Ochiai E. Chemicals for Life and Living

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222

18 Domain of Chemistry

There is one other important sub-discipline of chemistry; that is, analytical

chemistry. You would need to identify and quantify the material you deal with; i.e.,

chemical analysis. Analytical chemistry borrows ideas from all branches of chemis-

try, particularly physical chemistry, and devises means to do analysis. It is an applied

ﬁeld of chemistry. In real life, analytical chemistry is, perhaps, the most-often-used

chemistry. Depending on the material dealt with, you can have bioanalytical chem-

istry, environmental analytical chemistry and so forth. Another way of subdividing

analytical chemistry is based on the means used: “wet chemical analysis,” typically

dealing with solutions (hence wet) and using more traditional analytical methods,

and “instrumental analysis” which uses a number of different instruments to do

analysis: spectrophotometric, electrochemical, etc.

An entirely different kind of chemistry sub-discipline is nuclear chemistry.

It deals with chemicals all right, but its concern is quite different from all the

sub-disciplines mentioned above. It studies the nuclei of atoms in chemicals. The

nuclei obey quite different kinds of rules than the ordinary chemicals do, as the next

Chap. (19) explains. It treats the radioactivity, nucleosynthesis (how elements are

produced), nuclear ﬁssion and fusion, and extends to cosmochemistry.

223

E. Ochiai, Chemicals for Life and Living,

DOI 10.1007/978-3-642-20273-5_19, © Springer-Verlag Berlin Heidelberg 2011

This chapter is not meant to be a comprehensive, detailed, and systematic presentation

of all chemical concepts/principles and their applications. It tries to present a bare

minimum of the essential chemical concepts that may be necessary to appreciate the

content of this book.

One of the essentials discussed is the concept of Avogadro number, but its appli-

cation, i.e., the quantitative calculations involving “moles,” is hardly discussed in

the main body of the book. The author considers that grasping of such quantitative

calculations is unnecessary except for the practitioners of chemistry and the related

areas. The inclusion of this important concept is to only emphasize the nature of

material world; i.e., it consists of particles, atoms, ions, molecules, etc. And the

quantities of any material is fundamentally measured in terms of the number of the

particles, though in everyday life the quantities may be measured in terms of mass

(weight), volume, length, etc. Occasionally, such calculations do appear in the text

for sure, but it is only meant to be a reminder to those who understand or need

such a calculation. In general, quantitative calculations that appear in the text are

meant to give the readers a sort of feeling toward the quantities talked about. They

are not intended to be precise and detailed and often are not shown of the details of

assumptions, calculations, and data.

Other major omissions are the concept of atomic orbital hybridization vs shape

of molecule and the determination of the molecular structures. These topics would

require a lengthy discourse to do full justice to the subject. We entirely omit these

aspects and present all the chemical structures as given.

19.1 Atoms/Molecules/Mole/Avogadro Number

According to chemistry, a substance is an aggregate of a very large number of very

tiny particles that can be a kind of atom or a kind of molecule. A substance that is

made of a single kind of atom is called an “element,” while a substance consisting

Chemistry’s View of the Material World:

Basic Principles

224

19 Chemistry’s View of the Material World: Basic Principles

of a large number of a same kind of molecule or the same combination of different

atoms is a “compound.” We are using terms such as “atom,” “element,” and “mole-

cule” here without deﬁning them. We are relying on your knowledge of these terms;

it does not matter at this stage whether your own deﬁnition is precise or not. We will

eventually ﬁnd out exactly what they are.

Now take an ounce (28.35 g) of water. It turns out to consist of approximately

1000000000000000000000000 (there are 24 “0”s) molecules of water. [By the way,

300000 (5 “0”s) is often expressed as 3 × 10

; so the large number above with 24

“0”s can be written as 1 × 10

for short. This kind of expression is called “scientiﬁc

notation”]. A single molecule of water represents the smallest unit that has the char-

acter of water, and hence, weighs only about 3/10

g (i.e., 28.35 g (i.e.,1 oz)/10

)

[As 1/10

, i.e., 1/1000 one-thousandth, is written as 1 × 10

−3

, 3/10

can be written as

3 × 10

−23

and is a very tiny value]. When chemists look at water, they think in terms

of literally a sea of these very tiny molecules of water. They are moving around

randomly within a conﬁne at ambient temperatures and pressures, and thus the

aggregate of water molecules show a character of liquid at ordinary temperatures.

