The GED Science Exam - Physical Science


The Structure of Atoms
You and everything around you are composed of tiny particles called atoms. The book you are reading, the neu-
rons in your brain, and the air you are breathing can all be described as a collection of various atoms.
History of the Atom
The term atom, which means indivisible, was coined by Greek philosopher Democritus (460–370
B
.
C
.). He dis-
agreed with Plato and Aristotle—who believed that matter could infinitely be divided into smaller and smaller
pieces—and postulated that matter is composed of tiny indivisible particles. In spite of Democritus, the belief that
matter could be infinitely divided lingered until the early 1800s, when John Dalton formulated a meaningful
atomic theory. It stated:
■
Matter is composed of atoms.
■
All atoms of a given element are identical.
■
Atoms of different elements are different and have different properties.
■
Atoms are neither created nor destroyed in a chemical reaction.
CHAPTER
Physical Science
PHYSICAL SCIENCE includes the disciplines of chemistry (the
study of matter) and physics (the study of energy and how energy
affects matter). The questions on the physical science section of the
GED will cover topics taught in high school chemistry and physics
courses. This chapter reviews the basic concepts of physical-
science—the structure of atoms, the structure and properties of mat-
ter, chemical reactions, motions and forces, conservation of energy,

increase in disorder, and interactions of energy and matter.
23
223
■
Compounds are formed when atoms of more
than one element combine.
■
A given compound always has the same relative
number and kind of atoms.
These postulates remain at the core of physical science
today, and we will explore them in more detail in the fol-
lowing sections.
Protons, Neutrons, and Electrons
An atom is the smallest unit of matter that has the prop-
erties of a chemical element. It consists of a nucleus sur-
rounded by electrons. The nucleus contains positively
charged particles called protons, and uncharged neu-
trons. Each neutron and each proton has a mass of about
1 atomic mass unit, abbreviated amu. An amu is equiv-
alent to about 1.66 × 10
−24
g. The number of protons in
an element is called the atomic number. Electrons are
negatively charged and orbit the nucleus in electron
shells.
Electrons in the outermost shell are called valence
electrons. Valence electrons are mostly responsible for
the properties and reaction patterns of an element. The
mass of an electron is more than 1,800 times smaller
than the mass of a proton or a neutron. When calculat-

ing atomic mass, the mass of electrons can safely be neg-
lected. In a neutral atom, the number of protons and
electrons is equal. The negatively charged electrons are
attracted to the positively charged nucleus. This attrac-
tive force holds an atom together. The nucleus is held
together by strong nuclear forces.
A representation of a lithium atom (Li). It has 3 protons (p)
and 4 neutrons (n) in the nucleus, and 3 electrons (e) in the
two electron shells. Its atomic number is 3 (p). Its atomic
mass is 7 amu (p + n). The atom has no net charge because
the number of positively charged protons equals the number
of negatively charged electrons.
Charges and Masses
of Atomic Particles
Proton Neutron Electron
Charge +1 0 –1
Mass 1 amu 1 amu
ᎏ
18
1
00
ᎏ
amu
Isotopes
The number of protons in an element is always the same.
In fact, the number of protons is what defines an ele-
ment. However, the number of neutrons in the atomic
nucleus, and thus the atomic weight, can vary. Atoms
that contain the same number of protons and electrons,
but a different number of neutrons, are called isotopes.

