This sample chapter is for review purposes only. Copyright © The Goodheart-Willcox Co., Inc. All rights reserved.
A
food scientist must understand
and predict how food will react in processing, packaging, and preservation.
This requires an understanding of the
nature of the particles in food. The
branch of science that studies these particles and how they are categorized is
chemistry.
Chapter 4 looks at the basics of
chemistry needed for the study of food
science. Particles are defined and classified according to their physical and
chemical characteristics. You will study
how these particles combine and break
apart. You will also see how scientists
describe chemical reactions in writing.
Change in the structure or position of
these particles requires energy. Chapter 5
describes the types of energy. It
describes how energy is transferred and
measured. It also discusses energy’s
importance in food production.
Chapter 6 explores a category of particles called ions. The electrical nature of
ions is basic to many reactions that occur
in food mixtures. This chapter defines ions
and describes how they are measured. It
also examines some of the important
applications of ions in the food industry.
Water is key to life and a main component of most foods. Chapter 7 identifies the unique chemical characteristics
of the particles that make up water. It
discusses water’s role in chemical reactions. This chapter also examines how
water functions in food preparation and
in a nutritious diet.
This photomicrograph shows a substance that is
essential in many chemical reactions in food
products and in the human body—water.
76
Unit II
Basic Chemistry
4 Basic Food Chemistry:
The Nature of Matter
5 Energy: Matter in Motion
6 Ions: Charged Particles
in Solution
7 Water: The Universal Solvent
Chapter 4
Chapter
4
Basic Food
Chemistry: The
Nature of Matter
Objectives
Key Terms
After studying this chapter,
you will be able to
describe the basic structure of atoms.
identify symbols on the periodic table
commonly used in food science.
define ionic and covalent bonding.
explain the difference between pure substances and mixtures.
compare physical and chemical reactions in
laboratory experiments.
balance chemical equations to illustrate
simple chemical reactions.
chemistry
matter
atom
subatomic particle
nucleus
proton
neutron
electron
orbital
element
atomic number
atomic mass
atomic mass unit
compound
molecule
chemical formula
chemical bond
shell
ionic bond
ion
covalent bond
78
To understand why ingredients react the
way they do in recipes or formulations, you
need to understand some basic chemistry.
Chemistry is the study of the makeup, structure, and properties of substances and the
changes that occur to them. It is the study of
matter, which is anything that occupies space
and has mass.
The Basic Nature of Matter
Everything you encounter, whether plant,
animal, or mineral, is made up of atoms. An
atom is the smallest unit of any elemental substance that maintains the characteristics of
that substance. In other words, one atom of
iron has the same physical characteristics as a
chunk of iron. Atoms are extremely tiny. You
cannot see them even through a powerful
microscope.
Knowledge of atoms was once based on
scientific theories and indirect experiments.
Only recently have devices become available
that allow scientists to map the individual
locations and shapes of atoms. Fortunately, it
is not necessary to see individual atoms to
learn much about them.
Students often use ball-and-stick models of molecules to help them understand the basic nature of
matter.
Lewis structure
valence electron
double bond
pure substance
organic compound
inorganic compound
mixture
homogeneous
mixture
heterogeneous
mixture
solution
solute
solvent
physical change
phase change
chemical change
reactant
product
law of conservation
of matter
79
Basic Food Chemistry: The Nature of Matter
Subatomic Particles
Although atoms are the smallest unit of
any element, they are not the smallest particles
known. Each atom is composed of smaller
parts called subatomic particles. The nucleus,
or central core of the atom, contains tightly
clustered particles of protons and neutrons. A
proton is a subatomic particle that has a positive electrical charge. A neutron is a subatomic
particle that is electrically neutral. Protons and
neutrons have about the same mass.
The third particle in an atom is called an
electron. Electrons have a negative electrical
charge that is equal to, but opposite of, the
positive charge of protons. Electrons are much
smaller than protons or neutrons. It takes
approximately 1,836 electrons to equal the
mass of one proton. The reaction between the
positive and negative charges of protons and
electrons causes the electrons to spin around
the nucleus. Electrons prefer to move in pairs.
The space occupied by a pair of electrons in an
atom is called an orbital. See 4-1.
Parts of an Atom
–
–
nucleus
proton
–
+
+ + +
+
neutron
electron
–
–
4-1 Protons and neutrons are tightly clustered in
the nucleus of an atom. Electrons travel at
tremendous speeds in the space outside the
nucleus.
Elements
An element is a substance that contains
only one kind of atom. There are 90 naturally
occurring elements known on earth. All matter in the universe is composed of one or more
of these elements. Scientists have even used
some of these elements to create approximately
20 additional elements.
The number of protons in the nucleus
determines which element an atom is. Pure
oxygen is composed of atoms with 8 protons
in the nucleus. Calcium, one of the main components of bone tissue, is composed of atoms
with 24 protons in each nucleus.
A simple system of symbols is used to
identify the elements. These symbols are a
form of abbreviation. Learning these symbols
will make it easier for you to record chemical
reactions.
The symbol for many elements is the first
letter of the element’s name. This letter is capitalized. The symbol for carbon is C. The problem is that eleven elements begin with the letter C. A second letter from the name of the element is added in these cases. For example, calcium is represented as Ca. Note the second letter is lowercase.
You may wonder why the symbol for
potassium is K and iron is Fe. Some of the elements’ symbols come from their names in
80
Unit II
other languages. Kalium is Latin for potassium;
ferrum is Latin for iron. See 4-2.
The Periodic Table
In the nineteenth century, researchers
became aware of links between the physical
and chemical characteristics of elements.
These properties seemed to repeat in a regular
fashion. In an effort to classify elements by
these relationships, a Russian chemist, Dmitry
Mendeleyev, developed the periodic table. This
chart helps show how elements relate to and
react with one another. He was able to use the
chart to predict the existence and properties of
elements that were unknown at the time.