Some of the water molecules go out of the liquid into the surrounding space and

come back. All of the molecules there always keep going out and coming back, but

overall, the majority of water remains in the liquid. However, if molecules that

escaped into the space would not come (as in the case where water is in an open

container), eventually all water molecules disappear into space.

Let us now go to a little deeper level. That is, the idea of “atom.” The atoms

constitute molecules and hence are the basis of all materials. For example, a water

molecule has been shown to consist of two hydrogen atoms and an oxygen atom.

The idea of “atom” is as old as ancient Greece, where several philosophers sug-

gested that the material is made of atoms. Atoms are the smallest units of matter that

cannot be divided further. Democritus coined the word “atom” to mean what was

just said. He said: “The ﬁrst principles of the universe are atoms and empty space;

everything else is merely thought to exist.” There are different kinds of atom. An

aggregate of atoms of the same nature is called an “element”; so atoms of different

natures belong to different elements. Some Greek philosophers thought that the

whole world was made of four elements: for example, air, water, ﬁre, and soil. We

know now that these are not elements in the modern sense.

Today, science recognizes about 100 different elements in the whole universe.

“Elements” should be understood as “chemical elements” to distinguish it from

“elements” in ordinary usage. Everything that is present in this universe is made of

one or more atoms of these 100 or so elements. A water molecule is made of two

atoms of element hydrogen (symbol “H”) and one atom of element oxygen (symbol

“O”). And this fact is expressed as H

O. Chemists also have learned that in water

molecule, two hydrogen atoms are bound with the one oxygen atom as shown in

Fig. 19.1. This is the so-called structure of water molecule. In real life we would not

be able to deal with a single molecule, for it is too tiny. Instead we will be dealing

with, say, a gallon of water. A gallon of water contains a very large number of water

molecules as we told you earlier. But the properties of this large body of water or a

22519.1 Atoms/Molecules/Mole/Avogadro Number

small amount of water, say a drop of water, are determined by the molecule having

O composition; therefore, we will use “H

O” notation even when we talk about

this large number of water molecules, i.e., the substance or compound water. This is

the chemical formula of water (molecule).

Iron is another element (chemical symbol = Fe). Metal iron, say 100 g, is made of

a very large number (about 10

) of atoms of element iron. When iron rusts, it turns

into a substance called iron oxide, which may be described as consisting of two

atoms of element iron and three atoms of element oxygen (Fe

). (It might be

pointed out here, though it may be confusing at this stage, that there is no such

molecular entity having the composition Fe

. It shows only the composition).

Element carbon (chemical symbol = C) is made of atoms of a single element

carbon, but can take different shapes. All carbon atoms in these substances are

bound to each other, but their arrangements can be different. As a result, graphite

and diamond, though consisting of only carbon atoms, i.e., being elemental carbon,

have such disparate properties. Carbon atom(s) make million different compounds

(molecules), binding hydrogen atom(s), oxygen(s) and nitrogen(s, chemical

symbol = N); they are collectively called “organic compounds.” Some of them such

as proteins and DNA constitute our body.

All the elements are now arranged neatly in the so-called periodic table, which

consists of 18 columns and 7 rows plus 2 extra segments (Fig. 19.2). Each element

is given a number (atomic number) and then pigeonholed in each of the boxes.

Chemistry deals with all these elements, their compounds, and how they behave.

Remember that all the materials that exist are made of these elements.

As we said before, all the atoms and molecules are miniscule and it is very hard

to deal with a single atom or a single molecule. Therefore, chemists invented a unit

of measurable quantity of elements or compounds (substances).

