Why Calories Count
From Science to Politics
By Marion Nestle, Malden Nesheim
UNIVERSITY OF CALIFORNIA PRESS Copyright © 2012 Marion Nestle and Malden Nesheim
All rights reserved.
What Is a Calorie?
In embarking on an entire book about calories, we have to begin at the beginning —what to call them. Calories are units of work or heat, but what they are called depends on who is doing the calling. We think the name inconsistencies can be so confusing that we summarize them in table 1. We believe there is an easier, commonsense way to think about the definition, as we will explain. To get to that point, let's begin with the official definition used by chemists:
One calorie is the amount of heat energy needed to raise the temperature of one gram of water by one degree centigrade, from 14.5° to 15.5°, at one unit of atmospheric pressure.
As a first step toward arriving at a more intuitive definition, let's ignore the ambient temperature and atmospheric pressure, as they usually do not make much difference except to scientists. But pay close attention to the weight units. In table 1, we emphasize gram in this definition to distinguish it from definitions based on heating an amount of water 1,000 times greater—a kilogram. A gram is about one quarter of a teaspoon. A kilogram is 2.2 pounds.
Gram units are inconvenient for most Americans, and so are so-called small calories, based on a gram standard. If food labels listed calories by the chemists' definition, your daily calorie requirements might be about 2,500,000 a day rather than 2,500, and a carrot would provide 25,000 calories instead of 25. That is why nutritionists much prefer to use kilocalories (kcal), units that are 1,000 times larger and based on the heat required to raise the temperature of a kilogram of water by one degree centigrade.
Here is where things get tricky. Because it feels awkward to say kilocalories, nutritionists shorten it to Calories, spelled with a capital C. Food labels list Calories, but they really mean kilocalories (also known as large calories or kilogram-calories). When spoken, Calories (capital C) and calories (small c) sound the same. As a result, when discussing the energy value of food, the words kilocalories (kcal), Calories (Cal), and calories (cal) have come to mean exactly the same thing: 1,000 chemists' calories. Hence: confusion.
The peculiar result of all this is that the word calories can mean two things at the same time: chemists' calories and nutritionists' Calories, which are 1,000 times greater. In common practice, most people use calories, capitalized or not, to mean kilocalories or Calories, despite the confusion this causes.
James Hargrove, who has written a fine history of the evolution of the terms, argues that "it is untenable to continue to use the same word for different thermal units (gram-calorie and kilogram-calorie) and to use different words for the same unit (Calorie and kilocalorie). The only valid use of the Calorie is in common speech and public nutrition education."
We, as it happens, are in the business of public nutrition education, and we use common speech all the time. We do not view this situation as untenable, just confusing. The various definitions lead to problems that can be amusing or annoying, depending on how you look at them. For example, Hargrove points to the inconsistent use of the terms calories and kilocalories—in this case meaning the same thing—in the U.S. Code of Federal Regulations rules for food labels: "A normal serving of the food contains at least 40 kilocalories (that is, 2 percent of a daily intake of 2,000 kilocalories).... The food contains all of the following nutrients per 100 calories based on 2,000 calorie total intake as a daily standard."
Public nutrition educators that we are, we follow common speech. Throughout this book we use calories, spelled with a small c, to refer to kilocalories/ Calories, unless we have a specific reason to use a more precise term. Food labels do this too.
With this said, we must now introduce another source of confusion. We live in the United States. Practically everywhere else in the world, people use the metric system and follow the metric-based International System of Units (Système International d'Unités). In this system, food energy is expressed in joules or in their thousandfold-larger counterpart, kilojoules. We give the conversion factors in table 1. These terms also are used imprecisely, and we hear people talking about joules when they really mean kilojoules.
The easiest way to think of the conversions is to remember that 100 calories is about 420 kilojoules (or calories times 4.2) and that 2,400 calories is equivalent to 10,000 kilojoules. If you only want estimates, you can multiply calories by 4 to get kilojoules or divide kilojoules by 4 to get calories. This is obviously imprecise but should work well enough for most purposes. As you will soon see, "works well enough for most purposes" is a constant theme in this book.
And just for the record, let's add one more term. Nutrition scientists like to use megajoule (MJ) to describe energy intake or expenditure because it doesn't require as large a number. A megajoule is 1,000 kilojoules. A diet of 2,400 calories a day has 10,000 kJ or 10 MJ.
