Chemistry for Cooks: An Introduction to the Science of Cooking

Chemistry for Cooks

Chemistry for Cooks

An Introduction to the Science of Cooking

Sandra C. Greer

The MIT Press

Cambridge, Massachusetts

London, England

© 2022 Massachusetts Institute of Technology

All rights reserved. No part of this book may be reproduced in any form by any electronic or mechanical means (including photocopying, recording, or information storage and retrieval) without permission in writing from the publisher.

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Library of Congress Cataloging-in-Publication Data

Names: Greer, Sandra C., 1945- author.

Title: Chemistry for cooks : an introduction to the science of cooking / Sandra C. Greer.

Description: Cambridge, Massachusetts : The MIT Press, [2022] | Includes index.

Identifiers: LCCN 2022000773 (print) | LCCN 2022000774 (ebook) | ISBN 9780262544795 (paperback) | ISBN 9780262372589 (pdf) | ISBN 9780262372596 (ebook)

Subjects: LCSH: Cooking—Experiments. | Food—Composition. | Chemistry, Technical.

Classification: LCC TX545 .G74 2022 (print) | LCC TX545 (ebook) | DDC 664/.07—dc23/eng/20220725

LC record available at https://lccn.loc.gov/2022000773

LC ebook record available at https://lccn.loc.gov/2022000774

10 9 8 7 6 5 4 3 2 1

d_r0

Dedicated to

Andrew Greer,

Enrico Rotelli,

Michael Greer,

Kaliel Roberts,

Arlo Greer,

Mack Greer,

Melina Murray,

Gwendolyn Bush,

Judith Chandler,

Janet Holmgren,

Gail Locke,

Patricia Miller,

Margaret Palmer,

Sue Rosser,

and Linda Richards,

and to the love that we share in the kitchen and around the table.

Contents

Preface

Acknowledgments

1 Some Basic Chemistry

Matter

Salt, Sodium Chloride, NaCl

Recipe Analysis: Lemon Salad Dressing #1

Recipe Analysis: Orange Granita, Adapted from Amanda Hesser

2 Measurements in Cooking

Mass, Weight, Volume, Density

Measuring in the Kitchen

Tools and Methods

Recipe Analysis: Pound Cake, from Viola Williams Thomason

Recipe Analysis: Biscuits, Adapted from the White Lily Foods Company and Edna Lewis

3 Heat and Temperature

Heat and Thermodynamics

Temperature: How Do We Measure the Heat Energy at Any One Place?

Energy and Calories: How Do We Measure the Amount of Heat in Matter?

Heat Capacity and Heats of Phase Changes: How Much Heat Do We Need to Add to Change the Temperature?

Heat Transfer: How Does Heat Flow from One Substance to Another?

Thermal Expansion: How Much Do Substances Expand When They Get Hotter?

Heat and Food: What Does Heat Do to Food?

Recipe Analysis: Roasted Asparagus

Recipe Analysis: Steamed Asparagus

4 Water, the Miracle Molecule

The Water Molecule and the Hydrogen Bond

Hydrophilicity and Hydrophobicity

Pressure

The Phase Diagram of Water

Solutions

Water Purity

Cooking and Water

Recipe Analysis: Dried Pasta, from Lidia Matticchio Bastianich

Recipe Analysis: Whole Artichokes by Pressure Cooker

5 Acids and Bases

Acids

Bases

pH for Measuring Acidity and Basicity

The Reaction between Acids and Bases: Neutralization

Acids and Bases in Food

Recipe Analysis: Sweet and Sour Pork, from Sylvia Tsai

Recipe Analysis: Polenta

6 Just Enough Organic Chemistry

Five Classes of Organic Compounds

Four Reactions of Organic Compounds

Recipe Analysis: Pears Poached in Wine, Adapted from Mitchell Davis

Recipe Analysis: Chicken with Tomatoes, Adapted from Jamie Oliver

7 Fats and Oils

Saturated and Unsaturated Fats and Oils

Hydrogenation and Trans Fatty Acids

Rancidity and Reactions

Melting and Smoking Points

Solubility and Density

Heat Transfer and Heat Capacities

Recipe Analysis: Orange and Olive Oil Cake, Adapted from Danai Kindeli

Recipe Analysis: Indian Butter Chicken, Adapted from Allrecipes

8 Carbohydrates

Monosaccharides and Disaccharides: Simple Sugars

Polysaccharides: Starches and Fibers

Carbohydrates in Plants

The Structure of Plant Cells

Cooking Plants

Recipe Analysis: Caramel Sauce in the Microwave, Adapted from Shirley Corriher, Andrew Janjigian, and Dan Souza

