CHAPTER 1
In the Marketplace
From the street vendor to the glitziest mall, itâs the same old battle out there: people selling versus people buying. The sellers always have the advantage, because they know exactly what it is theyâre selling, while the emptors must be in a constant state of caveat. In many cases, the buyer not only doesnât know what the product really is, but canât even get a good look at it through the fog of promotion, packaging, and pitch.
In this chapter, we will take a clear-eyed look at what some products really are, beneath their glossy surfaces. We will visit a supermarket, a hardware store, a drug store, and a restaurant, with a stop or two at the local pub. Iâll even throw in a plug for the metric system.
A Foggy Day at the Bar
I must have opened thousands of bottles of beer. (No remarks, please; Iâm a bartender.) Many times, as soon as I pop the cap, wisps of fog appear in the neck of the bottle, and they sometimes even puff up above the opening. Iâve seen my share of foggy customers, but what causes foggy beer?
The fog is exactly the same as any fog: a collection of tiny particles of liquid water that have been condensed out of the air by a cold temperature, but that are too tiny to fall down like rain; they are kept suspended by being constantly bombarded by air molecules. They look white because they reflect all wavelengths of light equally.
Your puzzlement apparently stems from the fact that you canât see any fog inside the bottle until you open it, yet it is equally cold at both times. What is there about opening the bottle that makes the fog appear?
The space above the beer in the unopened bottle is filled with a mixture of compressed carbon dioxide, air, and water vaporâall gases. The water molecules in the vapor are content to stay that wayâfar apart from one another as an invisible gas, rather than clumping together as molecules of liquidâbecause they got there in the first place simply by leaping individually out of the beerâs surface. At the temperature of the beer, only a certain number of them will have had enough vigor to leap into the void as vapor. Another way of saying this is that the amount of water vapor in the air above the liquid is in equilibrium with the liquid water at that temperature. Those vapor molecules remain suspended individually as a vapor until you remove the cap and release the gas pressure.
When you release the pressure, the compressed gases are able suddenly to expand, and when gases expand, they lose some of their energy and cool down (p. 98). The water vapor is now cold enough to condense into droplets of liquid water, and thatâs the fog you see.
Then, if you put the bottle down on the bar without pouring it, youâand your observant customerâmay see some of the fog actually rising above the mouth of the bottle and spilling over onto the bar. With the pressure now released, carbon dioxide gas is leaving the beer and expanding as it hits the warmer air at the top of the bottle. As it expands, it lifts some of the fog above the bottleâs rim. Since carbon dioxide is heavier than air, it actually spills over like an invisible waterfall, carrying some of the fog along with it and flowing down the sides of the bottle.
Now see if you can explain all this to the guy who asks, âHey, man. Whyâs my beer smokinâ?â Alternatively, keep a few copies of this book on hand and let him read it. Good for a vigorous discussion.
And no offense, but if you worked in a higher-class establishment, you would notice exactly the same fog effect upon opening a bottle of champagne, and for exactly the same reasons.
Pressing Hot and Cold
When I sprained my ankle playing softball, somebody ran to a drug store and bought a cold pack. They squeezed it and shook it, whereupon it turned into an instant cold compress. Whatâs inside that package that makes it get cold so fast?
The cold pack contains ammonium nitrate crystals and a thin, breakable pouch of water. When the pack is squeezed, the water pouch breaks and, with a little shaking, the ammonium nitrate dissolves in the water.
When any chemical dissolves in water, it may either absorb heatâget coldâor release heatâget hot. Ammonium nitrate is one of those that absorb heat. It takes the heat right out of the water, thereby cooling it. And the amount of cooling is not trivial; that cold pack can actually get down close to freezing.
Because doctors keep blowing hot and cold about when to apply heat to an injury and when to apply cold, there are almost as many hot packs on the market as there are cold packs. The hot packs contain one of those chemicals that give off heat when they dissolve in water, usually crystals of calcium chloride or magnesium sulfate.
But why should a chemical absorb or release heat during the simple process of dissolving in water? After all, at home we dissolve crystals of two common chemicals, salt and sugar, in water time after time, yet we never see the sugar, for example, cooling off our hot coffee or heating up our iced tea. The fact is that salt and sugar are exceptions (see below).
