Goodbye cobwebs! Hello shiny, new blog post!
Since last posting I started grad school and joined The Sumner Lab at UC Davis and have began a project studying interesting metabolism of microbial mats in Antarctic lakes.
A few weeks ago, some minions from the DIB Lab started a reading group to go through the introductory biology textbook Campbell biology in an effort to learn from each other’s backgrounds and broaden our knowledge base. Our group has biologists and computer scientists, and a spread of experience levels. Each week, someone presents a chapter that we’ve all read. Turns out, biology is really fun to talk about! Discussions usually go off on interesting tangents about an interesting biology facet, details of definitions of foundational concepts, or how to effectively teach concepts to students. My motivation for these posts are to share tidbits of interesting discussion and potentially be corrected on anything I might have gotten wrong here. :)
This post is about chapter 3, “Water and Life”. It is the second chapter in Unit 1: The Chemistry of Life. Taylor Reiter began the discussion by having us carefully pour water on top of coins to demonstrate cohesion and spoke about the implications of that with adhesion and surface tension (ie bugs walking on water!).
We discussed connection between hydrogen bonding and how it relates to the unique properties of water. In short, hydrogen bonding is an attraction between highly electronegative covalently bonded atom (nitrogen, oxygen, or fluorine) and a hydrogen covalently bonded to another N, O, or F. The covalently shared electrons are pulled closer to the electronegative atom, giving it a partial negative charge and the hydrogen a partial positive charge. The partial positive hydrogen will form an interaction with a partial negative N, O, or F or neighboring molecule in an aqueous solution.
A common mistake that undergraduate students make when learning about hydrogen bonding is that it can happen between all hydrogens and other molecules, but this is only the case when covalently bonded to N, O, and F, which are electronegative and have a small atomic radius.
Hydrogen bonding causes the solid form of H2O (ice) to be less dense than the liquid form! In general it makes sense that in a solid form, molecules are packed together more tightly and thus has a higher density. However, the hydrogen bonding in water organizes the molecules to have more space in between them, resulting in a less dense crystalline structure.
Specific Heat of Water
The specific heat of a substance is the amount of energy to be absorbed or lost for 1 g of the substance to change its temperature by 1°C. Conveniently for water, 1 g = 1 mL because its density is 1 g = 1 cm^3 = 1 mL. The specific heat of water is higher than substances of structural similarity and composition because of the high amount of hydrogen bonding. Water molecules stick together closer than a molecule of hydrocarbon chains because of these interactions, which results in water requiring more energy to change its temperature than a liquid composed of molecules that don’t interact as strongly.
If you’ve ever had the pleasure of living in a sweltering climate, you may have taken the drastic measure of pouring water on your head in an effort to cool down. This works thanks to evaporative cooling, which allows humans to cool down by sweating and outrun horses who would overheat. Heat produced from your body is transferred to water molecules in sweat. These molecules become excited and evaporate, taking some heat with them. As a child I remember being told that sweating cools your off and always wondered how secreting a liquid causes this – but it’s the evaporation of water, not the production of it, that does the trick!
Hydrophobic and Hydrophilic: Language is Important!
Hydrophilic molecules have polar bonds (less equal electron sharing) and readily interact with water molecules through partial negative and partial positive interactions and thus are attracted to water and each other. Hydrophobic molecules are nonpolar molecules (more equal electron sharing) and don’t have attractive forces with water, polar molecules, or each other.
Karen Word pointed out a common misconception about hydrophobic interactions held by students first learning this material. Hydrophilic molecules are attracted to each other, but hydrophobic molecules are not attracted to each other. However, a common misconception is that both types of molecules are attracted to other molecules of the same type. Hydrophobic molecules are pushed out of solution by the stronger interactions of hydrophilic molecules. Of course Van der Waals forces exist and cause some dipole among nonpolar molecules, but these are much, much weaker than other dipole based intermolecular interactions. Karen also mentioned that she combats this misconception by using careful language and repeating this conception multiple times when teaching.
That is all for today! The next chapter is still in chemistry land, where I feel comfortable because of my background. Subsequent posts will likely have less explanation and more open ended questions.