Recipe Analysis: Collard Greens, Adapted from Louise Childress Thomason

9 Proteins

Amino Acids and Protein Structure

Protein Denaturation and Coagulation

Enzymes

Animal Proteins

Plant Proteins

Cooking Proteins

Recipe Analysis: Scrambled Eggs

Recipe Analysis: Pound Cake, from Gail Blasingame Locke

10 More Chemical Reactions Plus Fermentation

Oxidation–Reduction or Redox Reactions

Maillard Browning or Nonenzymatic Browning

Fermentation: Yeast Fermentation and Bacterial Fermentation

Recipe Analysis: Caramelized Onions

Recipe Analysis: Whole Wheat Bread, from King Arthur Baking Company

11 Colloidal Dispersions

Colloidal Dispersions in General

Solids Dispersed into Liquids or Solids: Sols and Suspensions

Liquids Dispersed into Solids: Gels

Liquids or Solids Dispersed into Liquids or Solids: Micelles and Emulsions

Gases Dispersed into Liquids or Solids: Liquid Foams and Solid Foams

Thickeners for Sauces and Gravies

Effects of Salt and Sugar on Colloidal Dispersions in Water

Recipe Analysis: Lemon Salad Dressing #2

Recipe Analysis: Banana Pudding, from Joy of Cooking

12 Diffusion and Osmosis

Diffusion

Osmosis

Diffusion and Osmosis in Cooking

Recipe Analysis: Oven-Fried Chicken, Adapted from Judy Hesser

Recipe Analysis: Honeyed Pork Loin, Adapted from Diana Henry

13 Nutrition

Digestion

Nutrition

Recipe Analysis: Roasted Sweet Potatoes

Recipe Analysis: Pan-Fried Rib-Eye Steak

14 Food and the Senses

Flavor

Herbs and Spices

Food Additives

Recipe Analysis: Penne alla Vodka

Recipe Analysis: Fruit Cobbler, from Louise Childress Thomason

Final Thoughts

Index

Preface

Cooking was probably the greatest [discovery], excepting language, ever made by [humans].

—Charles Darwin, The Descent of Man, and Selection in Relation to Sex

Charles Darwin, one of the greatest scientists of all time, recognized that the cooking of food changed human history. The reasons for the impact of cooking all involve science. Cooking makes food easier to eat and to digest by using heat or chemicals to change the structure of plant or animal tissue, to release nutrients from the food cells, and to break food molecules into smaller, more digestible pieces. When we human beings can make better use of food sources, then we can spend less time obtaining food and more time on other activities.¹ In addition,

• Cooking kills germs by heating or otherwise treating food so that microorganisms die.

• Cooking can destroy poisonous chemicals that make food inedible.²

• Cooking preserves food by changing the enzymes that cause food to decay.³

• Cooking makes food taste and smell better by causing chemical reactions that improve flavors and textures, by releasing aromas, and by adding salt, spices, and herbs.

Chemistry for Cooks is a science textbook for undergraduate college students and for high school students. It is meant for students who are not science majors, and no background knowledge in science is needed. Everything you need to know is explained in the book. General readers, home cooks, culinary students, and professional cooks will find the text understandable. Anyone who likes to cook and wants to understand cooking better will find this book helpful.

The goal of this book is for you, the reader, to learn some science, so that you will:

• Understand what matter is made of and how the structure of matter is altered by heat, by interactions among molecules, and by reactions that change molecules;

• Understand how to treat food in order to change its microscopic structure and make it more delicious and more digestible;

• Learn to ask why a particular procedure or method is used and to think critically about what you do in the kitchen;

• See cooking as an enjoyable and interesting application of science.