Each cell of the periodic table gives information about one chemical element. The format of this information can vary somewhat
from source to source. However, it will usually
include the symbol for the element. It will also
point out some physical features of the atoms
of the element. These features help distinguish
an atom of one element from an atom of
another.
Two characteristics of atoms shown in
many periodic tables are the atomic number
and the atomic mass. The atomic number is
the number of protons in the nucleus of each
atom of the element. The atomic mass is
Basic Chemistry
Elements Most Commonly Found
in Foods
Element
Symbol
Aluminum
Al
Ca
Calcium
Carbon
Chlorine
C
Cl
F
Fluorine
Hydrogen
Iron (Ferrum)
H
Fe
Mg
Magnesium
Nitrogen
N
O
Oxygen
Phosphorus
Potassium (Kalium)
P
K
Sodium (Natrium)
Sulfur
Zinc
Na
S
Zn
4-2 Learning symbols for elements commonly
found in foods will help you read and write
chemical equations in food science class.
Nutrition News
Elemental Nutrition
Metals and Nonmetals
Many periodic tables use a color key to
group elements into various classes. One main
classification that is often used groups elements as metals and nonmetals. Elements in
these groups have a number of properties in
common. They also tend to react with other
elements in similar ways.
Except for hydrogen, the elements on the
left side of the periodic table are metals. (Even
hydrogen is a metal at very low temperatures
and very high pressures.) Metals are usually
shiny solids at room temperature. (Mercury is
an exception. It is a liquid at room temperature.) They are good conductors of heat and
electricity. Metals can also be drawn into wires
or pounded into sheets.
Nonmetals are found on the right side of
the periodic table. Many of these elements are
gases at room temperature. They tend to be
poor conductors of heat and electricity, and
they are generally brittle when solid.
Compounds and Chemical Formulas
When two elements combine, they form a
new substance totally unlike the elements
from which it was made. Compounds are substances in which two or more elements have
chemically combined. The basic unit of any
compound is a molecule. Sodium (Na) is a soft
silver metal that is explosively reactive.
Chlorine (Cl) is a yellowish-green gas that is
poisonous. Both elements are dangerous in
large amounts. When Na and Cl combine,
they form sodium chloride. Do you recognize
the name? It is the chemical name for table
salt. This salt is a white granule that is safe to
eat. The chemical formula that represents a
single molecule of the compound sodium
chloride is NaCl.
A chemical formula is a combination of
symbols of the elements that make up a compound. The chemical formula represents one
molecule of a compound. If you know how to
read a chemical formula, you can quickly
identify the elements in a substance. For
instance, the chemical formula for water is
H2O. From this formula, you can see that
water is composed of hydrogen and oxygen.
See 4-4.
In some chemical formulas, you will see
subscript numbers written to the right and
slightly below some atomic symbols. These
numbers tell how many atoms of each kind are
in a molecule. The 2 in the formula for water
indicates there are two atoms of hydrogen in a
water molecule. If a number does not follow a
symbol, the molecule contains only one atom
of that element. Therefore, the formula H2O
tells you that a molecule of water has two
atoms of hydrogen and one of oxygen.
Chemical formulas can also tell something
about how the atoms are arranged in a molecule. A molecule of acetic acid (vinegar) contains two carbon atoms, four hydrogen atoms,
and two oxygen atoms. The chemical formula
could be written as C2H4O2. However, each
carbon atom can combine with up to four
other atoms. Therefore, many arrangements of
the atoms are possible. The structure of an
acetic acid molecule is similar to the figure
below.
H
O
H—C—C—O—H
—
nutrition. These elements make
up carbohydrates, fats, and proteins. (Proteins also contain nitrogen, and they often contain sulfur, too.) These three nutrients
have many functions in the body,
including meeting all your energy
needs.
approximately equal to the sum of the masses
of protons and neutrons in an atom. The mass
of a proton or neutron is defined as equal to
one atomic mass unit. The mass of an electron
is 1/1,836 of an atomic mass unit. Electrons are so
small their mass is insignificant.
Not all atoms of an element have the same
number of neutrons. Scientists have calculated
an average atomic mass for each element. This is
the number used on periodic tables. The atomic
number and atomic mass are not in the same
places within the cells of all periodic tables. It is
simplest to remember the atomic number is a
smaller number than the atomic mass.
The organization of the cells in the periodic
table is a key to how elements will interact
chemically. The table is arranged in columns
and rows. The columns, which are called
groups, are often numbered left to right from 1
to 18. The rows, which are called periods, are
numbered top to bottom from 1 to 7. See 4-3.
Later in the chapter, you will read more about
what an element’s group and period tells you
about the element.
—
—
strip some foods of these vital
elements. This is one reason
everyone needs to include fresh
fruits and vegetables and whole
grain products in their daily diets.
Besides the dietary minerals,
the elements carbon, hydrogen,
and oxygen are vital to good
81
Basic Food Chemistry: The Nature of Matter
—
A number of elements have
been found to be essential for
good health. Many of these are
the nutrients dietitians and other
health professionals call minerals.
