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You might expect that by this time in the history of modern civilization science would have avoided some of the confusion inherent in human languages, but as with any other language, words have entered science at different times from different places in different contexts. You'll find that there are sometimes different names for the same chemical, and sometimes different chemicals are called by the same name. It can get really complicated for organic compounds, those that contain carbon, but for inorganics at least, there is a standard naming convention that's relatively easy to understand. Put another way, a standard naming convention, or nomenclature, enhances the fidelity of the chemistry meme-plex. If you take the time to study this nomenclature, you'll have a much easier time reading the rest of the book.
The first thing to learn about inorganic nomenclature is that each compound has a first name and a last name. You probably already know the chemical name for common table salt, sodium chloride, NaCl. Silicon dioxide, SiO2, is the standard name for silica, the compound which makes up the mineral, quartz. Salt and silica are the old names, given before anyone knew what they were made of. The standard name, though, gives you a lot of information about the compound, which is why they standardized names in the first place.
The first name, under this system, denotes the cation, pronounced "cat-I-on," which carries a positive charge. As a mnemonic device, you can imagine that the "t" is a "+": ca+ion.The second name, the anion, "an-I-on," carries a negative charge. The key to understanding inorganic formulas is the old adage, "opposites attract." In any given formula, the number of positive charges will equal, or balance, the number of negative charges.
Take sodium carbonate, for example, which is composed of sodium cations and carbonate anions. The superscript in each ion represents the charge, or oxidation state of the ion. Since a sodium ion has a charge of +1 and a carbonate ion has a charge of -2, it takes 2 sodium ions to balance the charge of a carbonate ion. We include this information as the subscript in the formula: Na2CO3, 2 Na+ for each CO3 2-. What about the subscript in CO3 2-? Well, that means there are 3 oxygen atoms for each carbon atom in the ion. But for now, you should think of the carbonate ion as a whole rather than as parts. In the simplest reactions, the carbonate ion is never broken up, so it makes sense to treat it as a single entity rather than the sum of its parts.
But if you don't recognize carbonate ion, you won't be able to take advantage of this great simplification and you will be hopelessly confused. Consider the compound NH4NO3. This is a simple ionic compound. Every chemist knows that it has two parts, a cation and an anion. Every chemist recognizes the two parts. But to a beginner there seem to be four parts. To a beginner the compound seems more complicated than it has to be because she doesn't recognize the parts. To make any sense of it, you need to know what the parts are.
So there are twenty names for you to learn. I have promised to teach you as little as possible. I tried to make it nineteen names, but I just couldn't do it. We need all twenty of them, and if you learn them now, you'll have an easier time later. You need to learn the charges as part of each formula, and associate each formula with its name. Eleven of them are for cations and nine of them are for anions. You can match any cation with any anion and get a valid chemical compound, so with this short list you can recognize the names of ninety-nine compounds. That's quite a lot for the little I've taught, but you know a lot of little will do.
The formula for a compound has two parts, the cationic part and the anionic part. Sodium chloride is NaCl, for example. The first letter of an element is always capitalized, so you can distinguish Co, the single element, cobalt, from CO, the compound, carbon monoxide. Turning compound names to formulas is easy once you know Table 7-1; just write the formula for the cation (positive ion, first name) followed by the formula for the anion (negative ion, second name). Hang subscripts on each one to make the charges balance. For example, let's write the formula of aluminum oxide. From the table, the aluminum ion has a charge of +3 and the oxide ion has a charge of -2. We want the positive charges to balance the negative ones. Well, 2×(+3) = +6 and 3×(-2) = -6 so the formula of aluminum oxide is Al2O3.
As a second example, try lead nitrate. The lead ion has a charge of +2 and the nitrate ion has a charge of -1. Since 1×(+2) balances 2×(-1), the formula of lead nitrate is Pb(NO3)2. The formula for nitrate goes in parentheses to show that it is a single anion made up of four atoms. A subscript after a parenthesis applies to everything inside. So lead nitrate contains 1 lead atom, 2 nitrogen atoms, and 6 oxygen atoms. Once the charges are balanced, we no longer write them in the formula because they have exactly canceled out. But the charges show up again when these compounds dissolve in water.