While the title says chemistry, the world does not neatly divide into topics of chemistry and biology and physics, so we will also need to include some biology and a bit of physics. Chemistry for Cooks is set at a level between the books written for the home cook that include some science,⁴,⁵ and the books written for the professional food scientist that assume a full background of coursework in science.⁶,⁷ There are also culinary science books that are about trying different things to see what works best in the kitchen, but that is not science. Kitchen experiments become science only when the result is an understanding of why something works, not just what works. An understanding of the underlying science will make cooking more rational and less rigid, more adventurous, and even more fun.

We will begin with a brief introduction of what matter is made of, namely atoms and molecules. If you already know this basic material, you can skim over the first chapter. Do read the section in chapter 1 on the important food ingredient, salt (sodium chloride). Then the next chapters cover the other key topics of water, heat, acids and bases, carbohydrates, fats and oils, and proteins. Later chapters explain important reactions (Maillard, fermentation), colloidal dispersions (emulsions, suspensions, gels), diffusion and osmosis, the basics of nutrition, and the effects of food on the senses.

There is a recent trend in cooking that is called molecular gastronomy or modernist cuisine.⁸–¹² This trend came from a group of chefs, scientists, and food writers who were interested in understanding the scientific principles of cooking with laboratory devices such as constant-temperature baths, foam dispensers, liquid nitrogen freezing, and so on. Chemistry for Cooks shares with molecular cuisine the goals of understanding the science behind cooking and of being willing to rethink culinary traditions in the light of the science. It also shares the goal of using some simple scientific tools (such as digital thermometers) to improve cooking. However, this book will take a more traditional approach to cooking and to kitchen equipment.

Chemistry for Cooks is not a cookbook, but it does include recipes. Each chapter includes two cases of Recipe Analysis, in which a recipe is closely considered using the scientific principles in this book. Each chapter ends with Kitchen Hints that are useful ideas based on the content of that chapter. Each chapter includes exercises for the student, and also topics for projects that require more research and thought than do the exercises. While Chemistry for Cooks is not designed to support a laboratory course, many of the projects involve kitchen exercises that may be used at the discretion of the instructor. More resources for instructors are provided on the publisher’s website and include a solutions manual, a sample course syllabus, and PowerPoint slides for all the figures and tables.

You do not need any other books to use this book, but the references at the end of each chapter will allow you to read more about any subject that piques your interest. If you want to start your own shelf of books for cooking, I recommend that you start with two books (besides this one!). First, Harold McGee’s On Food and Cooking: The Science and Lore of the Kitchen,² while not a textbook, is an outstanding reference book to have on your shelf. Second, you will want to have at least one good general cookbook, such as the legendary Joy of Cooking,¹³ or one of the more recent books, such as The New Best Recipe¹⁴ or Cookwise: The Hows and Whys of Successful Cooking.¹⁵

When I was a girl, growing up in Greenville, SC, I was already very interested in chemistry. I was devoted to my chemistry sets and my microscope. My mother tried to interest me in cooking by saying that cooking is chemistry. She did not persuade me then, but later I taught myself to cook while I was a graduate student, because I could not afford to eat out all the time. Cooking has ever since been a way to make each day special with something nice to eat, and a way to show my love to my family and friends. Over time, I also became interested in the science of cooking, and how science can improve cooking. Learning about cooking science has made me a better cook. I hope that the same happens for you.

References

1. Richard Wrangham, Catching Fire: How Cooking Made Us Human (New York: Basic Books (Perseus), 2009).

2. Harold McGee, On Food and Cooking: The Science and Lore of the Kitchen, 2nd ed. (New York: Scribner, 2004).

3. Harold McGee, The Curious Cook: More Kitchen Science and Lore (New York: Macmillan, 1990).

4. The Editors at America’s Test Kitchen and Guy Crosby, The Science of Good Cooking: Master 50 Simple Concepts to Enjoy a Lifetime of Success in the Kitchen (Brookline, MA: America’s Test Kitchen, 2012).

5. The Editors at America’s Test Kitchen and Guy Crosby, Cook’s Science: How to Unlock Flavor in 50 of Our Favorite Ingredients (Brookline, MA: America’s Test Kitchen, 2016).