When the U.S. Food and Drug Administration wrote regulations for Nutrition Facts labels on food products in 1993, it seriously considered including both kilojoules and calories in the design but eventually rejected the idea, reasoning that no attempt to induce Americans to use the metric system has ever succeeded. The historian Hargrove is thoroughly fed up with all this. He recommends getting rid of everything having to do with calories and switching to joules instead. Good luck with that. We doubt this will happen within our lifetime or even that of our children. In the meantime, we think calories works pretty well, once you get a feel for what they are. To that end, we invite you to take a look at our approach to understanding them more intuitively.
GETTING A FEEL FOR CALORIES
Our first recommendation is to convert the chemists' definition of Calories/ kilocalories to something more convenient. We have already gotten rid of temperature and pressure. Let's now convert metric units to common household measures. A liter of water weighs about one kilogram and is almost the same volume as a quart (one liter equals 1.06 quarts). We think it is fair to say:
A calorie is about the amount of heat needed to raise the temperature of a quart of water by 1°C.
One food calorie isn't very much in diets of 2,000 to 3,000 a day. It makes more sense to deal with 100 calories. This is the amount in 100-calorie packs of snack foods, obviously, but also in a pat of butter, an apple, an 8-ounce soft drink, or two Oreo cookies. If 1 calorie raises the temperature of a quart of water by 1°C, then:
100 calories raises the temperature of one quart of water by 100°C.
One more tweak and we'll be done: 100°C is the same as 212°F—the boiling point of water. Here, at last, is our easy definition.
100 calories is the amount of heat needed to bring a quart of water to the boiling point.
What? You may feel warm after eating a big meal, but not that warm. Bodies contain about six quarts of blood. Why doesn't your blood boil when you devour a 600-calorie cheeseburger? The answer: metabolism. Metabolism, as we explain in chapter 5, taps off that heat in tiny increments to maintain body temperature and to power the digestion of food, the construction of new body molecules, and the action of muscles. Any energy you don't need right away gets stored, mostly as fat, to be used later, when you aren't eating.
Basic ideas about how calories work in food and in the body are quite straightforward, if somewhat abstract:
Calories are units of energy.
Energy is the capacity to do work.
Work can be chemical (biochemical, in this case) as well as physical (muscular).
Biochemical reactions and muscle activity produce heat.
Heat can be measured as calories.
Food provides energy that fuels the work that bodies do: breathe, circulate blood, keep warm, transmit nerve impulses, excrete waste, move. Scientists measure the overall energy required for this work by the heat it produces, which they report in units of calories. This is much the same as the use of temperature to indicate the heat of the surrounding air. You know intuitively that you are going to shiver when the outside temperature is 20 degrees Fahrenheit but will be sweating if it is 100 degrees. Similarly, you can intuit that 100 calories doesn't represent very much food energy, whereas 5,000 is more than you are likely to need in a day.
If the notion of calories as a measure of heat does not seem obvious, it is surely because you cannot have a physical sense of them in food. Calories, recall, have no smell, taste, or appearance. Invisible as they are, you can only measure them with complicated equipment or tests or deduce them from what they do to your body weight.
The calories in a food are measured through the heat it produces when burned. This leads to one other potential source of confusion—the difference between energy, heat, and calories. Energy, defined as the capacity to do work, exists in different forms: heat energy, electrical energy, chemical energy, mechanical energy. In physics, the first law of thermodynamics states that these various forms of energy are completely interchangeable. Any form of energy can be converted into any other form. The steam engine is the obvious example: it uses the heat generated by burning fuel to produce steam that creates mechanical work.
Bodies also interconvert forms of energy. The chemical structures of certain food molecules—proteins, fats, and carbohydrates (and alcohol, which we will get to in chapter 11)—store energy. Metabolism transforms some of the chemical energy stored in food components into heat energy through a series of biochemical reactions that involve the oxygen you breathe. These are called oxidation reactions. We will say more about them in chapter 5. Metabolism also transforms food energy into mechanical energy in muscles and into electrical energy in the brain and nervous system.
Heat energy may appear confusing because you can measure it in two ways. You can stick a thermometer into a roasting turkey or a feverish child to measure internal heat. But to measure the energy stored in foods or expended in physical activity, you must use special devices or chemical techniques (see chapters 3 and 4). To summarize:
Energy is the capacity to do work. Food stores energy in the chemical structures of some of its molecules, mainly proteins, fats, and carbohydrates but also alcohol.
