[BACKGROUND] Good afternoon. Not a big crowd for about three hours before an exam. I'm hoping at this point everybody knows what room they're supposed to be in at five o'clock and everything else. We'll see how it goes today. Last time we finished up what we're going to do in Chapter 10. Today we're taking a look at Chapter 11, the title of which is, Solutions and Colloids. We'll start with Section 11.1, the dissolution process. What things happen when substances dissolve in other substances and how we can make predictions about whether something will dissolve in something else. Now earlier, we made a reference to a rule that comes into play especially when you're dissolving a liquid in another liquid, and that's the rule of like dissolves like. Does anybody remember what that rule has to do with? Polarity versus non-polarity. In other words, if you think back to the first day of class when I had that one little demonstration I did mixing three liquids together in the cylinder getting three layers, all that stuff part of what made that work. Well, let me just draw a picture of that. Here's a graduated cylinder, we're wound up with three liquids in there. It turns out that the liquid that wound up in the middle was simply water, and by now we know that water is what? Polar or non-polar? [OVERLAPPING] Polar, of course. Now, a lot of people just know at the top of their heads that water is polar, but you should now be able to explain to somebody why it's polar. Because a molecule of water looks like this. That is to say, it's a bent molecule with polarity vectors pointing towards the oxygen, and the sum of those polarity vectors is a vector and that's what makes it polar. Now the other things we poured in there did not dissolve in the water. The principle of like dissolves like can be interpreted as saying polar things dissolve in polar things, non-polar things dissolve in non-polar things. But in general, polar things and non-polar things don't dissolve in each other very well. The reason that's true is that we're talking about what kind of attractive forces there are between molecules. Towards the end of last class, we were talking about the different kinds of attractive forces there are between molecules, especially in either the liquid state or the solid state. We talked about the fact that different kinds of attractive forces have different strengths. Specifically, hydrogen bonds are stronger than other dipole-dipole forces, which are in turn stronger than London or Van der Waals forces. The term miscible here simply means it's soluble in each other in all proportions. Now it turns out that the first liquid we poured in during the demonstration was carbon tetrachloride. Polar or non-polar? [BACKGROUND] Non-polar. You know that because if it had been polar, it would have dissolved in the water most likely. Now because you know what a molecule of carbon tetrachloride looks like. It looks something like this, tetrahedral geometry. That Larry vectors point towards the chlorine, is because of the symmetry of the molecule called the polarity vectors cancel out, the sum is zero and that's why carbon tetrachloride is a non-polar molecule. In the first layer, we had something that looks like this when you draw out its molecular structure. This is a material commonly known as ether. When I draw it in Lewis structure form this way it's a little bit misleading because each of the big atoms, the carbons and the oxygen are all sp3 hybridized, so it's misleading to suggest that the bond angles are 90 degree angles here. But suffice will say this, even though this molecule does have an oxygen in the middle, most of this molecule is carbon atoms and hydrogen atoms and therefore this molecule was largely non-polar. It's a little bit more polar than carbon tetrachloride is down here, but for the most part, this is non-polar, as an example or as evidence of that, we can see that the ether does not particularly dissolve in the water. Then we added iodine. Here's what iodine molecules look like. Polar or non-polar? [BACKGROUND] Say again? It has to be non-polar because you have two identical iodine atoms here, the electrons are shared equally between them. Neither one is more electronegative than the other. This is non-polar. What we saw when we tried to dissolve this stuff and everything, as the iodine did in fact dissolve in the non-polar substances but not in the water. That should make sense because non-polar things that use relatively weak London or Van der Waals forces for their intermolecular attractive forces, should be able to dissolve in each other, but not in polar things which use, well the case of water, we're talking about hydrogen bonding which is a relatively strong intermolecular force of attraction. The point is if a substance is you're trying to mix together have similar intermolecular attractive forces, similar in terms of how strong they are, they should dissolve in each other. But if you're trying to mix together substances that have significantly different intermolecular attractive forces, they will be immiscible, that is, they will not mix together. Here's an analogy I use that again takes us back to elementary school days and little kids playing on the playground. Now you tell me, because it's been a long time since I've been in an elementary school, I don't know they still let kids play games like this or not. But when I was in elementary school, there was a game