The atomic masses of elements in the periodic table are
weighted averages for different isotopes. This explains
why the atomic mass (the number of protons plus the
number of neutrons) is not a whole number. For exam-
ple, most carbon atoms have 6 protons and 6 neutrons,
giving it a mass of 12 amu. This isotope of carbon is
called “carbon twelve” (carbon-12). But the atomic mass
of carbon in the periodic table is listed as 12.011. The
mass is not simply 12, because other isotopes of carbon
have 5, 7, or 8 neutrons, and all the isotopes and their
abundance are considered when the average atomic mass
is reported.
Ions
An atom can lose or gain electrons and become charged.
An atom that has lost or gained one or more electrons is
called an ion. If an atom loses an electron, it becomes a
positively charged ion. If it gains an electron, it becomes
a negatively charged ion. For example, calcium (Ca), a
biologically important element, can lose two electrons to
become an ion with a positive charge of +2 (Ca
2+
). Chlo-
rine (Cl) can gain an electron to become an ion with a
negative charge of −1 (Cl
−
).
The Periodic Table
The periodic table is an organized list of all known ele-
ments, arranged in order of increasing atomic number,
such that elements with the same number of valence

electrons, and therefore similar chemical properties, are
found in the same column, or group. For example, the
last column in the periodic table lists the inert (noble)
gases, such as helium and neon—highly unreactive ele-
ments. A row in the periodic table is called a period.
3 p
4 n
e
e
e
Nucleus
Electron
shells
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224
Elements that share the same period have the same num-
ber of electron shells.
Common Elements
Some elements are frequently encountered in biologi-
cally important molecules and everyday life. Below you
will find a list of common elements, their symbols, and
common uses.
H—Hydrogen: involved in the nuclear process that
produces energy in the sun
He—Helium: used to make balloons fly
C—Carbon: found in all living organisms; pure car-
bon exists as graphite and diamonds
N—Nitrogen: used as a coolant to rapidly freeze

food
O—Oxygen: essential for respiration (breathing)
and combustion (burning)
Si—Silicon: used in making transistors and solar
cells
Cl—Chlorine: used as a disinfectant in pools and as
a cleaning agent in bleach
Ca—Calcium: necessary for bone formation
Fe—Iron: used as a building material; carries oxygen
in the blood
Cu—Copper: a U.S. penny is made of copper; good
conductor of electricity
I—Iodine: lack in the diet results in an enlarged thy-
roid gland, or goiter
Hg—Mercury: used in thermometers; ingestion can
cause brain damage and poisoning
Pb—Lead: used for X-ray shielding in a dentist
office
Some elements exist in diatomic form (two atoms of
such an element are bonded), and are technically mole-
cules. These elements include hydrogen (H
2
), nitrogen
(N
2
), oxygen (O
2
), fluorine (F
2
), chlorine (Cl

2
), bromine
(Br
2
), and iodine (I
2
).

Structure and Properties
of Matter
Matter has weight and takes up space. The building
blocks of matter are atoms and molecules. Matter can
interact with other matter and with energy. These inter-
actions form the basis of chemical and physical
reactions.
Molecules
Molecules are composed of two or more atoms. Atoms
are held together in molecules by chemical bonds.
Chemical bonds can be ionic or covalent. Ionic bonds
form when one atom donates one or more electrons to
another. Covalent bonds form when electrons are shared
between atoms. The mass of a molecule can be calculated
by adding the masses of the atoms of which it is com-
posed. The number of atoms of a given element in a
molecule is designated in a chemical formula by a sub-
script after the symbol for that element. For example, the
glucose (blood sugar) molecule is represented as
C
6
H

12
O
6.
This formula tells you that the glucose mole-
cule is contains six carbon atoms (C), twelve hydrogen
atoms (H), and six oxygen atoms (O).
Organic and Inorganic Molecules
Molecules are often classified as organic or inorganic.
Organic molecules are those that contain both carbon
and hydrogen. Examples of organic compounds are
methane (natural gas, CH
4
), glycine (an amino acid,
NH
2
CH
2
COOH), and ethanol (an alcohol, C
2
H
5
OH).
Inorganic compounds include sodium chloride (table
salt, NaCl), carbon dioxide (CO
2
), and water (H
2
O).
States of Matter
Matter is held together by intermolecular forces—forces

