tinywow_video_to_text_65209252.txt

In this first lesson on a chapter on cell and molecular biology here, we're gonna talk about the building blocks of living organisms It's gonna center around proteins, carbohydrates, nucleic acids, and lipids. Before we get there, we have a little bit of chemistry. We've got to introduce that's gonn...

In this first lesson on a chapter on cell and molecular biology here, we're gonna talk about the building blocks of living organisms It's gonna center around proteins, carbohydrates, nucleic acids, and lipids. Before we get there, we have a little bit of chemistry. We've got to introduce that's gonna be kind of central to some of these discussions. So we'll start talking about matter. We'll talk about bonding. We're gonna talk about intermolecular forces. Then we're gonna talk about the special role that water plays in all of biology. We'll start with matter. So what in the world is matter? Well, the technical definition is that anything that has mass and occupy space. So it has mass and volume. So pretty much everything in the world around you is matter. My big bald head is composed of matter, but the dumb ideas it produces are immaterial and are not composed of matter. And pretty much matter is gonna be made up of, you know, the hundred or so different elements on the periodic table. So now there's a fair amount more than a hundred, so but there's still some debate on some of the higher numbers of whether they really exist or not and stuff. We're not going there. But the elements on the periodic table, they can exist as substances all on their own, in which case we call them elements. So we use that word element in two different ways. We can use it when they're existing as a single substance. We'll call that an element. But also when they're part of a larger substance, we're still gonna refer to them as elements. So you gotta be a little careful on the context here. So, but you might have, you know, pure sodium metal. You might have pure chlorine gas. And chlorine gas, it turns out, is diatomic. Most of the elements will write one at a time. But seven of them come as a diatomic and you'd go over which of those seven there are in the general chemistry portion of this course. So suffice it to say though, you can have pure sodium, you can have pure chlorine. So, and the fact that this guy's diatomic does mean it has two atoms, but it's still just one element. Now, you can have substances that are made of combinations of elements more than one element and we'll call those compounds. And so you might have something like sodium chloride. Now, it turns out the elements on the periodic table, and again, you'll get this in the general chemistry portion of this course. On the left-hand side here, the bulk of the elements are metals. And then on the right-hand side, you've got some non-metals. So, and it turns out when you've got a substance like NaCl, where you've got a metal and a non-metal, sodium's over here and is a metal. Chlorine's over here and is a non-metal. That's gonna be an ionic compound. We'll get a little more into that in this lesson here in a sec. So, but as I said, we've got an ionic compound here. So, but if you make a substance of multiple non-metals, like in the case of carbon dioxide, carbon over here is a non-metal, oxygen over here is also a non-metal. So, now we get a molecular compound instead. So, and I bring this up only to say that molecular compounds can exist as individual molecules. Ionic compounds don't. So, a lot of students think that any compound can exist as a molecule. Well, it turns out only molecular compounds can exist as molecules hence the name. So, like if I said I had a NaCl molecule, well, there's no such thing. So, ionic compounds generally exist as big crystalline structures that have many of the cation and many of the anion, many of the positive ion, many of the negative ions, so many sodiums, many chlorines. Whereas carbon dioxide can exist as individual molecules, and we often can draw out lovely structures like Lewis structures for those individual molecules, or we might draw some sort of ball and stick. We represent the atoms as spheres, and so you might have a carbon bonded to oxygens or something like that. But we can't do that for ionic compounds. We can only do that for molecular compounds made up of multiple non-metals. All right, so atoms versus molecules, last point of discussion here. And so, like we saw here in the case of carbon dioxide, we have a carbon dioxide molecule, but that carbon dioxide molecules composed of actually three atoms. All right, that's about all I want to get to if matter. Let's talk a little bit about bonding. All right, so the two most prominent types of bonding here are going to be ionic bonding and covalent bonding. And ionic bonding, you've got cations, and you've got anions. Cations are positively charged ions, and anions are negatively charged ions. So, and again, we talk about most ionic compounds are going to be metals with non-metals. And the way that works is metals usually form cations. They usually lose electrons to end up with a positive charge, whereas non-metals generally gain electrons to end up with a negative charge and therefore be anions. And that's why it's often metal with non-metal. So, but in this context, in the biological context, you're also going to see oftentimes where you've got all non-metals, but