Heat is one form of energy, and the one most easily measured using a thermometer or devices that permit calories to be calculated.
Temperature is a measure of the heat energy produced by the internal atomic motion of molecules of air, water, or other matter, as indicated on centigrade (Celsius) or Fahrenheit thermometer scales.
Calories and joules are units of heat or work energy measured in experiments using calorimeter devices or through calculations based on chemical techniques.
Because heat is common to all of these measurements, we thought the diagram in figure 1 might help clarify the distinctions.
The diagram distinguishes energy produced by the motion of all molecules of food from the energy stored in the chemical bonds of the particular molecules that yield calories. The atoms in all molecules move to some extent, and the heat generated by that movement can be transferred to other molecules. Thermometers measure that transferable heat.
But only some molecules—proteins, fats, carbohydrates, and alcohol—store energy that can be metabolized and used to support life. The energy of motion and the energy stored in chemical bonds are different. Here is one way to think of the distinction: from the standpoint of metabolism and body weight, french-fried potatoes served steaming hot right out of the fryer have the same number of calories stored in their proteins, fats, and carbohydrates as they do when refrigerated or frozen.
In speaking about the energy content of foods and the effects of food energy in the body, we try to be consistent about using heat as one form of energy and calories as a way to measure the amount of heat energy. With the main sources of calorie-related confusion exposed, if not resolved, we can now take a brief look at how scientists came to understand all this.
From Ancient Greece to Modern Calorie Science
Although it may seem self-evident that food is essential to life, scientists did not have much real understanding about how food energy keeps bodies warm, growing, and functioning until the late 1700s. The earliest understanding of calories as energy released by the interaction of oxygen with food molecules is usually attributed to Antoine Lavoisier, who lived and died in the eighteenth century. Lavoisier's view of metabolism as an oxidation process—the "burning" of food molecules in the presence of oxygen—still holds true.
We explain how this works in later chapters, but to jump-start the discussion let's begin with a quick overview of the basic concepts. The three energy-producing molecules in foods—proteins, fats, and carbohydrates—are highly structured. Creating them takes energy, some of which gets stored in their chemical bonds. Once eaten, food molecules are attacked by digestive enzymes. Enzymes break the organized structures into pieces small enough to be absorbed through the wall of the digestive tract. As we illustrate in chapter 5, enzymes disassemble complex molecules of starch, for example, into simple sugars, a process that releases small amounts of energy.
During metabolism, enzymes further disassemble the absorbed sugars into even smaller pieces, releasing more energy. Other enzymes transfer some of this energy to small storage molecules. When the body needs energy, still other enzymes split the storage molecules to release it. The released energy powers the chemical reactions that create new body molecules or make muscles contract, and some of it keeps the body warm.
The chemical reactions involved in these processes are "oxidation" reactions. They require oxygen, which is obtained through respiration. You breathe in oxygen. You exhale carbon dioxide as waste. Your other metabolic waste product is water, which you mostly excrete in urine. Food molecules that are completely "burned" in the body end up as carbon dioxide and water. All of these processes yield heat.
Scientists now know a great deal about the chemical structure of food and each of the processes crucial to obtaining energy from food: digestion, absorption, metabolism, respiration, and excretion. We'll examine them all in subsequent chapters. But first, it's worth exploring why it took such a long time to figure out the role of food in the body.
ANCIENT IDEAS ABOUT FOOD ENERGY: HIPPOCRATES AND GALEN
The most important early writings about energy and body heat are attributed to the Greek physician Hippocrates (~460–370 B.C.). Because the authorship of ancient texts is uncertain, the writings of Hippocrates are generally considered to include not only those of the physician himself but also those of his contemporaries and followers. In that sense, Hippocrates is something of a collective term. The writings, however, come across as if written by a single person, one who tried hard to make sense of his observations but sometimes jumped to conclusions that do not always make sense in modern terms. Even when his observations are right on target, he can sound much like a present-day diet guru. (Continues...)
Excerpted from Why Calories Count by Marion Nestle, Malden Nesheim. Copyright © 2012 Marion Nestle and Malden Nesheim. Excerpted by permission of UNIVERSITY OF CALIFORNIA PRESS.
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