between different molecules. Three common states of
matter are solid, liquid, and gas. Matter is an atom, a
molecule (compound), or a mixture. Examples of mat-
ter in solid form are diamonds (carbon atoms), ice
(water molecules), and metal alloys (mixtures of differ-
ent metals). A solid has a fixed shape and a fixed volume.
The molecules in a solid have a regular, ordered arrange-
ment and vibrate in place, but are unable to move far.
Examples of matter in liquid form are mercury (mer-
cury atoms), vinegar (molecules of acetic acid), and per-
fume (a mixture of liquids made of different molecules).
Liquids have a fixed volume, but take the shape of the
container they are in. Liquids flow, and their density
(mass per unit volume) is usually lower than the density
of solids. The molecules in a liquid are not ordered and
can move by sliding past one another through a process
called diffusion.
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Examples of matter in gaseous form include helium
gas used in balloons (helium atoms), water vapor (mol-
ecules of water), and air (mixture of different molecules
including nitrogen, oxygen, carbon dioxide, and water).
Gases take the shape and volume of their container. They
can be compressed when pressure is applied. The mole-
cules in gases are completely disordered and move very
quickly. Gas density is much lower than the density of a
liquid.

Phase Changes
Change of phase involves the transition from one state of
matter into another. Freezing water to make ice for cool-
ing your drink, condensation of water vapor as morning
dew, and sublimation of dry ice (CO
2
) are examples of
phase change. A phase change is a physical process. No
chemical bonds are formed or broken. Only the inter-
molecular (physical) forces are affected.
Freezing is the process of changing a liquid into a solid
by removing heat. The opposite process whereby heat
energy is added to the solid until it changes into a liquid
is called melting. Boiling is the change of phase from a
liquid to a gas and also requires the input of energy. Con-
densation is the change from gas to liquid. Some sub-
stances sublime—change directly from the solid phase to
the gas phase without forming the liquid state first. Car-
bon dioxide is such a substance. Solid carbon dioxide,
called dry ice, evaporates into the gas phase when heated.
When gas changes directly into a solid, the process is
called deposition.
Phase changes between the three states of matter
The stronger the intermolecular forces are, the easier
it is for the molecule to exist in one of the condensed
states (liquid or gas). Molecules in which intermolecular
forces are strong tend to have high boiling points, since
more energy is needed to turn the molecules into the
gaseous state where molecular interactions are low.
Compounds and Mixtures

A compound is a homogeneous substance composed of
two or more elements, united chemically. Examples of
compounds include carbon dioxide (a product of respi-
ration), sucrose (table sugar), seratonin (a human brain
chemical), and acetic acid (a component of vinegar). In
each of these compounds, there is more than one type of
atom, chemically bonded to other atoms in definite pro-
portion. Compounds are made of molecules.
A mixture is a physical combination of its compo-
nents. In a homogeneous mixture, the components can’t
be visually separated. Homogeneous mixtures also have
the same composition (ratio of components) through-
out their volume. An example is a mixture of a small
amount of salt in water. A uniform mixture is often
called a solution. In a solution, one substance (solute) is
dissolved in another (solvent). In the salt and water mix-
ture, the salt is the solute, and the water is the solvent. In
a heterogeneous mixture, the components can often be
visually identified, and the composition may vary from
one point of the mixture to another. A collection of
dimes and pennies is a heterogeneous mixture. A mix-
ture of sugar and flour is also heterogeneous. While both
components (sugar and flour) are white, the sugar crys-
tals are larger and can be identified.
Miscibility is the term used to describe the ability of
two substances to form a homogeneous mixture. Water
and alcohol are miscible. They can be mixed in such a
way that the mixture will be uniform throughout the
sample. At each point, it will look, smell, and taste the
same. Oil and water are not miscible. A mixture of these