where one group of kids but hold hands in a circle and another group of kids would try to break through and get into the middle of the circle. Still play games like that? Don't net kids use brute force on each other anymore? It's that okay? Do they do that? Suppose the group of kids holding hands is sixth graders, and the group of kids trying to break through are first graders, who's going to win that game? [BACKGROUND] The sixth graders, of course, because they're stronger than the first graders are. Well, water isn't letting any of these other non-polar things in because the hydrogen bonds between water molecules are stronger than the London forces that any of these other things use. The hydrogen bonds of water molecule like the sixth graders, they hold hands really tight and all these other guys which are relatively weak, can't break through because they're not strong enough. Make sense? The principle of like dissolves like can be thought of in terms of polar and non-polar. But really you should think of it in terms of how strong the intermolecular forces of attraction happen to be [NOISE]. Maybe more time with this slide. Let's consider what happens when gases dissolve in liquids. We talked about the important properties of gases back in Chapter 9. It turns out that the two most important factors that dictate whether a gas is going to be soluble in the liquid. What are the pressure of the gas? What is the temperature of the solution? In general, the solubility of a gas in a liquid tends to increase as you increase the pressure of that gas above the sample of the liquid. That's something called Henry's law. But at least qualitatively, we hope this makes sense. For example, I typically have a Coca-Cola with me here in class or free plug for the guys at Atlanta. When you first open a soda can or a bottle of soda like this, you open the bottle and you hear the little wsssh come out, right? Carbonated beverages like this are basically just carbon dioxide dissolved in water and a few other things added in the case of Coca-Cola. But the point is what gives it that tangy taste when you open it up is the CO_2 that's dissolved in the water. The reason for that wsssh, when you first open the bottle. They pack it under an increased pressure of CO_2 to increase the solubility of CO_2 in the water and therefore more completely carbonate the beverage that we're talking about. Now suppose you open that bottle of soda and you drink about half it and then you put it somewhere and you forget about it. You come back a few hours later and you tried to drink it again. What do you notice? It's flat. Yeah, it is flat. Flat is the word we use to describe Henry's law has taken over. For when you have it exposed to the air, there's very little CO_2 in the air. Air is mostly nitrogen, oxygen, tiny bit of argon but very little CO_2 floating around in the air above your flat soda. What happens? Henry's law says, "If we don't keep pressing down on the CO_2 that means the CO_2 that's already there, it's going to be less soluble in it and it bubbles out of solution." If that happens for long enough, your soda that goes flat, that's Henry's law. Makes sense? The solubility of gas in a liquid increases as the temperature of the solution decreases. This may seem counter-intuitive because most of the time if you're trying to get more of something to dissolve in water or whatever your solvent is, you increase the temperature. That works pretty well for solids. Let me show you this one diagram. This is a graph of a solubility of various solid solutes in water as a function of temperature. You can see that most of these lines tend to slope in this direction, that is as you increase the temperature, the solubility of each of these things in water tends to increase. However, here's a similar diagram involving gases and as you can see here, the trend for gases is to go in exactly the opposite direction. At higher temperatures, the gas becomes less soluble in whatever the liquid happens to be. Has anybody here ever heard of thermal pollution? That sounds familiar at all. Okay. What do you know about thermal pollution? What is it? What causes it? Yeah. [inaudible] Let's talk about nuclear power plants. One thing nuclear power plants need is water to cool everything off. The water they use to cool things off never comes in contact with radioactive material. The water isn't itself radioactive. That doesn't mean there aren't problems associated with it. Let me draw you a picture. Suppose here's your nuclear power plant over here, big cooling tower, all that. Here's a stream running by and we're dumping water into the stream now, nothing wrong with the water, it's not containing any nasty chemicals. It's not radioactive or anything like that but it is hot because we're using that water to cool off some of the components in the nuclear reactors and that heats up the water. So suppose we're dumping hot water into the stream. Suppose the stream is flowing in this direction. Upstream, happy little fish with smiles on their faces. Downstream, fish are less happy. Fish are more dead. How come? There's no pollutants, there's no radioactivity. What's creating the problem? The temperature. Yeah. Hot water. The back of that graph I just showed you. As we increase the temperature of the stream, the solubility of oxygen decreases. How do fish breath? They use their gills to take oxygen from the water. If there's not enough