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c10 - 93 Chapter 10 AQUEOUS SOLUTIONS AND REACTIONS...

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Unformatted text preview: 93 Chapter 10 AQUEOUS SOLUTIONS AND REACTIONS: Background ©2004, 2008, 2011 Mark E. Nob/e At this time, we will start into coverage of several specific types of reactions. We already discussed one category, combustion reactions, in Chapter 6. There are zillions of reactions, and chemists categorize these in different ways. Different instructors cover different categories of reactions; I may not cover all of those and your instructor may add some. For now, I am going to cover aqueous chemistry. This Chapter begins a sequence of six Chapters dealing with different kinds of reactions in aqueous solution. We will devote additional Chapters down the road to even more aspects on this broad topic. Why such an emphasis on aqueous? It's because of the absolute, sheer, total importance. We live in a water world. There is no escaping this. That is how Nature developed the planet and that is how Nature developed your cells. The importance of aqueous chemistry to you cannot be overstated. Your cells do aqueous reactions. Your environment does aqueous reactions. You yourself were conceived in an aqueous broth, a vestige of the species origin. These things are part of your world. Whether you know it or not. Of the many kinds of reactions which are common in water, we will cover four types for now. These are in the following Chapters. We can't start them yet. We must first develop the picture of the water world a bit more in order to understand what is going on. 10.1 Water basics We grow up in a world surrounded by water as a normal experience of human life. We drink water. We drink aqueous beverages. We eat many foods which are predominantly water. Our own bodies are predominantly water. To us, water is a normal everyday liquid. But there is an irony here: the simple fact is that WATER IS NOT NORMAL AT ALL in the chemical and physical sense. Water, in fact, is the exception and not the rule for liquids in general. Water is weird. Its properties are strange. Compared to all other compounds which are close to water in size or mass, only water is a liquid at normal Earth conditions. All the others are gases. Why is one compound a gas and another compound is a liquid under similar conditions? This is a very interesting question, but we're not ready for the details yet. I'll tell you this much now: there is an inherent drive in all of Nature to favor the gas phase out of all three common phases (gas, liquid, solid). This applies to everything. That drive is part of "entropy" but entropy is bigger than this and I'm not going there right now. Simply for now, if a compound is NOT a gas, then something else is going on. In network compounds, this something else is chemical bonds and these bonds keep the atoms closely connected to one another. In molecular compounds, this something else is called "intermolecular forces". Intermolecular forces are the ways by which molecules interact with each other in their immediate surroundings. This can lead to molecules actually clinging to each other; this is how a liquid phase and a solid phase overcome Nature's preference for gas phase. This is all we need for now. We will get into more details on the battle and the balance of entropy and intermolecular forces beginning in Chapter 34. Notice that I said water is a liquid at normal Earth conditions. If the Earth were much hotter and/or if the amount of water on the planet was very small, then all the water would be in the gas phase. Or if the Earth was much colder, then the water would be mostly ice. Under either set of conditions, there would be no rivers, no oceans, etc. to help shape the planet. No water liquid for life as we know it. These more extreme conditions do apply to other planets, moons and other celestial bodies. I don't just mean our Solar System. Think for a moment. There are hundreds of billions of stars in our own Milky Way galaxy. Billions of those are estimated to have rocky planets. Some of those planets will have conditions for liquid phase water and possibly for life of an Earthen type. Many will not have liquid phase water, but then they might have life based on other compounds and not of an Earthen type. Even the word "life" will have to be redefined in order to accommodate these. Sound strange? Many strange things wait to be discovered. There's a whole universe out there! Curiously, you don't even have to leave Earth to discover new organisms. Much of our own planet remains unexplored. There could be more living creatures waiting to be discovered in the ocean depths and cave depths of Earth than there are in the remainder of our Solar System. Learn to question. Learn to doubt. Then discover. You are able to be here because Earth has suitable conditions for liquid phase water. Importantly, water is liquid phase under these conditions because