It is called “mole.” One mole of a molecular substance is deﬁned to be a collec-

tion of approximately 6 × 10

molecules. This huge number 6 × 10

(23 0s after 6)

is called “Avogadro Number.” Avogadro is an Italian chemist who attempted to

determine this number ﬁrst. The number is frighteningly large, but the idea of

“mole” is the same as, say, “dozen,” which consists of 12 bodies. So 1 mol of water

consists of 6 × 10

particles of water molecule H

O. As quarter a dozen is equal to

12/4 = 3, quarter a mole is 6 × 10

/4 = 1.5 × 10

. By the way, no signiﬁcance should

be attached to the fact that these two quantities 12 and 6 (×10

) appear somehow

related. This is a pure coincidence. “6” in the quoted Avogadro number is an approx-

imation; it has been estimated to be “6.022,” to be more precise. This value itself is

still an approximation, but practically accurate enough to be useful.

Fig. 19.1 Structure of water

molecule (often this is further

simpliﬁed as H–O–H)

226

19 Chemistry’s View of the Material World: Basic Principles

The Avogadro number has been decided in such a way so that 1 mol of a

substance, say, water, weighs an ordinary everyday number, such as 18 g in this

case. It turned out that 1 mol of many elements comes close to an integer number

if we use the Avogadro number of about 6 × 10

; for example, 1 mol of hydrogen

Fig. 19.2 Periodic Chart. (Ac, actinium; Ag, silver; Al, aluminum; Am, americium; Ar, argon;

As, arsenic; At, astatine; Au, gold; B, boron; Ba, barium; Be, beryllium; Bi, bismuth; Bk, berke-

lium; Br, bromine; C, carbon; Ca, calcium; Cd, cadmium; Ce, cerium; Cf, californium; Cl, chlo-

rine; Cm, curium; Co, cobalt; Cr, chromium; Cs, cesium; Cu, copper; Dy, dysprosium; Er, erbium;

Es, einsteinium; Eu, europium; F, ﬂuorine; Fe, iron; Fm, fermium; Fr, francium; Ga, gallium; Gd,

gadolinium; Ge, germanium; H, hydrogen; He, helium; Hf, hafnium; Hg, mercury; Ho, holmium;

I, iodine; In, indium; Ir, iridium; K, potassium; Kr, krypton; La, lanthanum; Li, lithium; Lr, law-

rencium; Lu, lutetium; Md, mendelevium; Mg, magnesium; Mn, manganese; Mo, molybdenum;

N, nitrogen; Na, sodium; Nb, niobium; Nd, neodymium; Ne, neon; Ni, nickel; No, nobelium; Np,

neptunium; O, oxygen; Os, osmium; P, phosphorus; Pa, protactinium; Pb, lead; Pd, palladium; Pm,

promethium; Po, polonium; Pr, praseodymium; Pt, platinum; Pu, plutonium; Ra, radium; Rb,

rubidium; Re, rhenium; Rh, rhodium; Rn, radon; Ru, ruthenium; S, sulfur; Sb, antimony; Sc, scan-

dium; Se, selenium; Si, silicon; Sm, samarium; Sn, tin; Sr, strontium; Ta, tantalum; Tb, terbium;

Tc, technetium; Te, tellurium; Th, thorium; Ti, titanium; Tl, thallium; Tm, thulium; U, uranium;

V, vanadium; W, tungsten; Xe, xenon; Y, yttrium; Yb, ytterbium; Zn, zinc; Zr, zirconium)

22719.2 Structures of Atoms

(symbol H) is about 1 g, carbon (C) 12, oxygen (O) 16, iron (Fe) 56, and so on.

Note that when you combine two hydrogen atoms and one oxygen, you will

come up with 2 × 1(H) + 16(O) = 18 g for 1 mol of water as indicated above. This

mass is called the “molar mass” of a substance. For example, the molar mass of

the rust Fe

would be approximately 2 × 56 + 3 × 16 = 160 g. There is a deeper

meaning to this fact; i.e., that the mass of 1 mol of an element comes close to an

integer number. In order to understand this, we need to look into the structure of

an atom.