oxygen in the water, there's going to be more dead fish. That's thermal pollution. You don't necessarily need particular chemicals to create pollution, sometimes temperature is enough. That's a function of the fact that the solubility of a gas in a liquid depends on cool temperatures preferred and also depends on the pressure of the gas above the liquid. The two most important things to bear in mind when you think about a gas dissolved in a liquid; pressure and temperature. High pressure, low temperature preferred for getting the gas to dissolve. Finally, what happens when you dissolve solids in water? Let's talk about ionic solids for now. Let me show you this little diagram. >> What this is a diagram of is what happens when a crystal of salt dissolves in water. Now we know that salt does dissolve in water. You sprinkle some salt into water, you stir it up, it disappears. [NOISE] What's going up? Well, here's the crystal of salt. We talked about crystal lattices before. In a crystal lattice of salt, we have sodium ions represented by these little gray things here, and chloride ions, which are the green things alternating back and forth. But when you dump this into water and you have water molecules starting to bump into these things. Eventually, what happens is a water molecule might knock a chloride ion or a sodium ion off the crystal. It's floating around. As these things float around, here's a sodium ion, here's a chloride ion. Notice that they tend to get surrounded by water molecules. I'll show you a different picture of that in just a moment. But eventually the whole crystal breaks down and the individual ions are surrounded by water molecules, so they can't get back together again. Now let me show you this other chloride of this. Here's a picture of a sodium ion surrounded by water molecules. Here's a picture of a chloride ion surrounded by water molecules. But one of the reason that water, which is obviously polar, is such a great solvent for something like this, is that the positive sodium ion can actually be stabilized by being attracted to the partially negative oxygen atoms in the water molecules and the chloride ion which is negative, can be stabilized by being attracted to the partially positive hydrogen atoms within the water molecules. The pictures we're showing you here, the title of this slide is hydration of ions. Hydration simply means we're surrounding each ion with a cage of water molecules and what makes the cage is the attraction between the ions and the dipoles, that is to say the positive ion or positive hydrogens here, the negative oxygens here for whichever ion happens to be being solvated by this. The more general term is solvation. This is hydration because it's water that's doing it. Concept make sense? Okay. Now, sodium chloride dissolves in water. But that doesn't mean that any compound that's made of ions dissolves in water. In fact, there are some ionic compounds that are really not all that soluble in water. What I want to share with you today is a set of rules that have been developed to help you figure out just by looking at the formula whether or not a particular compound is going to be soluble in water. The rules are called the solubility rules. They are in your textbook. They are also in the lecture notes on page 112. I'm going to walk you through the solubility rules and provide examples as we go. In fact, I'm just going to make two columns here. Compounds that are soluble in water, compounds that are insoluble in water. Now before we start showing you the solubility rules, would anybody like more time on this slide? There are seven solubility rules and it will benefit you to get some practice with these things. Although such practice can certainly wait until after today's exam. Rule number 1, Any compound that contains either sodium ions, potassium ions, or ammonium ions, is soluble in water. For example, we know that sodium chloride is soluble in water from our own experience. But if we didn't know that, we can look at this formula, NaCl and say, "aha, it contains sodium ions, therefore, it must be soluble in water." [BACKGROUND] Most people have considerably less experience with potassium bromide than they do with sodium fluoride. Nonetheless, you can look at the formula KBr and say, "oh, it has a potassium ion there." Solubility rule one says that, Anything that has a potassium ion, it should be soluble in water, and in fact it is. Are there any questions about rule number 1? Rule number 2. These are now negative ions. Any compound that contains nitrate ions, NO_3_- per chlorine ions, ClO_4_-, chloride ions, ClO_3_- or acetate ions, C_2H_3O_2_-, all of those are soluble in water. For example, silver nitrate, AgNO_3 contains nitrate ions. According to rule number 2, anything that contains nitrate ions is soluble in water, and therefore, silver nitrate must be soluble in water. Lead four acetate contains acetate ions. Anything that contains acetate ions is soluble in water. Therefore, lead four acetate must be soluble in water. In short, the goal here is to allow you to look at the formula of the compound and figure out whether it's soluble in water or not. Questions about rule number 2? Rule number 3. In general, compounds that contain helide ions, chloride, bromide, iodide ions are soluble in water, but there are a few exceptions and the exceptions are those halide compounds that contain either silver ions, mercuric ions, or