it has strange and unusual properties. There is one thing that water can do better than just about any other solvent at these conditions: it can dissolve many ionic compounds. Water can also dissolve many covalent compounds, too, but so can many other solvents. Why is this important? The ability to dissolve ionics opens up entirely different worlds of chemistry. On the one hand, water can dissolve many ionics, but not all. The oceans contain massive 94 Chapter 10: Aqueous Solutions and Reactions: Background amounts of dissolved compounds, mostly simple ionics. On the other hand, rocks are also primarily ionic compounds and these don't dissolve readily in water. Much of this is important for shaping Earth. It's also important for shaping biology. The chemistry of life is inextricably associated with the ability of water to dissolve some ionics but not others. Like the sea, your blood is a soup of dissolved ions. Just ask Hemo the Magnificent. Yet, your bones and teeth are also ionic compounds, although these don't dissolve. There are actually two phenomena which are critically important to the total aqueous picture. The two phenomena are solubility and dissociation. Solubility is the more general of the two: you can have solubility without dissociation, but you cannot have dissociation without solubility. Both ofthese can apply to solutes which are ionic compounds and to solutes which are covalent compounds. 50 don't think that we're leaving out covalents. These are still very important to the total picture. And don't think that solubility and dissociation are easy to do, because they are not. This is where water's weirdness really comes to light: the ability to dissolve and to dissociate many different kinds of compounds. For our present purposes, solubility refers to whether a compound does or does not dissolve to a significant extent. The underlined part is very important to understanding solubility, and I will explain this further in the next Chapter. For now, I give two simple examples. You know sodium chloride, NaCI, as the primary component of table salt. You know this dissolves very well in water. We say that sodium chloride is "soluble". You may not know barium sulfate, BaSO4, from experience, but we did use it in the last Chapter as the product of an aqueous reaction and we said at the time that it was not soluble and it formed a white powder. We can now state that barium sulfate does not dissolve to a significant extent and we therefore consider it "insoluble". Actually, some of you may know BaSO4 from experience. I mentioned in Chapter 2 that some barium compounds are used for taking GI (gastro-intestinal) X-rays. Barium sulfate is the so-called "barium swallow". They slurry it in water, add some flavoring and who knows what else, and tell you to drink it. It goes into your digestive tract and it blocks the X-rays (it's "radiopaque") when they shine them on you. Then it squirts out your navel. (OK, OK. It doesn't squirt out your navel. Ijust made that up. I'm seeing if you're paying attention.) The second phenomenon to consider is dissociation. Dissociation is a general term in chemistry and it refers to the break—up of one chemical unit into two or more others. Our concern here is ionic dissociation: the process of separating a solute into individual ions in solution. I'll illustrate this with sodium chloride again. Remember that sodium chloride is an ionic network compound with ionic bonds between sodium cations and chloride anions, extending over it 5 KO, three dimensions. When it dissolves in water, you get separated $9 Na+ cations surrounded by water molecules, and you get separated Cl“ anions surrounded by water molecules. I've drawn a bit of this at left, where the unlabelled molecules are H20. (In @ reality, there can be hundreds or thousands more HZO molecules fig) than there are ions. Also, things are in closer contact than what I've drawn. I mention these points but I'm keeping the picture simple here.) The reason that you get separated ions is that the - water molecules aren't just innocently surrounding the ions: (962) they're deliberately interacting with the ions. Before I describe this further, I want to introduce some terms. In general, for any solute of any kind dissolving in any solvent of any kind, the solvent molecules surround the solute and interact with it in some manner. This process is called "solvation". We say the solute particles are "solvated". Solvation and solvated are general terms for any solvent. In the specific case with water as the solvent, the term "hydration" means the same as solvation. We say that the solute particles in water are "hydrated". Historically, some of the early studies of ions involved electricity. This related to the fact that water solutions with ionic solutes can conduct electricity, whereas pure water does not significantly conduct electricity. One term that arose from this historical connection is still in use: "electrolyte". Compounds which