just ending up with some sort of positive charge or negative charge, and sometimes it's going to be in a pH-dependent way and stuff like this, as we'll see with some of the amino acids and stuff. So, but if long if you got a cation with a positive charge bonding to anion with a negative charge, we're going to call that an ionic bond. All right, now covalent bonding on the other hand is going to be between non-metals. So, like we draw throughout the molecule of carbon dioxide, the bonds holding the carbon to the oxygens, those are going to be covalent bonds. So, and generally we break those up into a couple different classes. We talk about polar. And we talk about non-polar. So, polar covalent bonds versus non-polar covalent bonds, and sometimes these are called pure covalent bonds, and it all comes down to a difference in electronegativity. And you might recall from your general chemistry days that fluorine is the most electronegative element on the periodic table. And the closer we get to the fluorine, the more electronegative the elements. So, and that electronegativity is just a measure of how much it's pulling the bonded electrons towards it. And the more it can pull the bonded electrons towards it and atom can, the more partially negatively charge is going to become. And the more electrons are pulled away from a different atom, the more partial positive charge is going to become. And that separation of charge with partial plus and minus charges is going to be the nature of some other intermolecular forces and have some profound effects on chemical properties. So, and because you're now going to have a partial positive, partial negative, that's a partial positive pull and a partial negative pull for that bond, we refer to those as being polar. And so, the big thing is you want differences in electronegativity, like, you know, a carbon fluorine bond is a very polar bond. A carbon oxygen bond is a very polar bond. And typically, in the realm of biology, it's easier to remember which bonds are not polar. So, and it turns out if you've got two identical atoms, that's probably going to be a non-polar bond. And so, oftentimes, the most common ones we're going to see though are carbon carbon bonds. But also, it turns out carbon and hydrogen. Carbon and hydrogen are very close in electronegativity. And so, even though they're not identical atoms, it's still considered a non-polar bond. And so, oftentimes, we're going to see big non-polar regions of molecules. And we're going to associate those with finding lots of carbon, carbon, and carbon hydrogen bonds, and not having any or very few electronegative atoms like nitrogen, oxygen, fluorine, things of that sort. All right. Yeah, one last type of bonding to talk about. That's a little more rare than these guys. And it's going to be called a coordinate covalent bonding. So, coordinate covalent bonding. Now, covalent bonding is generally the sharing of electrons between non-metals. But in coordinate covalent bonding, it's still going to involve the sharing of electrons. But it's going to be just a little bit different. And the most prominent place you're going to see this show up is with transition metals. And the transition metals are right in here. And when transition metals or transition metal ions are often going to bond to what we call ligands. And the big requirement of a ligand is just has to have a lone pair. And so, in some cases, you might see, like, an Fe2 plus ions pretty common. So, and normally think, oh, I've got Fe2 plus. He's positively charged. He can bond to negatively charged ions. And that's true. And that would be ionic bonding. That's not what's going to be new here. But it turns out sometimes you're going to see Fe2 plus bonding to neutral atoms as well. And the big thing here is this neutral atom, all it needs to actually have is a lone pair of electrons. And it might use that lone pair of electrons to form a bond in that case. And so, if you take a look at, like, the center of hemoglobin, there's a heme group. So, and you've got some nitrogen atoms that are all bonded to iron 2 plus or 3 plus, depending on the oxidation state in that particular hemoglobin. And what you'll find, though, is that also that is, turns out, in heme, where a molecule of oxygen comes and binds. And one of the lone pairs is used to make a bond to that iron ion as well. And these are coordinate covalent bonds. Notice, oxygen doesn't actually have a negative charge. There's no ionic bond going on here. And so, again, we'd say that the oxygen molecule here is acting as a ligand. And so, generally, you're going to have a central metal ion, which is usually going to be one of these transition metals in the middle of the table. And you're going to have some multiple number of ligands. So, in the case of hemeoglobin, you've got this big porphyrin ring, which is going to make four bonds to the iron ion. So, but that leaves room for an oxygen molecule to actually bond in another location as well. So, we've got three main types of intermolecular forces we're going to talk about in a biological context. And when you cover this in general chemistry, we'll include a fourth, which is going to be called ion dipole forces. So, but for a biological context, we don't really need to talk about those. So, we'll leave this just a hydrogen bonding, dipole-dipole forces, limited dispersion forces. And we're going