Each of these minerals has at
least one important function in
the body. Food processing can
Chapter 4
H
82
Chapter 4
83
Basic Food Chemistry: The Nature of Matter
Lawrencium
103
Lr
260.105
electron
shells
electrons
Einsteinium
99
Es
252.083
Berkelium
97
Bk
247.070
Plutonium
94
Pu
244.064
Neptunium
93
Np
237.048
Protactinium
91
Pa
231.036
Thorium
90
Th
232.038
Uranium
92
U
238.029
Americium
95
Am
243.061
Curium
96
Cm
247.070
Californium
98
Cf
251.080
257.095
Mendelevium
101
Md
258.099
Nobelium
102
No
259.101
Lutetium
Fermium
100
Fm
70
Yb
173.04
Ytterbium
Thulium
69
Tm
168.934
Ho
164.930
Er
167.26
Erbium
68
Holmium
67
Dy
162.50
Gd
157.25
Dysprosium
66
Terbium
65
Tb
158.925
Gadolinium
64
Eu
151.965
Sm
150.36
Pm
144.913
Europium
63
Samarium
62
Promethium
61
Nd
144.24
59
Pr
140.908
Neodymium
60
Praseodymium
Cerium
58
Ce
140.115
Uub
(unnamed)
112
(unnamed)
111
Uuu
(unnamed)
110
Uun
Mt
(266)
Hs
(265)
Bh
(262)
Meitnerium
109
Hassium
108
Bohrium
107
Sg
(263)
Seaborgium
106
Dubnium
105
Db
(262)
Rf
(261)
227.028
Rutherfordium
104
Actinium
89
Ac
Radium
88
Ra
226.025
Fr
223.020
7
Francium
87
71
Lu
174.967
222.018
Radon
86
Rn
Astatine
85
At
209.987
84
Po
208.982
Polonium
Bismuth
83
Bi
208.980
82
Pb
Lead
207.2
204.383
Thallium
81
Ti
79
Au
78
Pt
195.08
80
Hg
200.59
Mercury
Gold
Platinum
77
Iridium
Osmium
76
Os
190.2
Re
186.207
Rhenium
75
W
183.85
Ta
180.948
Tantalum
73
138.906
Hafnium
72
Lanthanum
57
La
Barium
56
Ba
Cesium
55
6
Cs
132.905
137.327
Hf
180.948
Tungsten
74
44
Ru
101.07
43
Tc
97.907
Ir
192.22
Silver
47
Ag
107.868
46
Pd
106.42
Ruthenium
Technetium
92.906
42
Mo
95.94
Zirconium
40
Zr
91.224
88.906
Yttrium
39
Y
Strontium
38
Sr
87.62
85.468
Rubidium
37
Rb
5
Ca
40.078
Niobium
41
Nb
Molybdenum
102.906
Rhodium
45
Rh
196.967
131.290
Xenon
54
Xe
Iodine
53
I
126.904
127.60
51
Sb
121.757
Tellurium
52
Te
Antimony
Tin
50
Sn
118.710
Indium
49
In
114.82
36
Kr
83.80
Palladium
48
Cd
112.411
Cadmium
Krypton
35
Br
79.904
As
74.922
Ge
72.61
78.96
Selenium
34
Se
Arsenic
33
Germanium
32
69.723
Gallium
31
Ga
Zinc
30
Cu
63.546
Zn
65.39
Copper
29
Nickel
28
Ni
58.693
Co
Mn
54.938
Cobalt
27
Iron
26
Fe
55.847
Manganese
25
Cr
51.996
23
V
50.942
47.88
Calcium
20
K
39.098
4
Potassium
19
Scandium
21
Sc
44.956
Titanium
22
Ti
Vanadium
Chromium
24
58.933
10
9
8
12
Mg
24.305
Basic Chemistry
Electron Shells in a Calcium Atom
7
Magnesium
Sodium
11
Na
22.990
3
Be
9.012
2
Li
6.941
2
Lithium
3
H
1.008
1
Hydrogen
1
Beryllium
4
3
4
Atomic mass
Symbol
Atomic Number
5
6
1.008
H
1
Hydrogen
Element
Groups
1
Bromine
Argon
18
Ar
39.948
17
Cl
35.453
P
30.974
Chlorine
Sulfur
16
S
32.066
Phosporous
15
Silicon
14
Si
28.086
Al
26.982
11
Nonmetal
Periodic Table of the Elements
Metal
12
Aluminum
13
10
Neon
Ne
20.180
O
15.999
N
14.007
10.811
C
12.011
F
18.998
He
4.003
17
Fluorine
9
16
Oxygen
8
15
Nitrogen
7
14
Carbon
6
13
Boron
5
B
18
Helium
2
Unit II
Periods
4-3 The periodic table helps researchers predict the chemical behavior of the elements. The elements outlined in red are those of greatest concern to dietitians.
nucleus
Note: The relative sizes of the particles and
shells in this model are inaccurate.
4-4 Water is a compound made up of the elements hydrogen and oxygen.
Formulas for carbon-based compounds
are usually written in a way that shows how
the atoms connect to the carbon. Thus, the
formula for acetic acid is usually written as
CH3COOH. Note how this formula is similar
to the arrangement of the atoms in the molecule.
Chemical Bonding
The lines in the diagram for an acetic acid
molecule represent chemical bonds. A chemical
bond is the force that holds two atoms together.
The subatomic particle that forms the bond is
the electron. Whether two elements will combine, and in what ratio, depends on their
electrons.
Electrons move in orbitals about the nucleus
of atoms in predictable patterns of space. An
area of space surrounding the nucleus that has
one or more orbitals is called a shell. (Shells
are referred to by their energy levels or by
their principle quantum number. This text will
use the more visual term shell.) The first shell
has one orbital and can hold two electrons.
The next shell can hold up to eight electrons in
four orbitals. See 4-5.
Atoms can have up to seven shells. The
number of shells in the atoms of an element
4-5 This model represents a calcium atom, which
has 20 electrons. These electrons are grouped in
four shells surrounding the nucleus.
determines the element’s period, or row, in the
periodic table. In other words, all the elements
in the third period have atoms with three
shells.
Atoms are most stable when the outer
shell of electrons is full. If the outer shell is one
electron short of being full, the atom will try to
gain an electron. If an atom has only one electron in its outer shell, the atom will try to give
away the electron. This will cause the shell to
be empty.
All the elements in a group, or column, of
the periodic table will react with other elements in similar ways. This is because all the
elements in a group have the same number of
electrons in their outermost shells. For
instance, each element in group 17 has one
electron missing in its outer shell. The elements in group 16 are missing two electrons
from their outer shells. Thus, an element’s
location in the periodic table helps to predict
how it will combine with other elements.