When a chemical substance dissolves in water, it is a two-step process: first, the chemicalâs solid, crystalline structure must be broken down, and then a reaction takes place between the water and the broken-down chemical parts. The first step invariably has a cooling effect, while the second step has a heating effect (see below). If step one cools more than step two heats, as in the case of ammonium nitrate, the overall effect is cooling. If itâs the other way around, as it is with calcium chloride and magnesium sulfate, the overall effect is heating. In the cases of salt and sugar, the two steps happen to be just about equal, so they cancel each other out and there is very little change in temperature.
Here is whatâs going on during the two-step process in which a solid crystal dissolves in water. FYI, a crystal is a rigid, three-dimensional, geometric arrangement of particles. The particles may be atoms, ions (charged atoms), or molecules, depending on the substance weâre talking about; weâll just call them particles.
Step 1: First, the particles must be released from their rigid positions in the crystal in order to be able to float about freely in the water. To break down any rigid structure requires the expenditure of energy; somebody or something has to supply the sledgehammer blows that knock the structure apart. During the breakdown of the crystalâs structure, therefore, some heat energy must be borrowed from the water, and the water cools down accordingly.
Step 2: The liberated particles donât just swim around in splendid isolation. They have a strong mutual attraction for water molecules. If they didnât, they wouldnât have been interested in dissolving in the first place. As soon as they are in the drink, they are literally attacked by water molecules, rushing to cluster around them like magnetic mines around a submarine. When magnets (or water molecules) are attracted to something, they expend energy in their rush toward their targets. This energy, the energy of hydration, heats up the water.
Now itâs just a matter of which effect is bigger: the cooling effect from the breakdown of the solid or the warming effect from the particlesâ attraction for water molecules. If the cooling is bigger, the net effect will be that the water gets colder when the solid dissolves. Thatâs how it is with ammonium nitrate. On the other hand, if the warming effect is bigger, the net result is that the water gets warmer when the solid dissolves; thatâs how it is with calcium chloride and magnesium sulfate.
Salt and sugar? In each case, itâs just an accident that the two effects are approximately equal and cancel each other out. So there is practically no net cooling or heating when salt or sugar dissolves in water. (Actually, saltâsodium chlorideâdoes cool the water very slightly when it dissolves.)
TRY IT
Ammonium nitrate is a common fertilizer and calcium chloride is a common dehumidifier, sold for drying out damp closets and basements. You may have some of these chemicals around the house or farm. Stir some ammonium nitrate into water and the water will get very cold. Stir some calcium chloride into water and it will get quite hot. (Donât cover and shake; the heat can make the liquid splatter.) A tablespoon of the solid in a glass of water will do.
Oysters on a Half-Shelf
Half the calcium supplements on the health-food shelves seem to be ground-up ânatural oyster shell.â Is oyster-shell calcium better than other kinds?
Clams and oysters make their shells primarily out of calcium carbonate. But chemically speaking, it doesnât matter whether the calcium carbonate in the supplement bottle came from an oyster bed or a bed of limestone, which is also made of calcium carbonate. Neither is more ânaturalâ (whatever that means) than the other. Calcium carbonate is calcium carbonate. Oysters incorporate a bit of non-mineral matter in their shells, however, so calcium carbonate from other sources might be a bit purer.
Calcium supplements are sold in other chemical forms besides calcium carbonate (read the labels). Weight for weight, though, these other forms contain less calcium than calcium carbonate does, and itâs the actual element calcium that youâre after; your metabolism doesnât care about the other stuff. Calcium carbonate contains 40 percent calcium by weight, while calcium citrate contains 21 percent, calcium lactate contains 13 percent, and calcium gluconate contains only nine percent calcium. Now you can figure out which supplement on the shelf gives you the most calcium for your money.
But bear in mind that different chemical forms of calcium may be absorbed to different degrees in different peopleâs bodies. Nutritionists argue incessantly about this.
The Great Fog Forgery
Why is dry ice dry? And what makes all those clouds of smoke around it?
Itâs not smoke; itâs fog. And it isnât carbon dioxide either, as some people think; carbon dioxide gasâCO2âis invisible. The fog surrounding the dry ice is made of tiny droplets of water, condensed out of the airâs natural humidity by the dry iceâs low temperature.