to start with dipole-dipole forces. Dipole-dipole forces are between individual molecules. So, notice intermolecular forces is forces between separate molecules. And this happens when they're both polar. So, if you got two polar molecules, they can be two of the same molecule. They can be two different molecules. But as long as they're both polar, they will have dipole-dipole forces between them. And the big driving factor of this force, these partial positive and negative charges. So, let's say you've got HCl. So, what chlorine is more electronegative than hydrogen, and as a result, chlorine's going to end up pulling the bonding electrons closer to him, making him partially negative, and that's going to leave hydrogen to be partially positive. And so, if I put another polar molecule, and in this case, I'm just going to make it another molecule of HCl. Well, once again, you're going to have a partially negative chlorine here and a partially positive hydrogen there. And what you're going to find is that the partially negative charge on this chlorine to this molecule is going to be attracted to the partially positive hydrogen on the adjacent molecule. And we're going to form a little attractive interaction there. Now, these intermolecular forces are typically much weaker than actual, either ionic or covalent bonds. So, don't think of them, you know, we might use the word hydrogen bonding, which turns out it's really much weaker than an actual like covalent bond. So, but this weak attractive force, that's what we're calling this intermolecular force. Again, much weaker than say the covalent bonds between hydrogen and chlorine within the same molecule. All right. So, and again, just do the nature of having partial positive, partial negative charges. And so, if you've got two polar molecules, whether they're two identical molecules or two different polar molecules, they'll have dipole-dipole forces between them. Now, in the case of hydrogen bonding, which is going to be the strongest of these three intermolecular forces, well, one, it's always going to involve hydrogen, hence the name, but not all hydrants are capable of actually carrying out. And so, it turns out one of the three bonds you have to have is you have to have either a fluorine bonded to a hydrogen, an oxygen bonded to a hydrogen, or a nitrogen bonded to a hydrogen within a molecule. So, it's not just enough to have hydrogen. You have to have hydrogen in a really, really polar bond, bonded to either fluorine, the most electronegative element, oxygen, the second most electronegative element, and nitrogen, which might be third or fourth depending on who you talk to. So, but in one of these three very polar bonds, that's when you're going to have, you know, it's still kind of like dipole-dipole forces, but it's so much stronger than all the other dipole-dipole forces that they give it a special name. And so, you're supposed to recognize in biological molecules when you see hydrogen bonded to one of these three, and I say one of these three. Well, the truth is, you're never going to see an FH bond, because the only molecule that actually has an FH bond is hydrofluoric acid HF. That's it. In a biological context, you're only really going to see OH and NH bonds. So, and that's when they're going to participate in hydrogen bonding. So, things like alcohols, which have an OH bond. So, if you look at like DNA, you're going to have ribosis being the backbone and has OHs, and there's hydrogen bonding involved with those OHs and think of the sort. So, holding the DNA bases together in the two strands of a double helix, that's hydrogen bonding going on right there. When proteins fold up in a big complex three-dimensional structures, lots of hydrogen bonding involve there as well. Not the only animal like the force involved there, but one of the major ones involved as well. So, definitely want to recognize when there's hydrogen bonding, and definitely want to know that the solvents of all of life, right, water, having OH bonds capable of lots and lots of hydrogen bonding. All right, last but not least, lung and dispersion forces. And it turns out, you learn in the general chemistry lesson that all molecules are capable of lung and dispersion forces, and the bigger they are, the more of these lung and dispersion forces they're going to have. So, and I don't want to really go into what these are and stuff too much in this lesson. You'll get that in the general chemistry lesson. So, but you should know that all molecules have these, but we really only talk about them when it's the only kind of force they have. And so, if you've got hydrogen bonding or your polar and have dipole dipole forces, you're probably not going to talk about these even though they're present. It's when you're non-polar and don't have hydrogen bonding and don't have dipole dipole forces that most of the time then we will talk about these lung and dispersion forces. And so, when you're non-polar, that's the only kind of intermolecular force you have. And it's just due to the motion of electrons causing very temporary partial positive and partial negative charges. And they're very temporary and they're very weak. And so, these lung and dispersion forces, sometimes just called London forces, sometimes called Van der Waals forces, much weaker than the corresponding hydrogen bonding and dipole dipole forces in most cases. So, now