The element at the top of group 1 in the
periodic table is hydrogen. Hydrogen has one
proton in its nucleus and one electron moving
in a shell around the nucleus. All the other elements in group 1 also have only one electron
in their outermost shells. These elements all
bond readily with other elements.
84
Unit II
The elements in group 18 of the periodic
table have their outermost shells full of electrons. These elements are extremely stable. A
stable element is one that is least likely to form
chemical bonds.
Ionic Bonds
There are two basic types of chemical
bonds. One type is an ionic bond, in which the
electrons are transferred from one atom to
another. This causes both atoms to have a
charge. The atom receiving an electron will be
negatively charged, while the one that loses
the electron becomes positively charged.
Therefore, the ionic bond is a result of the
attraction between a positive charge and a
negative charge, 4-6.
Ionic Bonding
–
–
–
–
sodium
atom
+
–
+ +
+
+
+ + +
+ + +
–
–
–
–
–
–
–
–
–
–
–
–
–
–
+
+
+ + +
+
+ +
+ +
+ +
+ + +
+ +
–
chlorine
atom
–
–
–
–
–
–
–
–
4-6 When two atoms form an ionic bond, one
atom gives an electron to the other.
Basic Chemistry
Table salt is an example of a compound
with an ionic bond. Remember that salt is
made of Na and Cl. A sodium atom is electrically neutral. It has 11 positively charged protons and 11 negatively charged electrons. Its
outer shell contains only one electron. A chlorine atom is also electrically neutral. It has 17
positively charged protons and 17 negatively
charged electrons. Its outer shell is missing
one electron.
The sodium atom transfers the electron in
its outer shell to fill the outer shell of the chlorine atom. The sodium atom now has 11 positively charged protons and 10 negatively
charged electrons. This gives the sodium atom
a net positive charge. The chlorine atom that
received the electron now has 18 electrons and
17 protons. It has a negative charge.
An atom or group of atoms that has a positive or negative electrical charge is called an
ion. An ion with a positive electrical charge
has a superscript plus sign (+) written beside
the chemical symbol. An ion with a negative
electrical charge has a superscript minus sign
(–) written beside the chemical symbol. Thus,
the sodium ion is written as Na+, and the
chlorine ion is written as Cl–.
Like charges repel each other and opposite
charges attract. Therefore, sodium ions will
push away from other sodium ions and pull
toward chlorine ions. This will cause sodium
and chlorine atoms to form an alternating pattern or structure. This regular arrangement of
atoms causes a crystalline structure or crystal.
See 4-7.
Substances with ionic bonds will tend to
dissolve in water. This is because one end of
the water molecule is slightly positive, whereas the other end is slightly negative. The positive sodium ion is attracted to the negative
end of the water molecule. The negative chlorine ion is attracted to the positive end of the
water molecule. This results in the ionic bonds
being pulled apart. The salt crystal is now a
salt solution in water.
Sodium, like all elements except hydrogen
in group 1 of the periodic table, is a metal.
Chlorine, which is found in group 17 of the
periodic table, is a nonmetal. Ionic bonds form
between metals and nonmetals. Compounds
like table salt that result from ionic bonds are
Chapter 4
85
Basic Food Chemistry: The Nature of Matter
Salt Crystal
4-8 Calcium chloride, sodium chloride, and potassium chloride are all compounds formed by ionic
bonds between metals and chlorine, a nonmetal.
Na+
Cl–
4-7 Ionic bonding results in crystals. In a salt
crystal, each chlorine ion is surrounded by six
sodium ions.
crystalline in structure. The crystalline shape
is a result of the interaction between the negative and positive ions.
The metals potassium and calcium will
also combine with chlorine. Potassium chloride is the compound used to make no-salt
seasonings. Calcium chloride is used as a drying agent. See 4-8.
Storage Tip
Some compounds will easily react with
water. Because air contains water, dry powders, such as baking powder, need to be
stored in airtight containers.
Covalent Bonds
The second type of chemical bond, a
covalent bond, is formed when atoms share
one or more pairs of electrons. Water molecules are formed by covalent bonds between
two hydrogen atoms and one oxygen atom.
Hydrogen atoms have only one electron. It is
in a shell that can hold up to two electrons.
Oxygen atoms have eight electrons. There are
two electrons in the inner shell and six in the
outer shell, which can hold up to eight. An
oxygen atom will share one of the six electrons in its outer shell with each of two hydrogen atoms. At the same time, each of the
hydrogen atoms will share its electron with
the oxygen atom. Thus, both hydrogen atoms
will have a full shell of two electrons. The
oxygen atom will have a full shell of eight
electrons. See 4-9.
This is shown better with the Lewis structure. The Lewis structure is a shorthand
method of diagramming electrons that are
likely to be shared. The Lewis structure helps
scientists have a better picture of how atoms
combine. This system was developed by
Gilbert Newton Lewis, a chemist from the
United States. It is another tool to help predict
how elements will react.
The electrons that are likely to be shared or
transferred are in partially full shells and are
called valence electrons. In the Lewis structure, each valence electron is represented by a
dot next to the symbol for the element. A shell
can hold a maximum of eight electrons.
Therefore, the dots are arranged in pairs on
the four sides of the element symbol. The
water molecule described earlier would be
diagrammed in the following way:
H plus H plus O
yields O H
H
86
Unit II
Basic Chemistry
Chapter 4
87
Basic Food Chemistry: The Nature of Matter
Covalent Bonding
The Classification of Matter
Matter
+
–
Pure Substances
–
hydrogen atom
Compounds
–
–
+
Mixtures
+
+
+
–
+
+
–
+
Elements
• iron
• calcium
• zinc
+
+
Homogeneous
Mixtures
• vinegar
• coffee
• apple juice
Heterogeneous
Mixtures
• tossed salad
• vegetable soup
• Italian dressing
–
–
hydrogen atom
–
–
oxygen atom
Organic
Compounds
• sugar
• fats
• proteins
Inorganic
Compounds
• alum
• baking soda
• baking powder
• salt
4-9 Covalent bonds are formed when one atom shares an electron from its outer shell with another atom.