lead ions. For example, magnesium chloride. Most chlorides are water-soluble, magnesium is not one of the exemptions. Therefore, magnesium chloride is soluble in water. But, silver is one of the exceptions. That's why, for example, silver bromide is not soluble in water. Another exception is lead with a plus two charge. That's why lead iodide is going to go into the insoluble in water list. Because, silver and lead are exceptions to that rule. Is this making sense so far? We'll give people time to finish with this slide before we show you the rest of the solubility rules. [BACKGROUND] Anybody need more time here? Those are the first three rules, there are four others. Solubility rule 4 has to do with sulfates, most of which are water-soluble. But there are about half a dozen exceptions, many of the exceptions involve ions in Group 2 of the periodic table. For example, zinc sulfate is water-soluble because most sulfates are, zinc is not one of the exceptions, therefore, zinc sulfate will dissolve in water. But barium is one of the exceptions to the sulfate solubility rule, that means barium sulfate for the most part is not soluble in water. Questions about rule number 4? Up to this point, we've been talking mostly about compounds that are water-soluble. Rule number 5 refers to hydroxides, OH minus, then oxides, O double minus, most of which are insoluble in water with the following exceptions, calcium, barium, and anything that falls under rule number 1. Rule number 1 takes precedence over all the other rules. For example, ferric oxide, rust, not soluble in water. But calcium oxide, also known as lime, is soluble in water because calcium is one of the exceptions to the rule that says oxide-containing compounds are insoluble in water. Questions about rule number 5? Rule number 6. Compounds that contain carbonate ions, CO_3, double minus, phosphate ions, PO_4 triple minus, sulfide ions, S double minus, or sulfite ions, SO_3 double minus, all insoluble in water with the exception of rule number 1. For example, aluminum phosphate, like most other phosphates, insoluble in water. But ammonium phosphate, NH_4_3PO_4 is soluble in water because ammonium, anything is soluble in water, according to rule number 1. That's about as close as I can come to getting both lists on the screen. Finally, rule number 7. Compounds that are insoluble in water according to rule 5 or rule 6, should be soluble in acids. Take a look at our list of compounds here that are insoluble in water and find one compound on that list that is soluble in acid. [NOISE] Yes. Rust. Okay. Rust. Ferric oxide, oxides fall under rule 5. But rule 7 says anything that's insoluble in water according to rule 5, should be soluble in acid. Rust is in fact soluble in acid. What's another compound that are insoluble in water list that dissolves in acids? Yeah. [inaudible] Yap. Phosphates insoluble in water according to rule 6, rule seven says, if it's insoluble, according to rule 6, it should dissolve in acids, so aluminum phosphate dissolves in acid. The idea here is to allow you to look at the formula and figure out whether the compound is soluble in water, not soluble in water, soluble in acid, not soluble in acid. Everything else that are insoluble in water list is also insoluble in acid because it doesn't fall under rule 5 or rule 6. Now, I brought along a demonstration today that I hope will illustrate all of these points, but I want to make sure that everybody's got these solubility rules written down first, so finish with this slide and I'll show you the demonstration. Question? What are the exemptions to determine the experiments with all the compounds? All of these things are determined experimentally. We'll talk later on in the semester about what exactly solubility means in more quantitative terms, but for a very qualitative list, these are the solubility rules. But yeah, people didn't just make this stuff up, they went into the lab and actually tried dissolving these things in water just to see what will work. Anybody need more time here? All right, here's the demonstration. As I do this demonstration, you'll probably want to refer to the list of solubility rules. Nickel (II) chloride, soluble or insoluble in water, and which rule tells you that? [NOISE] Yeah. Soluble, rule 3. According to rule 3, chlorides generally tend to be soluble in water. Nickel is not one of the exceptions to that rule, so nickel chloride should dissolve in water, and to prove it, I brought along a solution of nickel chloride dissolved in water. Turns out nickel ion is dissolved in water and makes this nice green color that you see here. Sodium carbonate, soluble or insoluble in water? What rule tells you? Yes. Soluble, rule 1. According to rule 1, anything that contains sodium ions is soluble in water. Here's a solution of sodium carbonate. We were able to dissolve it in water. The demonstration involves observing what happens. We mix these two solutions together. I think this is probably going to be easiest to see on the wood surface to give you a contrasting color. I'm going to pour a little bit of a sodium carbonate solution in here, just enough to coat the bottom. Then some of the nickel carbon or nickel chloride solution. I hope you can see the cloudiness that's forming there. That cloudiness comes from some solid substance that is forming here. In other words, even though these two guys are both soluble in water and when we mix them