Dry ice is carbon dioxide in solid form, just as regular ice is water in its solid form. Water ice cannot be heated beyond 32°F (0°C) without âdesolidifyingâ (transfiguring from the solid to a different state), while dry ice cannot be heated beyond â109°F (â78.5°C) without desolidifying intoâturning intoânot liquid CO2, but gaseous CO2, because it cannot exist in liquid form at normal atmospheric pressure. So a chunk of dry ice that is desolidifying (if I use that often enough, it may become a real word) into a gas is much, much colder than a chunk of ordinary ice melting into a liquid. Regular ice is wet because as it melts it becomes liquid water. Dry ice is dry because it doesnât melt; it changes directly into a gas without becoming a liquid first.
Why doesnât CO2 like to be in the liquid state?
Well, carbon dioxide molecules donât like one another very much; they donât stick together very well, the way water molecules do. Water molecules, H2O, have a central oxygen atom with two hydrogen atoms sticking out like devilsâ horns at an angle of 104.5 degrees to each other. In liquid water, those hydrogen horns form weak hydrogen bonds between adjacent molecules, binding them together with a mild sort of stickiness. This hydrogen-bond stickiness between water molecules is responsible for a number of unusual properties that make this common liquid categorically unique (p. 94).
You didnât ask, but . . .
Why does a CO2 fire extinguisher shoot out a blast of snow?
Itâs not water snow, but CO2 snow: flakes of dry ice.
A CO2 extinguisher is nothing but a high-pressure tank of liquid carbon dioxide with a squeeze valve. When you squeeze the valve, you let some of the liquid CO2 inside the tank escape. It instantly becomes a blast of very cold CO2 gas, accompanied by flakes of solid CO2 and a fog of water, condensed from the air. The extinguisher works in two ways: the coldness of the vapor can lower the temperature of the fireâs fuel below its ignition point, while the carbon dioxide smothers the fire because itâs a heavy, non-flammable gas that pushes away the oxygen.
Dry ice has been used on movie sets to fake fog. It is real fog, all right, because it consists of microscopic droplets of water suspended in the air. But you can always tell a fog forgery, because the water is very cold from the dry ice and the fog therefore lies on the ground like a blanketâunless it is blown around by an off-camera fan or stirred up by a mob of stumbling zombies. Real, weather-generated fog, on the other hand, hangs fairly motionless in the air.
Movies and plays use dry ice also to simulate cauldrons of boiling water. Just throw some dry ice into the water, and as the solid carbon dioxide changes to gaseous carbon dioxide, it rises through the water in fog-filled bubbles that break at the surface and are supposed to emulate hot steam. If you look closely, though, you can always tell that itâs fake. Steam goes straight up because hot air rises, while the cold dry-ice fog hangs low over the cauldron. Again, like a blanket.
While weâre on the subject of movie fakes, how about those scenes of storm-tossed ships? Are they just miniature models, shot at slow motion in a big tank? Thereâs a surefire way to tell. Check the size of the water droplets from the crashing waves. If theyâre the size of a porthole or a cannon ball on the ship, itâs a model in a tank. Water just doesnât break up into drops the size of cannon balls, unless the âcannon ballsâ on the ship are really BBs on a scale model.
But CO2 molecules are shaped like this: O=C=O. They have no sticky hydrogen horns to bind them together, so they cannot condense into the tightly crowded structure of a liquid unless forced together by a high pressure. Carbon dioxide is shipped around the country this wayâas a liquid under high pressure in steel tanks. When the tankâs valve is opened, the liquid instantly boils off into a burst of gas.
You didnât ask this either, but . . .
Why is the blast of snowy gas that comes out of a CO2 fire extinguisher so cold, even though the extinguisher may have been sitting around in the room for months?
When you release some of the pressurized, liquid carbon dioxide from the high-pressure tank into the low-pressure room, it instantly flashes off into a gas, which must then expand rapidly under the reduced pressure. In order to expand, the gas must make room for itself by knocking other stuff, say the air molecules in the room, out of the way. When the CO2 molecules knock the air molecules for a lo...