we're going to spend just a couple of minutes talking about water, the unique properties of water, and it's pivotal for life to exist. So, when they're examining other planets to find out if they might carry life, the first thing they ever look for is, does it have water? So, if it doesn't have water, they usually just rule it out right away. So, water absolutely pivotal for life to be possible. So, inside of cells, the medium inside there is water, extracellular mediums, typically water. So, when it turns out, water will dissolve a large variety of both ionic solutes as well as polar solutes. And it makes it a great medium for chemical reactions to take place. Now, typically these chemical reactions are only going to take place when water is in the liquid phase, and luckily for us, water exists as a liquid over a very wide range, all the way from zero degrees Celsius to 100 degrees Celsius. For a lot of other compounds, the range of which they exist as a liquid is often much smaller, and often it's not whether it's big or small, it's the only consideration, but also, is that a temperature that exists on that planet? It's on planet Earth. Zero degrees to 100 degrees is pretty common in a lot of parts of the Earth in a lot of different seasons. So, it makes it real convenient. So, it turns out with water, you should definitely recognize that it is capable of hydrogen bonding, and it turns out lots and lots of hydrogen bonding. So, and it turns out, that makes this very, very polar, which is why it can dissolve a lot, you know, a wide range of both ionic and polar solutes. So, also, it turns out that the strength of this, the end molecular force is involved here, also give it a rather high heat capacity. You can add a lot of heat to water, which is going to cause to expand a little bit and weaken these inner molecular forces a little bit, but because they're so strong, there's a wide range over which they can be expanded and weakened a little bit and stuff like that. And so, you can add a lot of heat. So, it turns out and not get a huge temperature change as well. So, water's got a fairly high heat capacity, again, all going back to hydrogen bonding. And, you know, lots of unique things about water and it usually is going to come down to the fact that it's got this great degree of hydrogen bonding present. Now, some vocab we've got to talk about here. So, adhesion and cohesion. So, it turns out water being very polar will bind a lot of other polar things. So, if you take a look at like, you know, a graduated cylinder. So, when you measure something in a graduated cylinder, you usually get this lovely meniscus that is curved. And that is because the glass here is actually polar itself. And so, the water molecules are highly attracted to the glass on the sides. There's white creeps up the sides forming that meniscus. That's an example of adhesion. So, water's also exhibits a fair degree of cohesion. And cohesion just means sticking to itself. And water molecules are highly attracted to each other through this network of hydrogen bonding. So, this gives water a high surface tension. If you notice, if you filled this lovely graduated cylinder, or if you filled a glass up with water, you'd find that you can actually fill it a little bit fuller than the glass will go because the water will form a bubble at the top, so to speak. And that's due to surface tension. So, you get this network of hydrogen bonding happening at the top of that water holding it together. And the surface tension has a wide variety of different biological relevances. If you had only water inside your lungs around your avioli, it would cause them to implode. And so, instead, your lung cells need to produce a surfactant to break up some of the hydrogen bonding that's going on to weaken the surface tension that's going on so that your avioli don't implode. We need them to stay inflated, so to speak. So, that's adhesion and cohesion. We also have hydrophilic and hydrophobic. Hydrophilic means water loving. Hydrophobic means water fearing. So, it's exactly that. When you have things that are relatively polar, they're going to like water because water is polar and like dissolves like, like, like, like, so to speak. And so, things that are polar tend to be hydrophilic. So, whereas hydrophobic, things that are relatively non-polar are going to be hydrophobic and hate water. So, another way to look at it is will they mix with water? Things that generally mix with water are polar things and we call them hydrophilic. Things that don't mix with water, like, say, oil. So, form a separate layer. If you put it in with water and stuff like that, we refer to those as being hydrophobic and, again, they're relatively non-polar. So, we'll find out that these are two common words used in describing protein structure. So, it turns out a lot of proteins, you know, being in the medium of the cell. A lot of them are just exposed to the medium of the cell water. And so, what they end up doing is having some of the amino acids they're composed of that are hydrophilic facing out towards the water, whereas the ones that are hydrophobic, well, they don't like water. They don't interact well with water. And so, they sequester them to the interior of the protein so that they don't have to interact with water. So, two great terms we're going to use pretty ubiquitously throughout the rest of this lesson.