When the compounds we have been discussing dissolve in water, they behave very peculiarly; they ionize, or fall apart into their respective ions. For sodium chloride, we write:
NaCl(aq) = Na+(aq) + Cl-(aq)
The sodium and chloride ions, which were next to each other in solid sodium chloride, separate from one another in solution and float around on their own. Compounds that behave like this, falling apart into ions in aqueous solution, are called electrolytes, or salts. The (aq) in the formula stands for aqueous, that is, a solution in water. The water itself does not appear in the equation; it is simply the medium in which the ionization takes place.
Suppose for a moment that we dissolve two different electrolytes in water, say, sodium chloride and potassium nitrate. The sodium chloride would fall apart into sodium and chloride ions; the potassium nitrate would fall apart into potassium and nitrate ions. So we are left with a solution containing sodium, potassium, chloride, and nitrate ions. But exactly the same ions would have formed had we dissolved sodium nitrate and potassium chloride in water. Once in solution, a potassium ion, for example, does not "remember" whether it came from potassium chloride or potassium nitrate. It's just a potassium ion floating around. As long as the electrolytes remain in solution, no reaction takes place; we just have an ion soup.
But suppose now that of all the ways of mixing and matching the first and last, the cationic and anionic names, one of them is insoluble in water. If we dissolve sodium chloride and silver nitrate, we get a soup of sodium, silver, chloride, and nitrate ions. When a dissolved sodium ion happens to bump into a dissolved chloride ion, they don't stick because, after all, sodium chloride is soluble in water. When a silver ion happens to bump into a nitrate ion, they don't stick because silver nitrate is soluble in water. When a sodium ion happens to bump into a nitrate ion, they don't stick because sodium nitrate is soluble in water. But when a silver ion happens to bump into a chloride ion, they stick together to form insoluble silver chloride. Solid silver chloride rains down, or precipitates, accumulating as a layer of insoluble powder at the bottom of the solution. In other words, a reaction has occurred.
This type of reaction, which accounts for about half the reactions discussed in this book, is called a metathesis. Just as we represented a compound with a formula, we represent a reaction with an equation:
NaCl(aq) + AgNO3(aq) = NaNO3(aq) + AgCl(s)
In Chapter 3 we used the equation as a mathematical sentence, with the equal sign as the verb. We extend that notion here. A chemical equation is used as a shorthand for a reaction, the transformation of one set of elements or compounds into another. In such a reaction, the number of each kind of atom does not change. The Law of Conservation of Mass requires that the number of each kind of atom be conserved. So in an equation, we require that there be the same number of each kind of atom on either side of the equal sign. Sometimes an arrow or other symbol is used in place of the equal sign. The arrow has the advantage of implying the direction in which the reaction proceeds, but in the spirit of "teaching you as little as possible," I will use the equal sign.
Consider a slightly more complicated reaction:
CaCl2(aq) + Na2CO3(aq) = 2 NaCl(aq) + CaCO3(s)
In the beginning, there were calcium ions, chloride ions, sodium ions, and carbonate ions, all floating around in the solution. But when a calcium ion finds a carbonate ion, they stick together forming an insoluble precipitate, calcium carbonate. The sodium ions and chloride ions stay in solution. Why is it 2 NaCl and not Na2Cl2? Well, the ions from the tables are Na+ and Cl-, so we know the compound must be NaCl. The 2 is placed out front to balance the equation, that is, to make sure that the number of each kind of atom is the same on both sides. If there is no number out in front of a formula, it is assumed to be 1. We call the number out front the stoichiometric coefficient, which, though a mouthful, is shorter than saying "the little number in front of each formula in a balanced chemical equation." The unit of the stoichiometric coefficient is the mole. 1 mole calcium chloride + 1 mole sodium carbonate yields 2 moles sodium chloride + 1 mole calcium carbonate. Don't worry, the more you use words like stoichiometric coefficient and mole, the more you'll grow to understand them.