together, some chemical reaction takes place to form something that is not soluble in water. To reason out what's going on here's probably the easiest way to go. [BACKGROUND] When we dissolve nickel chloride in water and we get nickel ions and chloride ions floating around in the water. When we dissolve sodium ion or sodium carbonate in water we get sodium ions and carbonate ions floating around in the water. We mix those two together, we get all four of those ions floating around in water. Is there some combination of a positive ion and a negative ion pair that would create a water-insoluble compound? If so, what is that compound? Nickel(ii) carbonate? Nickel carbonate. Yep. Nickel(ii) carbonate is absolutely correct. Most carbonates are insoluble in water. Nickel carbonate is insoluble in water. So this cloudiness that we're seeing here is solid nickel carbonate that is precipitating out of solution once we mix all these ions together. Now the byproduct here, of course, is sodium chloride. In fact, if we wanted to write a balanced equation here, we put a two in front of sodium chloride. But the main point is the sodium ions and the chloride ions, don't much do anything. They just float around in solution and nothing much happens. The sodium ions and the chloride ions here are sometimes referred to as spectator ions. Because like spectators at a football game, they're not actually participating. They're just sitting around watching what the other ions are doing. Okay, they don't really have eyes for spectating in. There was a lab experiment earlier this semester that asks you to write net ionic equations for reactions that were taking place , that sound familiar? All we mean by net ionic equation is an equation like this one, except we don't bother penciling in the spectator ions. So this would be a perfectly good way to write the entire balanced equation here. But if all we want is the net ionic equation, that's just the nickel ions reacting with the carbonate ions. The net ionic equation here would just be to say aqueous nickel(ii) ions react with aqueous carbonate ions to form solid nickel(ii) carbon. There we are. Now, nickel carbonate is obviously not soluble in water, but what should dissolve in? Acid? Rule Number 7 says, "Anything that's inside a little water, according to rule Number 6, like nickel carbonate should dissolve in acid. Well, I just happened to have brought one. Some six molar hydrochloric acid. Let's just add a few drops of six molar HCl here and we'll see what happens. Would you agree with me that it's less cloudy than it used to be? [BACKGROUND] Point is, by adding acid, we made the nickel carbonate precipitate go away. Specifically, what happened there as we add HCl, the chloride ions combined with nickel, we get our nickel chloride back. We know that soluble in water, the hydrogen ions combine with the carbonate ions to give us H_2CO_3, which is carbonic acid. But that rapidly decomposes to form water and carbon dioxide. You have to look very close to see the few little bubbles of carbon dioxide that are forming there, but you'll have to trust me, there are some bubbles in there. The point is hopefully the solubility rules can help you to make sense of situations like this. Sound reasonable? As you prepare for exam Number 3, which is four weeks from today, one thing you will want to be working on is learning the solubility rules. [NOISE] The title of Section 11.2 in your textbook is electrolytes. Who knows what electrolytes are? Who knows why they are called electrolytes? You either define or give me an example of that term. [BACKGROUND] You have heard the word before. What's it mean? Is that a hand up? No. Okay. Yes? [inaudible] You're on the right track. The reason electrolytes are called electrolytes, they are mostly salts, that is to say, ionic compounds. Electrolytes are simply compounds which when you dissolve them in water, create a solution that conducts electricity. Now there's an example of this that I'm sure it goes back to earlier days when you would go swimming grabs in a swimming pool, perhaps at the beach. Let's suppose you're swimming in the ocean or in a swimming pool and a thunderstorm crops up. What do the lifeguards tell you to do? Get out of the water and stay out of the water. Why? Yes. Because the ocean is salt water, so it'll give out electricity? Yeah. The ocean is a great big electrolyte solution. It is largely salt dissolved in water. If you're swimming in the ocean, both the lightning hits, it doesn't have to hit you, it just hits somewhere near you. The electricity gets conducted because salt, sodium chloride, whatever else is dissolved in there, is an electrolyte. That's why they tell you to get out. Even in a swimming pool. The chemicals they treat the water with in the swimming pool are ionic. Swimming pool water conducts electricity. By contrast, non-electrolytes are compounds whose aqueous solutions do not conduct electricity. The point is when you dissolve an ionic compound in water, it's the presence of the positive and negative ions in the water that are free to move around, that allows that aqueous solution to conduct electricity. But suppose you dissolve something like sugar in water. Sugar is made of molecules which are neutral. When a crystal of sugar breaks down, the individual sugar molecules becomes surrounded by water molecules. However, the individual sugar molecules are neutral, and if the only