentally bump the plate!)
Another way to think about it is that water molecules want to minimize their “unfullfilled” charges, which means minimizing their air exposure which means minimizing their surface area. BUT what if they’re in contact with a surface that can also fulfill those charges?
In this scenario, COHESIVE forces (attractions between like molecules) have to contend with ADHESIVE FORCES (attractions between unlike molecules) which try to get the molecule to spread out. If the surface of the water is in contact with a material it’s attracted to, that material will “pull it away” from its water friends
This is why water forms a concave MENISCUS (that “smile”) in a glass, straw, graduated cylinder, etc. The MENISCUS forms because the edges of water are adhering to container’s walls. ADHESIVE forces dominate at the RIM (because that’s where the “unlike” molecules are, whereas COHESIVE forces dominate in the CENTER (there’s only water or air to hang out with here, so the water “sucks in” to minimize air contact)
Note: when you’re measuring, measure from the BOTTOM of the meniscus! And make sure that you’re measuring with your eyes level with it – if you look at it from an angle, it can look like the bottom of the meniscus isn’t where it really is. This is called “parallax error” and it can lead to inaccurate measurements.
COHESIVE forces try to MINIMIZE surface area whereas ADHESIVE forces try to MAXIMIZE surface area
When COHESIVE forces “win” we say the liquid is “NON-WETTING” & our drop stays spherical
When ADHESIVE forces “win” we say the liquid is “WETTING” & our drop spreads out.
When GRAVITATIONAL forces “win” the drop falls
So instead of setting up our drops on paper towels, where adhesive forces are strong, we set them up on siliconized glass coverslips. Even more than it wants to avoid air, water wants to avoid hydrophobic things. NONPOLAR molecules don’t have the charges needed to “stick” to water. This is why we coat our coverslip “roofs” with silicon – it’s hydrophobic so it makes the ADHESIVE forces WEAKER. This lets COHESIVE forces dominate & the drop stays round.
Detergents lower surface tension because they disrupt the water-water interactions. So when you do your laundry the detergent can get deep into the clothes & not just hang out on the surface
And this is why we can biow bubbles with soapy water but not plain water. Bubbles are basically trapped pockets of gas. Yesterday we looked at how dry ice, the solid form of carbon dioxide, can sublimate – go directly from a solid into a gas without becoming a liquid along the way http://bit.ly/2U83SCC
So when you drop dry ice into water, the submerged solid starts subliming – and you get pockets of CO₂ gas forming – bubbles! But when those bubbles reach the surface, they burst because water has too high of surface tension to form bubbles – it’s kinda like trying to blow up a ballon made of cement. But when soap molecules get in there and break things up a big, the water molecules are able to spread out more and the bubble can expand. And the soap does more than just that….
In a soap bubble what you get is a film where you have 2 layers of soap molecules sandwiching a thin layer of water. This strange arrangement comes about because soap molecules are amphiphilic – they have a hydrophilic part (a polar head) and a hydrophobic part (a nonpolar, fatty tail). They arrange themselves so that their tails face the air (inside or outside the bubble) in order to allow as many heads as possible to hang out with the water. And this kinda insulates the water so that it doesn’t evaporate and “burst your bubble” and it limits the stickiness between the water molecules so the bubble doesn’t just “implode”
As a result, if you add dry ice to soapy water you can get bubbles that can survive for a time outside the water – until the surface tension gets overwhelmed by the vapor pressure inside the bubble caused by the gas molecules banging on the walls of the bubble as they try to escape.
Another way to decrease surface tension is by Increasing the temperature. Heat is energy & when molecules have more energy the bonds between them are weaker because the molecules have more energy to vibrate and squirm making it harder to stay connected. So washing your laundry at warm temperatures helps too.