What about the reaction:
2 NaCl(aq) + CaCO3(s) = CaCl2(aq) + Na2CO3(aq)
Isn't this also a balanced reaction? For reactions in aqueous solution, the reactants must go into solution and at least one product must come out of solution. In this case, the calcium carbonate is not soluble, so it just sits in a lump at the bottom of the container. How do you know which compounds are soluble? You guessed it, another table, Table 7-2
Solubility in a Nutshell
Nitrates and acetates melt into water.
I'm sure that dozens of people will benefit from the little ditty presented in the sidebar Solubility in a Nutshell. Whether you memorize the verse or the table, you will be able to predict in some detail an enormous number (over 9,000) of possible chemical reactions. Let's look at some examples.
A: Both sodium chloride and potassium sulfate are soluble in water. If I swap the names I get potassium chloride and sodium sulfate and these are both soluble. So if I pour solutions of sodium chloride and potassium sulfate solutions together, nothing happens. All four ions stay in solution and there is no reaction.
A: Both sodium sulfate and silver nitrate are soluble in water. If I swap the names I get silver sulfate and sodium nitrate. Silver sulfate is insoluble and sodium nitrate is soluble. So if I mix a solution of sodium sulfate and a solution of silver nitrate, an insoluble precipitate of silver sulfate will form and sodium nitrate will remain in solution. Two go into solution, and one comes out. The balanced equation is:
Na2SO4(aq) + 2 AgNO3(aq) = Ag2SO4(s) + 2 NaNO3(aq)
A: Both lead sulfide and calcium carbonate are insoluble in water. If I put them in water, they both just fall to the bottom like sand or stones. Since they don't dissolve in water, there is no reaction.
A: Both copper sulfate and calcium hydroxide are soluble in water. Both copper hydroxide and calcium sulfate are insoluble. Two go in, two come out.
CuSO4(aq) + Ca(OH)2(aq) = Cu(OH)2(s) + CaSO4(s)
Notice that I put the OH in parentheses to make it clear that the subscript, 2, applies to the whole hydroxide ion, not just to the hydrogen.
A: This is a bit of a special case. When hydrogen ions bump into hydroxide ions in solution, they stick together and form water. This is a very energetic reaction, the basis of what we will later call acid-base reactions.
H2SO4(aq) + 2 NaOH(aq) = Na2SO4(aq) +2 H2O(l)
We use (l) for water instead of (aq) because (aq) would imply that water is dissolved in water; (aq) is always attached to a solute, the thing that is dissolved, not to the solvent, the thing in which it is dissolved.
Acute toxicity, the kind that makes you sick immediately, is often spoken of in terms of a material's LD50. This is the dose which would be expected, on average, to be lethal for 50% of test animals. Oh, it sounds morbid, I know. And it is. It isn't at all pleasant to contemplate cute little bunnies and mousies and doggies being force-fed chemicals to see how many of them croak. And since animals differ in their tolerance to different chemicals, it isn't an extremely reliable way to judge human toxicity. But it turns out to be very difficult to recruit human subjects for these tests, given that half of them are going to snuff it. So we use what we can get.
The LD50 is usually expressed in grams or milligrams of chemical per kilogram of body weight. If this looks like a unit factor to you, you haven't completely slept through the lessons of Chapter 3. You can use UFA to determine what would be a toxic dose for 50% of animals your size. But to judge relative toxicity, we can use the raw LD50 directly, recognizing that a big number means the chemical is relatively safe (it takes a lot to kill you) while a small number means it is hazardous (a little bit will do you).
Sometimes LD50 values are given in the MSDS. A reliable way to search for information online is to search for the CAS number, the keyword "MSDS," and the keyword "LD50." LD50's always refer to the animal which was tested (rat, mouse, dog, etc.) and the route by which it was introduced (oral, inhalation, etc.). An LD50 that is missing such details has been copied from an anonymous source at best, and may be completely fabricated at worst. Particularly when looking online, it's important to consider the source of your information.
Though this project involves no materials, you're not off the hook for material safety. To get your bearings, look up oral, rat LD50's for caffeine (CAS 58-08-2), sodium chloride (CAS 7647-14-5), sucrose (CAS 57-50-1), and sodium cyanide (CAS 143-33-9). Using these values, arrange these substances in order of toxicity.
|Research and Development|
So there you are, studying for a test, and you wonder what will be on it.