particles you have floating around in the water are neutral, then that's not an electrolyte. Then there are some compounds called weak electrolytes. We will talk later on this semester about weak acids and weak bases, I think we already introduced that concept, we will say more about it later on. Suffice it to say this, a weak acid or a weak base when you dissolve it in water, creates a few ions, but not very many. Those things are called weak electrolytes because their solutions can conduct electricity a little bit, because they have a little bit of ions floating around in there. But compared to a strong electrolyte like sodium chloride, they don't conduct electricity as well as solutions of sodium chloride or other strong electrolytes do. There's a demonstration that goes with this too, but I anticipated that this is what today's crowd would look like. I think I'm going to save that demonstration for Tuesday when maybe we have a little bit better attendance. [NOISE] I'll try to do that for you on Tuesday. Do the words make sense? Does anybody need more time with this slide? A lot of the pictures and concepts we've shown you up to this point are at various places in Chapter 11 in your textbook. The general concept of solubility is discussed in section 11.3, for example. Then they talk about Henry's law and things like that. Then we get to section 11.4, the title of which is Colligative Properties. Let me just define what this means and give you some examples, and we will wrap up for the day. We're going to let you go a little bit early today so that those of you who were nice enough to show up in the first place can have a little bit of extra time to review for your exam. What do we mean by colligative properties of solutions? Well, what this refers to is, any property of a solution that depends only on how many solute or solvent particles you have floating around in the solution, but not on what those things are. Now by contrast, being an electrolyte would not fall into this category, because obviously whether something is a strong electrolyte or a weak electrolyte or a non-electrolyte depends on what the solute, the stuff that's actually dissolving in the water, actually is. These are properties that do not depend on what the substance is, just how many particles, those solute and solvent you have. There are four examples of colligative properties, and I'm willing to bet you're familiar with at least one of them. Freezing point depression refers to the fact that when you dissolve something in water or dissolve something in any other solvent, the freezing point of the solution is lower than the freezing point of the pure solvent. Now I'm confident you're familiar with that one, here's why. A few months from now when it gets to the winter time, and fortunately for us, the forecast is for a relatively mild winter, hoping that comes true, but there have already been parts of the country that have been hit by feet of snow already in the last month or so. What happens when people get hit with feet or even inches of snow? They start throwing salt around to try it to melt the snow. Well, that's the common misconception. Salt doesn't actually melt snow. What salt does is dissolve in the snow and create a solution which has a lower freezing point than it would have had previously. We know that water freezes at zero Celsius. Suppose you throw enough salt into snow to create a solution whose freezing point is, let's say, minus 5 Celsius. Well, if the temperature that day is minus 2 Celsius, then that's going to melt. That's what happens, and then hopefully you can drive your car on the roads and hopefully not worry about it too much. Adding a solute also has an impact on the boiling point of the solution, it tends to raise the boiling point above what it would be. Now, this is perhaps not the best example of that, but some people when they put a pot of water on the stove with the idea that they're going to cook spaghetti or something like that, they can throw a little bit of salt into the boiling water. Maybe that's just for taste or enhancing the flavor. But adding a little bit of salt to the water actually raises the boiling point of the water, which means that if you've raised the boiling point to, let's say, 101 degrees instead of 100 degrees, you might cook the spaghetti a little faster. It has that impact as well. We said before that the definition of boiling point is the temperature at which a substance's vapor pressure is equal to whatever the atmospheric pressure that day would happen to be. If the boiling point is going up, it may not be too surprising to find that the vapor pressure is going down. If you dissolve something in water or some other solvent, the vapor pressure of the solution is going to be less than what it was before you dissolved whenever that solute was in it. We're not going to talk about these three very much, but we are going to talk about the last one. Now, I don't know if you've heard of osmotic pressure before, but I suspect you have heard of a related term. What is osmosis? [NOISE] Can anyone define or give me an example of osmosis? The word sounds familiar, right? Yeah. The movement of water. Movement of water from where to where, through what? From a hypertonic to a hypertonic solution between- Go ahead, I'm liking what I'm hearing so far. -a semipermeable membrane. Okay. The magic words I was looking for there were, through a membrane. Let me