19.2 Structures of Atoms

A lot of research on the structures of atoms has come up with a picture shown in

Fig. 19.3. An atom consists of a nucleus at the center and electrons surrounding it.

The nucleus carries a positive electric charge and is very minute, while electrons are

negatively charged. And the magnitudes of the positive charge on the nucleus and

of the negative charge on electrons are the same (in a neutral atom). Since an elec-

tron is regarded as a particle (one of the fundamental particles), we can assign an

electric charge to a single electron and use that charge as a unit of electric charge;

let’s say it is “e.” Suppose that a neutral atom has “n” electrons, then the nucleus

must have “n” of something that carries a positive electric charge of magnitude “e.”

This something is called “proton.” A proton, though carrying the same magnitude of

electric charge positive (instead of negative), turned out to be about 1,800 times as

heavy as an electron.

The simplest atom consists of a single proton in the nucleus and an electron sur-

rounding it. This is hydrogen, symbolized as “

.” It turned out that another atom

exists that has one proton and one electron, but, in addition, has one more particle

called “neutron” in the nucleus. “Neutron” is approximately the same in mass as

proton but carries no electric charge. This atom is about twice as heavy as the regu-

lar hydrogen atom and is called “heavy hydrogen”

(or often called “deuterium D”

(

)). In this notation, the superscript is the total number of protons and neutrons

in the nucleus and the subscript is the number of protons (or the positive electric

charge in units of “e”). The former is called “atomic mass number” and the latter

“atomic number,” or more generally electric charge number. The reason that this is

Fig. 19.3 The atomic

orbitals that the electrons

will occupy. These are

actually more like fuzzy

clouds of different densities

and shapes

228

19 Chemistry’s View of the Material World: Basic Principles

still regarded as a kind of hydrogen is that deuterium (D) behaves almost like hydrogen,

as the chemical behaviors of an atom are determined mostly by the number of

electrons and the positive charge of the nucleus. Scientists invented a word “iso-

tope” to designate two or more atoms that have the same number of protons but

different number of neutrons in the nucleus. The isotopes with an atomic number

(i.e., the same number of protons) are chemically similar and hence designated as

belonging to the same element. Thus, both “

” and “

” (or

) are isotopes and

belong to the element “hydrogen.” A third hydrogen isotope

[how many protons

and neutrons here?] is also known. This is often called “tritium, T.” Not all the

isotopes of an element exist in an equal quantity in nature (on the Earth). In hydro-

gen case, the predominant isotope is

, whose natural abundance on the Earth is

99.985%, while the heavy hydrogen is present only 0.015% in nature (on the Earth).

Tritium can be produced only artiﬁcially.

Now let us move on. The second simplest element is helium, He or rather

;

that is, its nucleus contains two protons and two neutrons. The third element is

(lithium, how many neutrons are here?). There is also an isotope

. The natural

abundance of the former is 92.58% and the latter is 7.42%. Element iron, Fe, con-

sists of isotopes of

(natural abundance = 5.82%),

(91.66%),

(2.19%), and

(0.33%) (how many protons and neutrons here?). Since a proton

and a neutron weigh about the same, the mass of an isotope, say

would weigh

56 g (per mole), if we assign 1 g for 1 mol of hydrogen (i.e.,

). It usually comes

close to that, i.e., an integer number, but it is not quite an integer number. This is

because some mass (weight) is lost when protons and neutrons gather together to

form a nucleus. This weight lost is turned into an energy, through the famous

Einstein’s equation E = mc

, where E is the energy, m is the mass, and c is the speed

of light.

