What Einstein Didn't Know: Scientific Answers to Everyday Questionsby Robert L. Wolke
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Discover how cricket chirps can tell us the temperature, why you can't unburn a match, why ice floats, and a host of mysteries of modern living — including some riddles that maybe even Einstein couldn't solve. From the simple (How does soap know what's dirt? How do magnets work? Why do batteries die?) to the more complex (Why does evaporation have a cooling effect? Where does uranium get its energy?), this book makes science more understandable and fun.
Author Robert Wolke provides definitive and easy-to-comprehend explanations for things that we take for granted, like the illumination behind neon signs and the mysteries of beverage carbonation. Wolke also dares readers to explore and conduct their own experiments with food, kitchen utensils, and common household products. This fifteenth anniversary edition of his bestselling popular science classic has been completely revised and expanded.
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What Einstein Didn't Know
Scientific Answers to Everyday Questions
By Robert L. Wolke, Diana Zourelias
Dover Publications, Inc.Copyright © 2014 Robert L. Wolke
All rights reserved.
In the Marketplace
From the street vendor to the glitziest mall, it's the same old battle out there: people selling versus people buying. The sellers always have the advantage, because they know exactly what it is they're selling, while the emptors must be in a constant state of caveat. In many cases, the buyer not only doesn't know what the product really is, but can't 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. I'll 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; I'm 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. I've 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 can't 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 vapor—all gases. The water molecules in the vapor are content to stay that way—far apart from one another as an invisible gas, rather than clumping together as molecules of liquid—because they got there in the first place simply by leaping individually out of the beer's 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 that's the fog you see.
Then, if you put the bottle down on the bar without pouring it, you—and your observant customer—may 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 bottle's 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, "Hey, man. Why's my beer smokin'?" 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. What's 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 heat—get cold—or release heat—get 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 chemical's 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 it's 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 what's 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 we're talking about; we'll 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 crystal's structure, therefore, some heat energy must be borrowed from the water, and the water cools down accordingly.
Step 2: The liberated particles don't just swim around in splendid isolation. They have a strong mutual attraction for water molecules. If they didn't, they wouldn't 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 it's just a matter of which effect is bigger: the cooling effect from the breakdown of the solid or the warming effect from the particles' attraction for water molecules. If the cooling is bigger, the net effect will be that the water gets colder when the solid dissolves. That's 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; that's how it is with calcium chloride and magnesium sulfate.
Salt and sugar? In each case, it's 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, salt—sodium chloride—does cool the water very slightly when it dissolves.)
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. (Don't 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 "natural oyster shell." Is oyster-shell calcium better than other kinds?
Clams and oysters make their shells primarily out of calcium carbonate. But chemically speaking, it doesn't 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 "natural" (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 it's the actual element calcium that you're after; your metabolism doesn't 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 people's 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?
It's not smoke; it's fog. And it isn't carbon dioxide either, as some people think; carbon dioxide gas—CO2—is invisible. The fog surrounding the dry ice is made of tiny droplets of water, condensed out of the air's natural humidity by the dry ice's 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 "desolidifying" (transfiguring from the solid to a different state), while dry ice cannot be heated beyond –109°F (–78.5°C) without desolidifying into—turning into—not 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 doesn't melt; it changes directly into a gas without becoming a liquid first.
Why doesn't CO2 like to be in the liquid state?
Well, carbon dioxide molecules don't like one another very much; they don't stick together very well, the way water molecules do. Water molecules, H2O, have a central oxygen atom with two hydrogen atoms sticking out like devils' 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 didn't ask, but ... Why does a CO2 fire extinguisher shoot out a blast of snow?
It's 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 fire's fuel below its ignition point, while the carbon dioxide smothers the fire because it's 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 blanket—unless 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 it's 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 we're 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? There's a surefire way to tell. Check the size of the water droplets from the crashing waves. If they're the size of a porthole or a cannon ball on the ship, it's a model in a tank. Water just doesn't break up into drops the size of cannon balls, unless the "cannon balls" 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 way—as a liquid under high pressure in steel tanks. When the tank's valve is opened, the liquid instantly boils off into a burst of gas.
You didn't 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 loop, they lose some energy and slow down, just as a billiard ball loses some energy when it collides with another ball. And molecules that have slowed down are lower-temperature molecules by definition—low enough in some parts of the cloud to freeze into flakes of dry-ice snow.
This is not just a phenomenon of carbon dioxide. Any gas, when expanding (except when expanding into a vacuum), will lose energy and cool down. We will see this situation again when we let some air out of our overinflated car tires (p. 98).
My True Love Is Nothing but Skin and Bones
Somebody tried to tell me that clear, shimmering, sparkling-bright Jell-O, my childhood true love, is made from pig skins, cow hides, bones, and hoofs. Yuck! Can that possibly be true?
Of course not. Just the skins and bones. No hoofs.
Jell-O and similar desserts are about 87 percent sugar and 9 or 10 percent gelatin, plus flavoring and coloring. Kids love three things about it: it is brightly colored, it is very sweet, and it jiggles. (Love that jiggle.) Mothers don't mind, because gelatin is pure protein.
The gelatin, which of course is the jiggler, really does come from pig skins, cattle hides, and cattle bones. But stop squirming. Every time you have made soup or stock that jelled in the refrigerator, you have made gelatin out of chicken hides or beef bones.
The skin, bones, and connective tissue of vertebrate animals contain a fibrous protein called collagen. There is no collagen in hoofs, hair, or horns. When treated with hot acid (usually hydrochloric or sulfuric acid) or alkali (usually lime), collagen turns into gelatin, a somewhat different protein that dissolves in water. The gelatin is then extracted into hot water, boiled down, and purified.
Purified? You bet. But you don't want to see (or smell) the early stages of the process. By the time the gelatin leaves the factory, it has been thoroughly washed at various stages to get the acid or alkali out, then finally filtered, deionized (a way of removing chemical impurities), and sterilized. What eventually leaves the factory is a pale yellow, brittle, plastic-like solid in the form of ribbons, noodles, sheets, or powder. When solid gelatin is soaked in cold water, it absorbs water and swells up; then, when heated, it dissolves to form a thick liquid, which jells on cooling.
Excerpted from What Einstein Didn't Know by Robert L. Wolke, Diana Zourelias. Copyright © 2014 Robert L. Wolke. Excerpted by permission of Dover Publications, Inc..
All rights reserved. No part of this excerpt may be reproduced or reprinted without permission in writing from the publisher.
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Meet the Author
Robert L. Wolke is Professor Emeritus of Chemistry at the University of Pittsburgh and a food columnist for The Washington Post. The recipient of numerous awards, he enjoys nationwide renown as an educator, lecturer, and interpreter of science for lay audiences.
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its a cracker and a dream of a book which provides answers for all those question asked but no one knew the answer, well this books does. Questions like; What makes super glue - super, How come you can melt sugar but not salt, Why do water beds need heaters, How does steel rust, How does soap know what dirt is, Why do ocean waves roll in parallel to the shoreline and this dream which I really liked, Why are the earth, moon and sun all spinning. You will find all the answers here that have mystified you with plain explanation, by using scientific principles so that anyone can understand . HIGHLY RECOMMENDED - A MUST READ
This book offers very clear explanations to things we are very used to see but don't know really why. Not only this book will teach you, it will also make you laugh with its humor. (I read the spanish translation, but the original, if different, will be even better)