just find this one picture. Now, two words that get pathed into your head in high school, semipermeable membrane. What does semipermeable mean? Yeah. Not everything passes through. Some stuff gets through, other stuff doesn't get through. Usually what makes the difference is the size of the particles that we're talking about. Semipermeable membranes include things like cell membranes, cell walls, stuff like that. Water molecules, which are tiny, can generally get through. Bigger molecules like whatever these green things are, don't get through. What this picture shows, suppose you have a concentrated solution of something on one side of your membrane, and a more dilute solution on the other side. There's more water per unit volume in the dilute solution than there is in the concentrated solution. The net flow of water molecules is going to be in this direction. If you just let everything flow, you will eventually reach this state where you have two different heights of the columns, but the concentrations should now be closer to the same. That's osmosis. That sound familiar? Okay. The concept of osmotic pressure is related to osmosis, but I'm going to let you go at this point and we'll talk more next time about how it's related to osmosis and how this ties in with the concept of colligative properties and other things we've seen before. Good luck to everybody today. Make sure you go to the right room, make sure you bring your calculator, all the important stuff. [BACKGROUND] [NOISE] [BACKGROUND] Here's the problem. Even though you might put distill water the pool [NOISE], if you jump into it, there's going to be enough coming out of your sweat to make it not be distilled water almost immediately. If you can somehow manage to keep it distilled, no ions whatsoever. Yeah. If you say and I hope to be able to show you that in the demonstration when I doing hands-on, [BACKGROUND] go to pushing force. Okay. If you look up here, this is the energy diagram for arginine. Now all we did here was to pencil in the values and the point is S orbitals must have zero for M. The P orbitals come in groups of three, the values of minus one zero. This question says, "How many electrons have a value of zero for M?" Look for each box, that has a zero. Two electrons here, two electrons here, two electrons in the middlebox there. Two electrons there, two electrons and complex there. That's a total of 10, or 10 electrons. [inaudible]. Well, it's most helpful if you draw out the Lewis structure and then you have a good idea of what it looks like in three dimensions. Because once you know what it looks like in three dimensions, then you can think about electronegativity and think about which direction that the polarity vectors go. The best way to figure out whether the molecule is polar or non-polar is to add up the polarity vectors, get the vector sum up to zero. It's not homework [inaudible]. But none of that has any meaning if you don't understand things like electronegativity, molecular geometry, Lewis structure. Which is why you have to go through all that stuff. It's hard to just look at the formula and figure out what it's for. [OVERLAPPING] Yeah, it's very difficult. Okay? Thank you. [BACKGROUND] [OVERLAPPING] You're welcome. We're trying to come up with this. We want to [OVERLAPPING]. There had been a lot of NO2 on the list. Two moles of NO2, that's on the right side. And it's going to cut this in half and analyzed reversing amounts with ion changes there. Now when we do that, that also puts one mole of NO on the right side. We'll have three moles of NO on the right side. And some just add this equation as is. Because when we add two more moles of NO to the one mole of NO then we'd have to reverse it the other way, we get three. That's why we just take this number as it is. Then here's the other number that we're going to go into that right now. Okay. Okay. Thank you. Alright. [BACKGROUND] What I ended up doing is I subtracted. That's my [inaudible] of two and we got positive 238. We're going to be negative as it has been good. You're not supposed to do that. You're supposed to do this thing in bonds broken minus bonds formed. In other words, see this triple bond right here. You don't see it over there. So that means a triple bond was broken. You see the water here. The water has two OH bonds, but there are no OH bonds over here because the oxygen isn't connected to any hydrogens. So you have to break both of those OH bonds. Scroll up to say that's what's come up with that. Okay, here's, my arithmetic. Point is 812 comes from the triple bond, 460 comes with the two OH bonds, those are the bonds we can break. Scroll back then. But over here, we form a CO double bond. In four bonds we get energy out specifically similar to 95 kilojoules. We also form the carbon-carbon single-bond reflex, the carbon-carbon triple bonds and it's 347 more. And there's a total of six carbon-hydrogen bonds here as well. There's only four of the starting material so they gain two of those. That's where all those negative [inaudible] come from. Bonds broken minus bonds formed give you the answer. Thank you. Would there be an easier way to figure out these three without memorizing the chart? Right. The way I do it is first draw the Lewis structure. Yeah, I drew it here. Then what I do is figure out what the hybrid orbitals are. In other words, when it says what's the hybridization? Yeah. Because when I look at something