dissolved in water and allowed the solution to conduct electricity (e.g., table salt, NaCI) were called electrolytes. Compounds which dissolved in water and did not give rise to significant conductivity (e.g., table sugar, (:12sz011) were called "nonelectrolytes". Many people equate electricity with some electrons flowing through a metal conductor but, in an ionic solution, it's the ions doing the electricity, not electrons. Chapter 10: Aqueous Solutions and Reactions: Background 95 Electricity is really about charge flow, not just electron flow. Since ions have charges, too, then you can conduct electricity by letting ions move around: cations go one way and anions go the other way. That's electricity. That's how your nerves send electrical signals: it's how you breathe, it's how you eat, it's how your heart beats, it's how you move. Electricity by ion flow, not by electron flow. These things are part of you. Whether you know it or not. Since it's really ions that allow the solution to conduct electricity, an electrolyte is ANY compound which dissolves in water to give ions. A nonelectrolyte is any compound which dissolves in water but does not give ions. Watch the wording! Many ionic compounds are good electrolytes, but so also are many covalent compounds. Yes, some covalent compounds can dissolve in water and can form ions, although they have no ions themselves when pure. It's the water that causes this. This is part of the weirdness of water and it's part of what makes water so special. OK, you need to be aware of where we're at so you can see where we're going. I've just set the scene for some of the most important aspects of water: dissolving and dissociating a variety of ionic and covalent compounds. Let me recap this. > Water can dissolve some ionic compounds and separate them into individual ions. Dissociation is not a simple task. In order to separate the ions, water must overcome the strong ionic bonding between the cations and anions within the ionic network. > Water can dissolve some covalent compounds; water can separate some (not all) of these to form ions. Again, dissociation is not a simple task. Remember that a covalent compound has no ions to begin with. It only has covalent bonds. In order for dissociation to happen, water must break a covalent bond in the solute molecule. In both cases, water overcomes chemical bonding (ionic or covalent) in order to form the separated ions. But chemical bonds are very strong. SO HOW CAN THIS BE? Water is not a compound we would normally consider to react so strongly with something. Well, like I said above, WATER IS WEIRD. 10.2 Why is water weird? Water's weirdness stems from the interactions between the water molecules and the solute particles. In all honesty, these interactions also occur with some other solvents, but what makes water different is that the interactions are so strong. These interactions are part of intermolecular forces which I mentioned early in the Chapter. We don't need everything about those right now and we can't do them all right now anyway. So I will limit to the essentials. Although other aspects do contribute, we are going to focus only on "polarity" for now. A bit of background first. Before doing water, I'm going to do a simpler example with just a diatomic molecule. I'll do hydrogen fluoride, HF, since we did it in Chapter 3 as our first example of a covalently bonded molecule. We need to talk about covalent bonds again. In Chapter 3, we described covalent bonds as the sharing of electrons by two or more GE) atoms. For hydrogen fluoride, we said there were two electrons shared between the H and the F. Now I will expand on that simple picture. Yes, a covalent bond involves shared electrons, but the fact is that the sharing is usually not 50—50. In HF, the sharing is very uneven with F taking more of the share; that leaves H short in the share. So it's not 50—50. What is the share? The actual numerical distribution doesn't matter at this point. It only matters that the share is not even. MOST COVALENT BONDS DO NOT INVOLVE AN EVEN SHARE. This has serious consequences. Remember that electrons carry negative charges. If the bond electrons are not evenly distributed over the bond then the charges overall are not evenly distributed either. For hydrogen fluoride, the end of the molecule towards F has more than 50—50 of the bond share. This makes the F end a bit negative overall. The end of the molecule towards H has less than 50—50 of the bond share. This makes the H end a bit positive overall, because now there's not enough electron charge to offset the hydrogen's nuclear charge. This uneven distribution of charge is called "polarity". The adjective is "polar". These terms can be applied to a covalent bond or, in general, to an entire molecule. We say that HF has a polar covalent bond. The two ends of the bond carry unbalanced charge. The actual value of the charge is not important for our purposes, but typically the charges due to a bond polarity are some fraction of ii. We call these "partial charges" because they are _n_o_t a full charge. WARNING! Don't confuse these with ion charges: those are completely different. Remember the distinctions. 