4-10 This chart illustrates how the various classifications of matter are related.
Another example of the Lewis structure is
shown below for methane (CH4). A molecule
of CH4 has four hydrogen atoms surrounding
a carbon atom. Hydrogen is always on the
outside or end of molecules. This is because
it has only one electron and can only form
one bond.
H
HCH
H
Matter can be classified into the two general categories of pure substances and mixtures. Each of these general categories can be
divided into two subcategories. Referring to
Chart 4-10 can help you see how these categories break down.
Pure Substances
The definition for covalent bonds states
that more than one pair of electrons can be
shared. When two atoms share two pairs of
electrons, a double bond is formed. An example of this is carbon dioxide, CO2.
C plus O plus O
The Classification of Matter
yields O C O
The number of valence electrons around
the atoms equals the number of valence electrons in the molecule. When shared pairs are
counted, each of the atoms has eight electrons
around it.
O C O
A pure substance is matter in which all the
basic units are the same. You can group pure
substances as elements and compounds.
Elements important in food science are those
needed for good health. These include iron,
calcium, and potassium. Most of the pure substances food scientists work with are compounds. Common examples include salt
(sodium chloride) and baking soda (sodium
bicarbonate).
Organic and Inorganic Compounds
Chemists divide the study of pure substances even further by dividing compounds
into two main groups. This method of classifying compounds is based on the source of the
compound in nature: living or nonliving substances. Scientists have discovered that all living substances contain carbon. Most nonliving
substances do not.
Chemists group compounds into two
main categories: organic and inorganic.
Organic compounds contain chains or rings of
carbon. Most organic compounds also contain
hydrogen and oxygen. All the sources of energy
in your diet are organic compounds. These
compounds (carbohydrates, fats, and proteins) are the main components of the food
you eat. Inorganic compounds either contain
no carbon or have only single carbon atoms.
Examples of inorganic compounds in foods
are table salt, water, and minerals.
Mixtures
Most of the substances you will work with
in the foods lab are not pure substances. They
are mixtures. Mixtures are substances that are
put together but not chemically combined.
Calcium is an element and table salt is a compound. Both are pure substances. However,
few people sit and eat pure calcium or pure
salt. A glass of milk contains both calcium and
salt. Milk is a mixture of these two substances
plus many others.
You can categorize mixtures as homogeneous or heterogeneous. A homogeneous mixture
has a uniform distribution of particles
throughout the sample. Visually, you cannot
tell one part of the mixture from another.
Examples of homogeneous mixtures are tea,
mayonnaise, and soft drinks. Homogenized
milk is also an example of a homogeneous
mixture.
A heterogeneous mixture has a nonuniform distribution of particles. A bowl of
vegetable soup is a heterogeneous mixture.
When you look at a spoonful of the soup,
you can see corn, beans, carrots, and onions
in a broth. If you pureed this soup in a
blender, it would become a homogeneous
mixture. See 4-11.
Sometimes, homogeneous and heterogeneous mixtures can be hard to tell apart. For
instance, at first glance, hot cocoa appears to
be a homogeneous mixture. However, when a
cup of cocoa sits without stirring, the heavier
cocoa molecules will gradually settle to the
bottom. Although cocoa looks uniform, the
88
Unit II
Basic Chemistry
sugar is the solute. Water is also the solvent in
plain coffee and fruit juice.
The substances in heterogeneous mixtures
can be separated by mechanical means. For
example, you could strain the vegetable soup
and then hand sort the vegetables. Separating
homogeneous mixtures is more difficult but
not impossible. For instance, salt water is a
homogeneous mixture. You cannot separate
the salt from the water by hand. However, if
you heat salt water, the water will turn to
steam and evaporate, leaving salt crystals
behind.
Physical and Chemical Changes
When scientists analyze what happens in
the world around them, they describe changes
they observe. Whether you are looking at
chemicals or food, you will observe two basic
kinds of changes. These are physical changes
and chemical changes. It is important to
understand how these changes differ to interpret what you see.
Physical Changes
4-11 Maple syrup, corn syrup, soft drinks, and
apple juice are examples of homogenous mixtures. Vegetable soup is an example of a heterogeneous mixture.
When you chop onions, the pieces become
smaller, but the substance is still onion. It has
the same color, flavor, and aroma as onion.
This is an example of a physical change.
Physical changes involve changing shape,
physical state, size, or temperature without
changing the chemical identity. If you freeze
water and then crush the ice, you still have
H2O. Melting ice is also a physical change.
Water goes from a solid state to a liquid state,
but it is still water. Dissolving salt in water is
another physical change. You can still taste the
salt and feel the wetness of the water.
Chapter 4
Basic Food Chemistry: The Nature of Matter
a definite shape and volume. Examples
include salt and ice.
As a solid is heated, the atoms and molecules begin to move farther apart as they gain
energy. As a result, solids lose their structure
and become liquid. The particles flow or slide
past one another. The substance has no definite shape of its own. It will take the shape of
the container. Liquids do have a definite volume. Examples are water, milk, and fruit juice.
As more heat is added to a liquid, the
atoms and/or molecules gain enough energy
to escape into the air. This represents a change
to the gas phase of matter. Gases will expand
to fit any closed container in which they are
stored. They have no definite shape or volume. See 4-12.
Any phase change is an example of a
reversible physical change. You can freeze
water to make ice—the solid phase of water.
When left at room temperature, the ice returns
to the liquid phase. If more heat is added, the
water boils and evaporates as steam—the gas
phase. Steam, or water vapor, returns to the
liquid state as it cools.