like this I say, "Okay, there's five things, it must be sb3d hybridized. Now when sb3d tells me that the basic geometry is trigonal bipyramidal, which means the bond angles are 90 degrees and 120 degrees. We made a 100 and that would mean 100 point groups? That's just one of the basic angles of the trigonal bipyramidal geometry. I'll show you something from electrodes. Thank you. Okay. This five geometry is down here. And here are the bond angles that are associated with them, right? This is two things, three things, four things, five things, six things. Right. Okay. So for example, if I have five things, I know the basic bond angles are 90 and 120 and other basic geometry is trigonal by gram. These other geometries all derived from that. Yeah, that's all we see in the chart. Okay, well, this is a small chart. Some people like to memorize the big chart. Yeah, we'll take those unreasonably charted out. It's up to you. What I do is instead of worrying about this, I just know the five basic ones then I derive all the other ones in my head, by just replacing things with lone pairs, right? Some people prefer to memorize that. Okay, that's fine. Since starts five things, you're your own aqua deck. Yeah. I can do that on my fingers. I hold up five fingers. Yeah. Now you see it in 3D, okay. Alright, so this NCL3 you draw it out, you got four things that tells you it's sb3 hybridized. Sb3 is tetrahedral with bond angles of 109 degrees, that's why the answer is 109. Okay? And then same thing with this? Okay, draw this one out. You got six things. Two of them are lone pairs, but yeah, don't forget the lone pairs. Other words you have the four fluorines, but there's also two lone pairs, right [OVERLAPPING] that's six things. Okay? The point is six things [OVERLAPPING] after zero. I don't count to z9 itself, no, but any lone pairs that are on it, you do count. Yeah. Okay. Now you've got that. You're not counting that, I thought you count that. Now you're like you only count the lone pairs. [OVERLAPPING] You only count the lone pairs [OVERLAPPING]. All right. Okay. The point is six things would be octahedral 90-degree angles. If two of them are lone pairs, then it's square planar. That's why the answer there is square planar. So I would choose C. C? Okay. So what do we have to memorize, figured out that we had to memorize a chart? Well, okay, I'm going to let you decide how much of this piece of paper you want to memorize. Sure. Okay. What I knew is I mostly operate from the five basic geometries that derive the others based on those. Page 81. Okay, and that's page 8? [OVERLAPPING]. Page 81. Alright. Thank you very much. [BACKGROUND] right now. I will see you in five. Somebody live along student. Thank you. Hi there. So I hadn't pushed to adequately. [inaudible] [OVERLAPPING] I really enjoyed it. Well, that's your decision. Obviously, you have to juggle the rest of your life along with this course. Okay. Do you plan to continue to attend your labs or anything like that? I just don't know. That's my question [OVERLAPPING]. As an auditor, you're going to get an L for listener at the end of the semester, a better one. But one thing that might be worth doing, especially since we're more than halfway through the semester at this point, you'd go into lab. Because if you complete the lab, get a grade for Lab at the end of the semester, then we'll have that grade on a file for you. That way I can take the course again at some future time, you might not necessarily need to repeat the lab as we already have a lab grade on file for you. So you can do that. Thank you very much. Okay. Sure. Have a great day. You too. Yes. I have a question. Yes? Is there any way I can take a picture of that page. I do not have my [inaudible]. You may take a picture. Thank you very much. In fact, actually, I'll do you one better than that. Since I have my textbook, I'm going to show you where the corresponding picture is in your textbook. I only have the app for OpenStax. Well, that's okay. All right, in other words, here's the picture of the five basic geometries. Then here's the picture, it's more like the chart that I'm showing you on page 81. If you want to take a picture of that, that's fine or you want to take a picture of page 81, that's fine too. Okay. Here you go. Thank you. That's page whatever it is. It's hard to find pages in this book. Yeah, I've come to the realization that they might have done a second printing and changed some of the page numbers, which is really annoying. Yeah. I would think you really see the numbers in there. That's why I was like [inaudible] All right. Thank you very much. [NOISE]. This beautiful [inaudible] considered [inaudible]. Yes. [BACKGROUND] The exam is today right? I'm assuming that [inaudible]. Okay, what section number do you want me to [inaudible]. I believe, twenty two [inaudible]. If you are in my 2 o'clock lecture, it's a number in the nineties. If you are in my 3:30 lecture, it's a number in the seventies. Yes. The Excellent school. I can tell. I'm in the 3:30 one. Okay. So that way your lab section number should be somewhere in the seventies. [BACKGROUND]. All right. [OVERLAPPING] I'm going to suggest that you go to brown lab. Did you take the first [inaudible]? Yes. What [inaudible] went for that. I was in [inaudible]. Okay. [inaudible].
chem103-080-20191024-140000.mp4
From Dana Chatellier October 24, 2019
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