6. Srinivasan Damodaran and Kirk L. Parkin, eds., Fennema’s Food Chemistry, 5th ed. (Boca Raton: CRC Press, 2017).

7. H.-D. Belitz, W. Grosch, and P. Schieberle, Food Chemistry, trans. M. M. Burghagen, 5th German ed. (Berlin: Springer, 2004).

8. Nathan Myhrvold, Chris Young, and Maxime Bilet, Modernist Cuisine: The Art and Science of Cooking, 5 vols., vol. 1: History and Fundamentals (Bellevue, WA: The Cooking Lab, 2011).

9. Hervé This, Molecular Gastronomy: Exploring the Science of Flavor, trans. M. B. DeBevoise (New York: Columbia University Press, 2006).

10. Hervé This, Kitchen Mysteries: Revealing the Science of Cooking, trans. Jody Gladdening (New York: Columbia University Press, 2007).

11. Ferran Adria, Modern Gastronomy: A to Z (Boca Raton, FL: CRC Press, 2009).

12. Peter Barham, Leif H. Skibsted, Wender L. P. Bredie, Michael Bom Frøst, Per Møller, Jens Risbo, Pia Snitkjær, and Louise Mørch Mortensen, Molecular Gastronomy: A New Emerging Scientific Discipline, Chemical Reviews 110, no. 4 (2010): 2313–2365.

13. Irma S. Rombauer, Marion Rombauer Becker, Ethan Becker, John Becker, and Megan Scott, Joy of Cooking: 2019 Edition Fully Revised and Updated (New York: Scribner, 2019).

14. The Editors of Cook’s Illustrated Magazine, The New Best Recipe, 2nd ed. (Boston, MA: America’s Test Kitchen, 2004).

15. Shirley O. Corriher, Cookwise: The Hows and Whys of Successful Cooking (New York: William Morrow, 1997).

Acknowledgments

I am grateful to the Sonoma County Public Library and to the University of Maryland College Park Libraries for providing information resources. Mills College supported my teaching of a course called The Chemistry of Cooking in 2014. I thank Ruth E. Fassinger for suggesting the book title. I thank these friends for reading and commenting on the manuscript:

Svetla Baykoucheva,

Gwendolyn T. Bush,

John C. Gilbert,

Andrew Sean Greer,

Charles M. Knobler,

Sue V. Rosser,

and especially Gail Blasingame Locke.

Permission for the epigraph to chapter 8 is as follows:

You Can’t Hurry Love

Words and Music by Edward Holland Jr., Lamont Dozier, and Brian Holland

Copyright © 1965 Stone Agate Music

Copyright Renewed

All Rights Administered by Sony Music Publishing (US) LLC, 424 Church Street, Suite 1200, Nashville, TN 37219

International Copyright Secured All Rights Reserved

Reprinted by Permission of Hal Leonard LLC

Sandra C. Greer

San Rafael, CA

1

Some Basic Chemistry

Und die Chymie ist noch immer meine heimlich Geliebte.

Chemistry is still my greatest love.

—Johann Wolfgang von Goethe, Letter to Katharina von Klettenberg

In these few pages, we introduce the basic ideas from chemistry that you will need to understand this book. You will learn more chemistry as we go along, especially in chapter 6, Just Enough Organic Chemistry.

Matter

Our world and our universe are made of just two things: matter and energy. The ideas of both matter and energy are important for cooking. We will proceed here with matter and come back to energy in chapter 3.

Matter is basically the stuff the world is made of. Physicists formally define matter as that which has both mass and volume. Mass denotes the amount or quantity of matter. We usually measure mass by using weight, which is the pull of the earth’s gravity on matter (see chapter 2). Volume is the space taken up by matter. For example, water is a kind of matter and a cup of water has a volume of one cup and a weight of about half a pound.

For all ordinary purposes, matter and energy never vanish, no matter what we do to them. Both matter and energy can change form, but that does not mean that any matter or energy has disappeared, or that any new matter or energy has appeared. For example, water can change from solid ice to liquid water to gaseous steam, and its volume does change, but the amount (the mass) of the water does not change. The only time matter or energy vanishes or is generated is in nuclear reactions, and we do not want nuclear reactions in the kitchen!