Chemical Changes
When bread is heated, it gets warmer. It
may also lose some water content. These are
physical changes. When bread is toasted,
browning occurs. The color and flavor change
because the starch molecules have undergone
a chemical change.
A chemical change occurs whenever new
substances with different chemical and physical properties are formed. Chemical changes
can produce changes in color or odor. Other
evidences of a chemical change include flavor
changes and the release of gas. Mixing water
and baking powder results in a chemical
change. The mixture will foam and fizzle as a
gas is released. Fermenting grapes is also an
example of a chemical change. When grape
juice is fermented to make wine, yeast
changes the sugar in the juice into alcohol.
Identifying Physical Versus Chemical
Changes
There are times when it will be difficult to
tell physical and chemical changes apart. For
example, when you open a carbonated beverage, it begins to fizz. This change is caused by
the carbon dioxide physically separating from
the water. If you combined baking soda and
an acid ingredient to make muffins, the batter
would begin to bubble. In this case, carbon
dioxide is being formed through a chemical
reaction. See 4-13. Chemical changes will often
States of Matter
Phase Changes
very small particles will not stay evenly distributed throughout the milk. Therefore, hot
cocoa is a heterogeneous mixture.
Most homogeneous mixtures are solutions. A solution is a homogeneous mixture of
one material dissolved in another. The material
that is dissolved is called the solute. The material that does the dissolving is the solvent. In
sweetened drinks, water is the solvent and
One example of a physical change is a
change in the phase or physical state of a substance. The phases or states of matter are
solids, liquids, and gases. A shift from one of
these states to another is called a phase
change. A phase change is a physical change in
the visible structure of matter without changing the molecular structure.
In the solid state, atoms and molecules are
close together in a rigid structure. Solids have
89
4-12 As a substance goes from the solid to the liquid to the gas state, the amount of space between
molecules increases.
90
Unit II
Basic Chemistry
Item of Interest
Non-Newtonian Fluids
You may think the term fluid
has the same meaning as the
term liquid. However, fluids are
substances that have characteristics of liquids and gases. A
unique type of fluid is a nonNewtonian fluid. Non-Newtonian
fluids have characteristics of liquids and solids. You can make a
substance that has these characteristics by combining two parts
cornstarch with one part water. If
you tried to pour this mixture
from a beaker, it would act like a
slow-moving liquid. However,
suppose you placed this mixture
on a lab table and struck it with
your fist. It would instantly resist
4-13 The fizz that forms on a soft drink is the
result of dissolved carbon dioxide physically separating from the soft drink. The fizz that forms from
combining baking soda and vinegar is the result
of a chemical change.
involve other noticeable changes in odor,
color, or taste.
Permanent and Reversible Changes
Physical and chemical changes are similar
in that they may or may not be reversible. You
can change fruit juice from liquid to solid back
the force and act like a solid. This
is why stirring a mixture of cornstarch and liquid is difficult
unless there are at least equal
amounts of both substances.
This is important to remember
when dissolving cornstarch in a
small amount of liquid before
adding it to a gravy or sauce.
to liquid and it is still fruit juice. That physical
change is reversible. However, you cannot
chop an onion and then put it back together
again. Although the substance is still onion, its
physical state has been permanently changed
into tiny pieces.
The way your body uses food provides an
example of a reversible chemical change. Your
body breaks down carbohydrates from your
diet into glucose, which can be used as a shortterm energy source. If you do not have any
short-term energy needs, your body can convert excess glucose into body fat for storage.
Later, when you need more energy, your body
can convert the body fat back into glucose.
Burning cookies illustrates a permanent
chemical change. If you severely overbake
cookies, they will turn black. You cannot
reverse this chemical change to make the
cookies edible again.
Chemical Equations
You can describe a chemical change
using a chemical equation. In a chemical equation, chemical formulas are used to represent
the compounds involved. Chemical formulas on the left side of the equation are called
Chapter 4
91
Basic Food Chemistry: The Nature of Matter
reactants. They are the substances that exist
before a chemical change takes place. A plus
symbol (+) is used to indicate that substances
are combined. An arrow represents a chemical
change, or reaction. One or more chemical formulas on the right side of the equation are the
products. They are the substances that are
formed.
When sodium hydroxide (NaOH) is
mixed with hydrochloric acid (HCl), a reaction will occur. This reaction will produce salt
(NaCl) and water (H2O). The chemical equation representing this reaction would be written as follows:
NaOH + HCl
NaCl + H2O
sodium hydroxide (lye) + hydrochloric acid yields salt + water
The law of conservation of matter states
that matter can be changed but not created or
destroyed. The conservation of matter is
shown in chemical equations. There must
always be the same number of atoms on the
right side as there are on the left. Although
elements can recombine to form new compounds, the atoms themselves will not
change. Therefore, there must also be the same
kind of atoms represented on the right as were
on the left. Look at the following equation:
C12H22O11 + O2
C12H22O11 + O2
12CO2 + 11H2O
12 Cs 22 Hs 13 Os
12 Cs 22 Hs 35 Os
As you can see, the carbon and hydrogen atoms
are in balance, but the oxygen atoms are not.
Twenty-two more oxygen atoms on the
left will balance the equation.
C12H22O11 + 12O2
12CO2 + 11H2O
Now the equation balances. For every
molecule of sugar, you need 12 molecules of
oxygen. This will result in 12 molecules of carbon dioxide and 11 molecules of water.
Careful measurements and repeated trials
of this reaction would show that this ratio is
constant. There will always be 12 molecules of
carbon dioxide for every molecule of sugar
digested. When chemical formulas are known,
scientists can use them like recipes or formulations to combine the exact amounts of ingredients needed. See 4-14.
Notice that in balancing the equation, the
arrangement and number of atoms in the
molecules was not changed. Only the number
of each kind of molecule can be changed when
balancing equations.