two substances is not homogeneous, since the oil floats
on water. In a mixture of oil and water, two layers con-
taining the two components are clearly visible. Each layer
looks, smells, tastes, and behaves differently.
Gas
Liquid
Vaporization
Solid
Sublimation Deposition
Condensation
Melting
Freezing
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Chemical Reactions
Removing stains from clothes, digesting food, and
burning wood in a fireplace are all examples of chemical
reactions. Chemical reactions involve changes in the
chemical arrangement of atoms. In a chemical reaction,
the atoms of reactants combine, recombine, or dissociate
to form products. The number of atoms of a particular
element remains the same before and after a chemical
reaction. The total mass is also preserved. Similarly,
energy is never created or destroyed by a chemical
reaction. If chemical bonds are broken, energy from
those bonds can be liberated into the surroundings as
heat. However, this liberation of energy does not consti-

tute creation, since the energy only changes form—from
chemical to heat.
Writing Chemical Reactions
A chemical reaction can be represented by a chemical
equation; the reactants are written on the left side and
the products on the right side of an arrow, indicating the
direction in which the reaction proceeds. The chemical
equation below represents the reaction of glucose
(C
6
H
12
O
6
) with oxygen (O
2
) to form carbon dioxide
(CO
2
) and water (H
2
O). Your body runs this reaction all
the time to obtain energy.
(C
6
H
12
O
6
) + 6 (O

2
) → 6 (CO
2
) + 6 (H
2
O)
The numbers in front of the molecular formulas indi-
cate the proportion in which the molecules react. No
number in front of the molecule means that one mole-
cule of that substance is reacting. In the reaction above,
one molecule of glucose is reacting with six molecules of
oxygen to form six molecules of carbon dioxide and six
molecules of water. In reality, there are many molecules
of each of the substances and the reaction tells you in
what proportion the molecules react. So if you had ten
molecules of glucose react with 60 molecules of oxygen,
you would obtain 60 molecules of carbon dioxide and 60
molecules of water. In many ways, chemical equations
are like food recipes.
2 Bread + 1 Cheese + 2 Tomato → Sandwich
With two slices of bread, one slice of cheese, and two
slices of tomato, you can make one sandwich. If you had
six slices of bread, three slices of cheese, and six slices of
tomato, you could make three sandwiches. The same
principles of proportion apply in chemical reactions.
Heat of Reaction (Enthalpy)
Breaking molecular bonds releases energy stored in those
bonds. The energy is released in the form of heat. Simi-
larly, forming new bonds requires an input of energy.
Therefore, a chemical reaction will either absorb or give

off heat, depending on how many and what kind of
bonds are broken and made as a result of that reaction. A
reaction that absorbs energy is called endothermic. A
container in which an endothermic reaction takes place
gets cold, because the heat of the container is absorbed by
the reaction. A reaction that gives off energy is called
exothermic. Burning gasoline is an exothermic
reaction—it gives off energy.
Increase in Disorder (Entropy)
Disorder, or entropy, is the lack of regularity in a system.
The more disordered a system, the larger its entropy. Dis-
order is much easier to come by than order. Imagine that
you have 100 blue beads in one hand and 100 red beads
in the other. Now place all of them in a cup and shake.
What are the chances that you can pick out 100 beads in
each hand so that they are separated by color, without
looking? Not very likely! Entropy and chaos win. There
is only one arrangement that leads to the ordered sepa-
ration of beads (100 blue in one hand, 100 red in the
other), and many arrangements that lead to mixed-up
beads (33 blue, 67 red in one hand, 33 red and 67 blue in
the other; 40 blue, 60 red in one hand, 60 blue, 40 red in
the other . . .). The same is true of atoms. Sometimes,
arrangement and order can be achieved. Atoms and mol-
ecules in solids, such as snowflakes, have very regular,
ordered arrangements. But given enough time (and tem-
perature), the snow melts, forming less ordered liquid
water. So, although reactions that lead to a more ordered
state are possible, the reactions that lead to disorder are
more likely. The overall effect is that the disorder in the

universe keeps increasing.
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