Nevertheless, this fact that a nucleus is made of an integer number of particles

(protons and neutrons) is the basis for the choice of the Avogadro number. Avogadro

number has been chosen in such a way that the molar quantity of an element (an

isotope to be more precise) comes out as a number very close to an integer [because

the mass loss in the form of mc

is relatively small]. Historically, the Avogadro

number has been reworked several times; currently the Avogadro number (i.e.,

6.022 × 10

) is chosen so that 1 mol of the isotope

weighs exactly 12.000 g. The

mass of 1 mol of an atom is called “(molar) atomic mass” (given in units of g/mol).

It should be and is close to an integer number (i.e., the mass number, the superscript

ﬁgure of isotope), if it is deﬁned for a single isotope. However, the “atomic mass” is

usually deﬁned for an element. Hence the term “(molar) atomic mass” is a misno-

mer and should be called “(molar) elemental mass.” The atomic mass for an element

can differ signiﬁcantly from an integer number, because the atomic mass is the aver-

age of 1 mol of all the naturally occurring isotopes of an element. [Try to calculate

the (molar) atomic mass of iron based on the isotopic distribution data mentioned

above for a fun, assuming that each isotope has an integer number for its atomic

mass, though this assumption is not quite right. Compare your result with the

approximated value of molar atomic (elemental) mass given in the periodic chart,

Fig. 19.2. If you ﬁnd a discrepancy, think why].

22919.4 Nuclear Chemistry

19.3 Ordinary Chemistry and Nuclear Chemistry

An atom consists of a nucleus that is made of positively charged protons and neutral

neutrons, and electrons surrounding it, as we outlined above. Two entirely different

types of chemistry stem from this structure. One is concerned with the nucleus and

the other with how electrons behave. The former is “nuclear chemistry,” with

“radiochemistry” as its important sub-discipline. The latter is the ordinary “chemis-

try.” The basic reason for this division is that the nuclear forces binding protons and

neutrons in the nucleus are enormously stronger than the electrostatic force binding

the electrons to the nucleus. When one applies a force to a substance and induces a

change, a certain amount of energy may be expended or gained. Hence, an energy

change always accompanies a change in substance. “Energy” is often used as a

measure of a change in science, particularly in chemistry (see Sect. 19.8). In terms

of energy, then, a nuclear reaction (change in general) is greater by several orders of

magnitude (typically a million times) than a typical chemical reaction, as the nuclear

reaction involves changes in protons/neutrons in the nucleus while chemical reac-

tions involve changes in electrons. Therefore, ordinary chemical reactions would

not be able to cause a change in nucleus (i.e., nuclear reaction). As a result, it is

quite safe to deal with nuclear chemistry as separate from “ordinary” chemistry. As

a corollary, all isotopes that belong to an element, though they have different atomic

masses, can be assumed to behave (approximately) the same chemically. However,

isotopes behave very differently in terms of nuclear reactions. It is now obvious that

principles governing nuclear reactions are quite different from those operating in

the ordinary chemical reac-tions.

19.4 Nuclear Chemistry

19.4.1 Radioactivity: Spontaneous Nuclear Reactions

A nucleus that consists of neutrons and protons can be stable or unstable.

Understanding of the factors that affect the nuclear stability is beyond the scope of

this discourse, and it is not necessary to really understand for our purpose here.

However, we have the facts; i.e., which nuclei (isotopes) are stable and which are

unstable. The unstable isotopes would not remain as such and spontaneously change

into stable isotopes. In this process, the unstable nucleus emits particles and/or

energy in the form of radiation (gamma (g)-ray). Particles that are emitted include

electron (called b-particle), and helium nucleus (a-particle). Hence, these unstable

isotopes are called “radioactive,” emitting a, b, and/or g radiation. There are other

kinds of radiation, as well, but these three are the major ones.

There are exactly 264 stable isotopes known on the Earth. They lie on or slightly

above a line that corresponds to the composition of equal numbers of neutrons and

protons, as shown in Fig. 19.4. The area that is occupied by stable isotopes is coined

as “stability peninsula” (the darkened area; the solid black curve represents a sort of

central area where most of stable isotopes are distributed). In the ﬁgure, each of the

230

19 Chemistry’s View of the Material World: Basic Principles

colored areas consist actually separate dots, each of which represents an isotope

(stable or unstable), and it represents an approximate area where these isotopes

are distributed.