CO2 + H2O
This formula represents a simplified version of the digestion of sugar. The left side of
the equation has 12 carbon atoms and the
right only 1. This equation is not balanced.
There has not been a conservation of matter.
Every time a sugar molecule is digested, there
must be more than one molecule of CO2 made.
For the equation to balance, there must be 12
carbon atoms on each side.
C12H22O11 + O2
12CO2 + H2O
Now the carbon atoms balance, but the
oxygen atoms and hydrogen atoms still do
not. If there will be 11 molecules of water for
every molecule of sugar, the hydrogen atoms
will balance.
Agricultural Research Service, USDA
4-14 Food scientists use chemical formulas when
preparing formulations of food products.
Chapter 4
Review
Summary
A study of food science requires knowledge of the basic nature of matter. You need a
mental picture of how subatomic particles fit
into the structure of atoms. You need to know
that matter is made up of chemical elements,
which are identified in the periodic table. This
information will be your basis for understanding how atoms form ionic and covalent bonds
to create molecules and compounds.
Food scientists must be able to predict
how food products will perform during processing and preservation. Learning to classify
products as pure substances or mixtures will
help you make such predictions. As you
observe the behavior of food compounds, you
will need to identify physical and chemical
changes. Then you will use chemical equations as you record what you observe. This
course will help you see how food preparation
techniques relate to the chemical structure of
ingredients in foods.
Checking Your Understanding
1. Name the three subatomic particles. Give
the atomic mass, charge, and location in
the atom for each particle.
2. Describe the chemical symbols used to
represent elements.
3. What is the difference between the atomic
number and the atomic mass of an
element?
4. What does the 2 in the formula CO2 (carbon dioxide) indicate?
5. What is the most stable arrangement for
electrons?
6. How can an element’s location in the
periodic table help predict how the element will combine with other elements?
7. How do ionic and covalent bonds differ?
92
8. Identify the two categories of pure substances and give an example of each.
9. Give two examples of homogeneous mixtures and two examples of heterogeneous
mixtures.
10. Explain the main difference between
physical and chemical changes.
11. Compare the shape and volume of the
three states of matter.
12. How do you know that a chemical equation will always have the same number of
atoms of each type on each side?
Critical Thinking
1. What is the Lewis dot structure for each
4. Identify whether each of the following is
a physical or chemical change:
a. brewing tea
b. sweetening lemonade
c. browning pork chops
d. basting a turkey as it roasts
e. slicing tomatoes
f. caramelizing onions
g. cooking pancake batter
h. melting chocolate
i. simmering spaghetti sauce
5. List 10 food items in your refrigerator.
Identify whether they are compounds,
heterogeneous mixtures, or homogeneous
mixtures.
Explore Further
1. Science. Obtain a chemistry model kit or
materials from home, such as marshmallows and toothpicks. Use these materials
to construct models of each of the following molecules: water, acetic acid, and
methane.
2. Math. Mass 100 mL of sugar. Combine
the sugar with 100 mL of water. Water
has a mass of 1 g/mL. What is the combined mass? What is the combined
volume? Explain the results of conservation of mass and your understanding
of molecules.
of the following chemical formulas? Each
of these compounds is formed with covalent bonds. The elements are all
nonmetals.
a. C2H6
b. NH3
c. CH3COOH
2. Chlorine is found in column 17 of the
periodic table. The elements in column 17
are known as halogens. From your
knowledge of chlorine, what characteristics would you expect to find in the other
halogens?
3. Using the periodic table, identify how
many bonds with other atoms can be
formed by each of the following
elements:
a. calcium
b. fluorine
c. magnesium
d. nitrogen
e. phosphorus
f. potassium
g. sulfur
93
Experiment 4A
Physical Qualities of Food
Safety
O
O
O
Wear safety glasses when heating
glass.
Do not taste food samples.
Use hot pads or beaker tongs to move
hot glass beakers.
Purpose
All food products are made of chemical compounds. Each compound has measurable
characteristics, including boiling point, freezing point, color, aroma, and density. Becoming
familiar with these characteristics will help
you predict how ingredients will react in food
mixtures. In this experiment, you will examine the boiling points and densities of several
common food products.
Equipment
2 or 3 100-mL graduated cylinders
5 cups or bowls
1 or 2 250-mL beakers
3 150-mL beakers
thermometer
beaker tongs
Supplies
100 mL water
100 mL corn syrup
100 mL vegetable oil
100 mL cooking sherry
6 chocolate chips
100 mL rice
3 miniature marshmallows
3 ice cubes
1 drop food coloring
94
Procedure
1. Tare a 100-mL graduated cylinder. Use
this graduated cylinder to mass 100 mL of
water, corn syrup, vegetable oil, and
cooking sherry. Record measurements in
a data table. After massing each liquid,
pour it into a cup or bowl and set aside.
Be sure to wash and dry the graduated
cylinder after massing each liquid.
2. Mass the chocolate chips. Use rice to
measure the displacement volume of the
chocolate chips in a 100-mL graduated
cylinder. (Review how to measure the
volume of irregularly shaped objects from
Experiment 2C.) Record measurements in
the data table.
3. Mass the marshmallows. Use rice to
measure the displacement volume of the
marshmallows in a 100-mL graduated
cylinder. Record measurements in the
data table.
4. Working quickly to reduce melting, mass
the ice cubes. Measure the volume of the
ice cubes using the water displacement
method in a 250-mL beaker. Record measurements in the data table. Then place
the ice cubes in a cup or bowl and return
them to the freezer until they are needed.
5. Heat 50 mL of the water measured in
step 1 in one of the 150-mL beakers on a
hot plate or burner until it boils. Measure
and record the temperature. Discard the
heated water.
6. Heat 50 mL of the corn syrup in another
150-mL beaker until it boils. Measure and
record the temperature. Discard the heated
corn syrup and wash the thermometer.