96 Chapter 10: Aqueous Solutions and Reactions: Background Ion charges are due to different numbers of electrons and protons in a specific chemical unit. The values of the charges are always a d: whole number. Partial charges are due to uneven electron distribution in a bond or over a whole molecule. These are typically some fraction of :tl per bond. Ion charges are symbolized by 1+, 2+, 1—, 2—, etc. Partial charges are a fractional charge and are symbolized using a lowercase Greek delta, 6. The actual amount of the partial charge can often be measured or calculated, but that is not important for us right now. For HF, we symbolize the partial charge of the H end as 6+ and the partial charge of the F end as 6—, as shown on the 6+ H—F 5— left. There's another symbolism which uses an arrow with a positive tail. The positive tail lies at the 6+ end and the arrow points to the 6— end, as shown on the right. The two pictures are equal; you can use 6+/5— symbols or you can use the arrow H_F to represent the same thing. I'm going to refer back to these illustrations in a later Chapter, "_’ so you need to flag them: write "HF flag" in the margin. As I mentioned, polarity can refer to a covalent bond or to the whole molecule. We can say that hydrogen fluoride has a polar covalent bond and we can say that hydrogen fluoride is a polar molecule. I'm only using HF as an easy introduction to polarity. Now, let's get back to water. Each H20 molecule has two covalent bonds and both of them are polar. The electrons in each bond are not shared 50—50. The oxygen end of each bond has more than an even share; each hydrogen has less than an even share. (If you're wondering, you'll see in Chapter 25 how to know 5'05" which atom gets the bigger share of electrons. For now, I'm just telling you.) The H/ \H polarity for each bond in water is shown on the left. The two individual 6- + + bond polarities add to give the molecule an overall polarity, shown on 6 6 the right with the atoms drawn as circles. The grand result is that each 0 and every molecule of water has an uneven charge distribution. The side of the molecule 0 a with the two H's is partial positive and the side of the molecule with the O is partial negative. Flag this part also for later referral: write "H20 flag" in the margin. 5+ Thus, water is polar. Now, that's an understatement. Most compounds are polar, but what sets water aside is that it is very polar. Water as a liquid is much more polar than most other liquids. Although other intermolecular forces add to water's weirdness (as will be seen in Chapter 37), this strong polarity is the important property for our purposes presently. This polarity enables water to dissolve and to dissociate a vast assortment of ionic and covalent solutes, a feat which very few other liquids can even approach at ordinary conditions. This is what I needed to introduce. Now we shall see how all of this falls into place. Opposite charges always attract each other. Those charges can be whole ion charges or they can be the fractional, partial charges of polar things. The net result is that a water molecule's partial charges will attract ions. The 6— portion of the water molecules will attract cations. The 6+ portion of the water molecule will attract anions. The illustration on the left shows the 0 side of one water molecule interacting with a generic cation (labeled C+ 6- for cation, not to be taken for carbon). On @IIIII 6+ the right, the illustration shows the 6+ on one H interacting with a generic anion (labeled A"). A string of parallel lines (lllll) is often used to emphasize an interaction and I've included it in the illustrations here. (They're not always required and I won't always use them unless I want to for emphasis.) Notice how the ion charges and the partial charges are interacting, either with the polarity of the molecule overall or with the polarity of one bonded atom. These interactions are an essential feature of the hydration of ions. The attractions between a polar molecule and an ion can be good but they're not always great. Some are better than others, but that depends on which ions are involved. For example, a water molecule is more attracted to a 3+ cation than to a 1+ cation. Even so, the attractions are often not as strong as full chemical bonds. However, they work very well for water for two reasons. First, because water's polarity is stronger than the polarity of most other liquids, each of these attractions is stronger in water than in most other solvents. Second, it's not just one water molecule which is interacting, it's a bunch. Each dissolved ion is surrounded by a bunch of water molecules, all piling around, even in layers, trying to get in on the charge a...
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