In understanding the radioactivity, we need to recognize some changes that can

occur with neutron and proton. That is, they can interchange in the following

manner:

1 10

0 11

n neutron p pro( )( ) (ton e elect )ron

®+

(19.1)

101

11 0

pen

+®

(19.2)

(

)

1 10

1 01

p n e positron

®+

(19.3)

[In these equations we omitted another particle “neutrino”].

Unstable isotopes that lie above the stability peninsula have too many neutrons.

Hence they become stable by reducing the relative number of neutrons. This can be accom-

plished by emitting an electron (i.e., converting a neutron to a proton through (19.1)).

Radioisotopes

Stable Isotopes

Atomic Number (Number of Protons)

Number of Neutrons

100

60 70 80 90

120

110

–emission

neutron–emission

Radioisotopes

b-emission

neutron-emission

mostly a-emission

Line of equal

number of

neutrons and

protons

Fig. 19.4 Stable isotopes and unstable (radioactive) isotopes

23119.4 Nuclear Chemistry

Therefore, most of these neutron-rich unstable isotopes are b-emitters. What would

be an alternative way to reduce the neutron/proton ratio?

On the other hand, the unstable isotopes that lie below the stability peninsula would

want to reduce the excess protons, by e.g., either (19.2) or (19.3). Process (19.2) is

called electron capture and would not emit a particle, but usually emit g-radiation.

Process (19.3) would emit a positron, which has the same character as an electron

except for the positive electric charge (instead of negative); this is called b

-radiation.

As you see in Fig. 19.4, there are no stable isotopes for elements whose atomic

numbers are 84 or higher. These heavy elements usually decay by emitting a-parti-

cles. An example is

235

, which is used for the nuclear power generation.

The speed of the radioactive decay (disintegration) is governed by only the nature

of the nucleus, and often is characterized by the time that it takes for one half of a

sample of a radioactive isotope to disintegrate. This “time” is called “half-life.”

Some isotopes have very short half-life: e.g., 2.4 s in the case of

. They disinte-

grate very quickly. Others have very long half-life: e.g. 4.5 × 10

years in the case of

238

. This means that it takes 4.5 billion years (approximately the life span of the

Earth) for 1 g of radioactive

238

to become half a gram of radioactive

238

; in this

process, the end product is lead

206

, through a multistep process.

19.4.2 Induced Nuclear Reactions: Modern Alchemy

Nuclei, stable or unstable, behave on their own. It is difﬁcult to make them change

their courses of destiny. This is because the energies involved in the change are

quite large. And this is the reason that alchemists had never succeeded in converting

lead into gold. However, if large enough energy is provided, one can cause changes

in a nucleus. Such nuclear reactions are taking place in nature (in the universe) and

also can now be induced by artiﬁcial means.

In the upper portion of the atmosphere on the Earth, nitrogen isotope

is con-

stantly bombarded by (energetic) neutrons of the cosmic ray and is undergoing the

following nuclear reaction:

14 1 14 1

7 06 1

NnCp+® +

[Note that the sum of the superscripts on the left-hand side (14 + 1 = 15) is equal

to that on the right-hand side and also that the sum of the subscripts on the left-hand

side (7 + 0 = 7) is equal to that on the right-hand side (6 + 1 = 7). This conservation of

mass number and electric charge always holds true for nuclear reactions. Check this

with (19.1)–(19.3) above as well]. By the way, the resultant nucleus

is radioac-

tive, though

is not, and the radioactivity of

is used to determine the age of

the archaeological artifacts.

British physicist Rutherford succeeded in 1919 for the ﬁrst time in human history

in changing an element into another. He bombarded

with a-particles and

converted it into oxygen. That is:

14 4 17 1

7 2 81

N He O p Check the balance of super/subscr( .[]t) ip sa+ ®+