Elements and Atoms: What Is Matter Made of?

The basic kinds of matter from which all other types of matter are constructed are called the elements.¹ There are 94 naturally occurring elements, and these 94 elements are enough to construct all of the many different kinds of matter in the world, from inanimate matter (salt, air, water) to living matter (plants and animals). Examples of the natural elements are oxygen, nitrogen, sulfur, and carbon. Table 1.1 is a list of all the naturally occurring elements that are important to food, nutrition, and cooking equipment, along with the standard symbol for each element. (There are also about 24 synthetic—not natural—elements, but they are not important for cooking.)

The smallest bit of each element is an atom of that element. The element symbol is used to represent one atom of that element. The atom itself is made (for our purposes) of three other particles: protons, neutrons, and electrons. The proton has a positive electrical charge (+1), the neutron has no electrical charge, and the electron has a negative electrical charge (–1). The atom consists of a central nucleus containing the neutrons and the protons, and a surrounding volume where the electrons exist.

Table 1.1 Natural elements that are important for cooking and nutrition

The identity of the atom is completely determined by the number of protons in the nucleus, which is called the atomic number. For example, the element hydrogen has one proton, the element helium has two protons, and so on. The atomic number is also equal to the number of electrons in the atom for a normal—electrically neutral and uncharged atom—because the negative charges of the electrons balance the positive charges of the protons. The number of neutrons, however, can vary for a given atomic number. Atoms with the same number of protons and electrons but different numbers of neutrons are called isotopes of that element; they are just different versions of the same element.

Figure 1.1 shows the architecture of an atom. The electrons are shown as a cloud of charge around the nucleus, in keeping with the ideas from physics that their exact positions are not known, but only the probabilities of each being in a particular place.

Figure 1.1

The structure of an atom. The red center is the nucleus with its positively charged protons and neutral neutrons. Negatively charged electrons are in a cloud (blue) around the nucleus. The relative sizes shown for the nucleus and the electron cloud are not to scale. If you were to use a golf ball to represent the nucleus, the whole atom would be about 3 miles in diameter!²

An atom can lose or gain electrons and become electrically charged. Then it is called an ion of that same element. Atoms that have lost electrons are positively charged (because they have more protons than electrons) and are called cations. Atoms that have gained electrons and are negatively charged (because they have more electrons than protons) are called anions. Ions are symbolized with either a + or – superscript on the symbol of the atom. For example, the sodium cation is Na+ and the chlorine anion is Cl–.

Again, the number of protons defines the element, the number of neutrons can vary for a given element, and, in a neutral atom, the number of electrons equals the number of protons. This is the basis of the Periodic Table, one of the great achievements of civilization. Figure 1.2 shows the Periodic Table. Each block on the table shows the name of the element and the symbol for the element. The number at the top of the block is the atomic number, equal to the number of protons in the nucleus, and thus equal to the number of electrons in the electrically neutral atom. The number at the bottom of the block is the atomic mass number or atomic weight number, which equals the number of protons plus the number of neutrons (but does not include the electrons since electrons do not have enough mass to matter). The atomic mass number is not a whole number because it is an average over the various isotopes of that element. For example, the atomic mass number of carbon is 12.011, which is the sum of six protons plus six neutrons, with a bit more added because carbon can sometimes have seven neutrons.

Figure 1.2

The Periodic Table. Each entry gives the atomic number of that element at the top, then the symbol, and then the name. At the bottom of each entry is the atomic weight (or mass—see chapter 2). In some cases, the table also gives a range for the atomic mass, since different samples from nature are different. In this book we will just use the upper, average value for the atomic mass. [Copyright © 2021 International Union of Pure and Applied Chemistry. Reproduced by permission of International Union of Pure and Applied Chemistry.]

There are four basic things to know about the Periodic Table. (Of course, there is much more to know, but for that you may want to take some courses or read some books on chemistry.)

1. The elements appear on the Periodic Table in the order of their atomic numbers. As the atomic number increases, the number of protons in the nucleus increases, and the number of electrons increases to keep the electrical charge balanced. For example, a sulfur atom has one less proton and one less electron than a chlorine atom; an argon atom has one more proton and one more electron than a chlorine atom, and so on.

