CHAPTER 4: MATTER
Of the “space”, “time”, “matter” and “energy” that
make up the physical world, “matter” is probably the
most obvious and most tangible. We can see, feel, touch
(and often smell) the “stuff” around us, particularly
when it is in a solid or a liquid state (more later about
“states” or “phases” of matter).
As we will see, matter includes not only the concept of
mass, but also electric charge. On a subatomic level,
there are other distinguishing characteristics as well (e.g.
“spin” and “color” which will be mentioned briefly
later).
When the term “atom” (meaning indivisible) was first
introduced in ancient Greece, it was believed to be the
smallest possible unit of matter. Only much later (in the
early 1900s), was it realized that atoms themselves
consist almost entirely of empty space, with a tiny, dense
nucleus of protons and neutrons orbited by even tinier
electrons. It was in discovering and understanding this
“atomic structure” that the branch of physical science
known today as chemistry really began to emerge, and it
is on this atomic level that we will focus most of our
discussion of matter in this course.
It must be recognized, however, that in the last 40-50
years, a multitude of new discoveries have altered and
continue to challenge our understanding of matter in its
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most fundamental forms. Even protons and neutrons,
we now realize, are formed from other, even tinier
particles. In a strange way, this search for the tiniest bits
of matter is also a quest to understand the origins of the
entire universe itself. That is, the biggest question of all
has a great deal to do with understanding the most basic
constituents of matter. A little later, we will briefly
discuss the most recent, revised view of matter.
STATES (OR PHASES) OF MATTER
Everyone has heard at some point that matter can exists
as a solid, a liquid or a gas. Conventionally speaking,
this is true, but we must also include a fourth state (or
phase), namely a plasma, which exists in more extreme
conditions of heat such as is found in the core of the Sun
or other stars.
Solids
In a solid, the atoms or molecules are (relatively) closelypacked, and arranged in a lattice structure. The atoms
still move, but their motions (largely vibrational) are
extremely limited. Solids have a high resistance to
compression and to being pulled apart (though less
resistance to the latter), which is why they retain their
shape, and they only expand or contract very slightly
when temperatures increase or decrease.
The other three states are all considered fluids, meaning
they have the ability to “flow”.
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Liquids
In a liquid, the atoms or molecules are much less closely
packed, and are not held within a lattice. As a result,
they are much more free to move. Liquids also strongly
resist compression, but they offer little if any resistance
to being pulled apart. Consequently, liquids readily take
on the shape of their containers. Like solids, liquids
expand or contract only a little (though slightly more
than solids do) as a result of temperature changes.
Gases
The atoms or molecules of a gas are restricted in their
motion only by the limitations of their containers. So,
gases have indefinite volumes and can have almost
infinitely variable densities. Gases will normally expand
to fill their containers evenly (at least “locally”),
regardless of shape, and unlike liquids or solids, they are
readily compressible. For gases, it turns out that volume
V, temperature T and pressure P are interrelated (by
what has come to be called the gas law, PV/T = constant),
and hence, for example, a change in volume can be
accompanied by a change in pressure or temperature, or
both.
Plasmas
At extremely high temperature, molecules and atoms
break apart into a gas-like ionized “soup” of charged
particles (mostly atomic nuclei) which are not only much
more energetic than in ordinary gases, but, because they
are charged, act differently than in ordinary gases. For
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this reason, plasmas have little or no relevance or
application in conventional chemistry. (Plasmas are still
critically important, though, since they explain much
about the life and death of stars, and hence the about the
long-term evolution of the universe.)
For now, we will focus on matter as it is usually
categorized by chemists, and thus concentrate on solids,
liquids and gases.
THE CHEMICAL DIVISION OF MATTER
Chemists agree that matter has mass (and sometimes
charge) and occupies space. Beyond that, they divide
matter into two primary categories: pure substances and
mixtures.
Pure substances have a definite known composition and
definite properties, but their constituent “parts” or
components can’t be physically separated. In turn, pure
substances can consist either of: (i) single chemical
elements (e.g. hydrogen, carbon, oxygen) which also
can’t be decomposed by “chemical reactions”; or (ii)
chemical compounds (e.g. water, carbon dioxide, sodium
chloride) which can be decomposed into their constituent
elements by “chemical reactions”. (For more on
“chemical reactions”, see below.)
Mixtures contain two or more pure substances which can
be separated by physical means. Mixtures may be either
homogeneous or heterogeneous. In homogeneous
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mixtures (e.g. seawater, air, bronze), the substances mix
uniformly, while the substances which make up
heterogeneous mixtures (e.g. pancake mix, concrete) are
mixed unevenly or non-uniformly.
As you probably know, pure substances can often exist
in two or more states (solid, liquid or gas) at different
temperatures, with water being especially important and
interesting to us since its “freezing” and “boiling” points
are so close to one another (0 C and 100 C), and by
virtue of its being able to occupy all three states
simultaneously, and/or to move directly from solid to gas
(called sublimation) without an intermediate liquid state.
Chemical reactions (a term introduced above) include
reactions in which chemical compounds decompose into
their constituent elements, or (somewhat in reverse)
reactions in which elements or compounds combine to
form other, different compounds. For example, water
can be decomposed into hydrogen and oxygen, while
methane and oxygen gases can combine (during
combustion) to produce carbon dioxide and water
(vapour).
Everything in conventional chemistry begins with basic
chemical elements. Moreover, each chemical element in
turn possesses a different atomic structure (which is
presumably why each has unique and different
properties). So, let us turn to the atomic structure of
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matter and the notation we use in identifying chemical
elements and their place in the so-called periodic table.
CHEMICAL NOTATION
Each element is named with a 1- or 2-letter symbol (e.g.
H-hydrogen, O-oxygen, Na-sodium, Cl-chlorine).
Each element has a unique atomic number (denoted Z)
which is the number of protons in its atomic nucleus.
The total number of protons and neutrons in the nucleus
is called its mass number (denoted A). An element’s mass
number may not be unique, however. The same element
(e.g. carbon with atomic number 6) can exist with several
different mass numbers (or numbers of neutrons). These
different forms of an element are called isotopes, and
they will typically differ in terms of their stability (and
hence their abundance).
A
The standard chemical notation for an element is Z
X
,
i.e. the mass number and atomic number appear as a
superscript and a subscript, respectively, immediately to
the left of the symbol for the element.
Examples
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1
27
H ,12H , 24He,126C ,147N ,168O,13
Al ,197
79 Au