7. Heat 50 mL of the cooking sherry in a
third 150-mL beaker until it boils.
Measure and record the temperature.
Discard the heated cooking sherry.
8. Using the formula density = mass ÷ volume,
calculate the density of the four liquids
and three solids.
9. Read steps 11 and 12 and predict what
will happen.
10. Add a drop of food coloring to the
remaining 50 mL of water and pour it
into the 250-mL beaker.
11. Slowly pour the remaining 50 mL of corn
syrup into the 250-mL beaker with the
colored water. Then slowly pour 50 mL of
the vegetable oil into the 250-mL beaker.
Finally, slowly pour the remaining 50 mL
of cooking sherry into the 250-mL beaker.
Observe what happens.
12. Add the chocolate chips, marshmallows,
and ice cubes to the liquids in the beaker.
Questions
1. Which of the liquids or solids is the
densest?
2. Which of the liquids or solids is the least
dense?
3. Was there a relationship between a
liquid’s density and its boiling point?
4. How accurate were the predictions you
made in step 9?
5. How can you apply this information to
food preparation?
95
Experiment 4B
Chemical Changes
Safety
O
O
O
Wear safety glasses when heating
glass.
Do not taste unknown chemicals unless
specifically directed to do so.
Use test-tube tongs to move the hot
watch glass.
Purpose
When a chemical change occurs in a mixture,
some signs of that change can be observed.
Some of the signs that a chemical change has
occurred include color changes, the forming of
a gas (bubbling or foaming), temperature
changes, and the forming of a precipitate (a
solid settling out of a solution).
Equipment
magnifying glass or microscope
100-mL beaker
glass stirring rod
2 watch glasses
test-tube tongs
Supplies
1 g sodium chloride
1 vitamin C tablet, crushed
2 g sodium bicarbonate
40 mL distilled water
Procedure
1. Use the magnifying glass or microscope
to examine a small amount of sodium
chloride. Describe the appearance in your
data table.
2. Taste a few crystals of sodium chloride.
Describe the taste in your data table.
96
Experiment 4C
The Chemical Detective
3. Repeat steps 1 and 2 twice, the first time
with the crushed vitamin C tablet and the
second time with the sodium bicarbonate.
4. Combine 1 g of sodium chloride, half of
the vitamin C tablet, and 20 mL of distilled water in the 100-mL beaker. Stir
with a glass rod until the sodium chloride
has dissolved.
5. Pour a small amount of the solution on a
watch glass. Discard the remaining solution and wash the beaker and glass rod
with warm, soapy water.
6. Heat the watch glass on a range over
medium heat until the liquid boils away.
7. Remove the watch glass from the heat
with the test-tube tongs and allow it to
cool.
8. Examine the residue on the watch glass
with a magnifying glass or microscope.
Taste the residue. Record your observations on appearance and flavor in your
data table.
9. Combine 2 g of sodium bicarbonate, the
remaining half of the crushed vitamin C
tablet, and the remaining 20 mL of distilled water in the cleaned 100-mL beaker.
Stir with the cleaned glass rod until the
sodium bicarbonate has dissolved.
10. Repeat steps 5 through 8, being sure to
use a clean watch glass in step 5.
Questions
1. What happened when water was added
to the powders?
2. Describe any physical changes that
occurred during this experiment.
3. Describe any chemical changes that
occurred during this experiment.
4. How can you apply information from this
experiment to food preparation?
Safety
O
Do not taste food samples.
O
Dispose of powders and indicators as
directed by your teacher.
Purpose
During this experiment, you will observe the
physical and chemical properties of four common household substances. You will be given
four unknown powders and three liquid indicators. You are to record any changes you see
after adding each indicator to each powder.
You will then test an unknown mixture of two
to four of the powders with the indicators and
try to determine which powders your mixture
contains. A good detective observes small
details.
Equipment
glass plate
wax pencil
small scoop
magnifying glass
Supplies
indicators I, II, and III in eye dropper bottles
powders A, B, C, and D in small portion cups
mystery powder
Procedure
Part I
1. Use the wax pencil to draw a threecolumn, four-row table on the glass plate.
Label the columns I, II, and III. Label the
rows A, B, C, and D.
2. Use the small scoop to place a pea-sized
sample of each powder in each cell of the
appropriate row of the table.
3. Use a magnifying glass to examine each
powder. Record your observations of its
appearance in the physical properties column of your data table.
4. Place one drop of Indicator I on each powder in column I. Record any reaction
observed in the column marked Indicator I
on the data table.
5. Repeat the procedure with Indicator II in
the second column and Indicator III in the
third column. Record observations.
6. Clean the glass plate.
Part II
1. Obtain a mystery powder from your
teacher.
2. Use the wax pencil to draw a threecolumn, two-row table on the glass plate.
Label the columns I, II, and III. Label the
rows mystery and test mix.
3. Place a pea-sized sample of the mystery
powder in each of the cells in the first
row of the table. Use a magnifying glass
to examine the mystery powder. Record
your observations of its appearance in the
mystery powder data table.
4. Place a drop of each indicator on the
mystery powder in the appropriate column. Record your observations in the
mystery powder data table. Hypothesize
which of powders A, B, C, and D are in
the mystery powder.
5. Combine equal amounts of the powders
you believe are in the mystery powder.
Place a pea-sized sample of the mixture
in each of the cells in the second row on
the glass plate. Use a magnifying glass to
examine the mixture. Compare the appearance with that of the mystery powder.
6. Place a drop of each indicator on the mixture in the appropriate column. Compare
the results with the results for the mystery powder.
97
Questions
1. Which test, if any, was the most helpful in
identifying the mystery powders? Explain
your answer.
2. What other tests could be done to make
identification of these powders easier or
more accurate?
98
3. What differences, if any, were there in the
mystery powder and the mix you prepared based on your hypothesis?
4. Name a career that would use these identification skills.