2. The elements in the same column or group of the Periodic Table have very similar chemical properties. This is because the electrons are configured in similar ways for elements within the same group, and it is the electrons that determine the chemical properties. We will not go into the details of the electron configurations here. For example, sodium and potassium are both in group 1 and have similar chemical properties.

3. Across a row of the Periodic Table, the properties of the elements change very gradually, so potassium (row 4, group 1) will behave chemically more like calcium (row 4, group 2) than like copper (row 4, group 11).

4. Elements on the left side and in the middle of the Periodic Table are metals, and elements on the right side are nonmetals. Metals such as copper conduct both heat and electricity, whereas nonmetals such as sulfur do not conduct either.

Molecules, Compounds, and Chemical Bonds: Where Are the Electrons?

Some elements exist in nature as single atoms, including helium and the other elements below helium in the right-most group (group 18) of the Periodic Table—called the inert or noble gases.

However, most elements exist naturally in combination with other elements as molecules, the building blocks of compounds. Molecules are made by the connection—the bonding—of two or more atoms. The atoms in the molecule can be the same or different. Some elements exist in nature as a twosome: two atoms of the same kind, bonded together into diatomic molecules. Oxygen, nitrogen, and all the elements in group 17 of the Periodic Table (fluorine, chlorine, etc., called the halogens) form such diatomic molecules. Molecules can be as simple as diatomic oxygen molecules, and as complicated as the DNA (deoxyribonucleic acid) molecules that make up the genes of living things.

All of chemistry is about the sharing of electrons among molecules, and about changes in how the electrons are shared. The bonds between atoms in molecules are made by sharing electrons. The electrons can be shared equally between two atoms, they can be shared unequally between two atoms, or they can be given away entirely from one atom to another.

When two atoms form a chemical bond by sharing electrons equally, the bond is called a nonpolar bond, where nonpolar means that the negative electrical charge of the shared electrons is equally distributed around both atoms and there is no electrical charge asymmetry or polarity. This is most likely to happen where the two atoms are the same, as in the bond between two oxygen atoms to make a diatomic oxygen molecule. Since the atoms are the same, there is no tendency of the electrons to reside more around one atom than around the other atom.

When two atoms form a chemical bond by sharing electrons unequally, the bond is called a polar bond, where polar means that the negative electrical charge of the shared electrons is unequally distributed around the two atoms. Then an electrical polarity occurs: it is more negative on the side of the bond that attracts the electrons, and more positive on the side of the bond that lacks enough electrons to balance the positive charge of the nucleus. Most bonds in molecules have some degree of polarity, just because the electrons in a bond are not equally attracted by the nuclei of two different atoms. Electrons tend to be more attracted to elements toward the right-hand side of the Periodic Table than to elements toward the left-hand side of the Periodic Table. For example, in a molecule of water, two hydrogen atoms are bonded to one oxygen atom. The electrons that are shared between each hydrogen atom and the oxygen atom are more attracted to the oxygen atom, forming two polar bonds. Nonpolar and polar bonds both involve sharing electrons between atoms and both are called covalent bonds.

When the two atoms that form a bond are very different in the way in which they attract electrons, then the electrons can be completely transferred from one atom to the other, making an ionic bond. While the two atoms do not share electrons, they are bonded in that the one that has lost an electron has a positive charge, the one that has gained an electron has a negative charge, and the positive and negative charges attract one another. Ionic bonds are often between a metal atom that gives up electrons and a nonmetal atom that attracts electrons. For example, sodium (on the left side of the Periodic Table) and chlorine (on the right side of the Periodic Table) come together to make sodium chloride, NaCl, which has an ionic bond with an electron from the sodium atom transferred entirely to the chlorine atom to form two ions: Na+Cl–.

Polar, nonpolar, and ionic bonds can involve more than one electron. If only one electron is shared in the bond, it is called a single bond. If two electrons are shared, it is a double bond, and three shared electrons make a triple bond. For example, an atom of carbon can share four electrons with other atoms, and so can make four bonds;