Ochiai E. Chemicals for Life and Living

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242

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

The electrons around the nucleus actually ﬁll up the space between the atoms.

So, another way of representing a molecular structure is to ﬁll the space that repre-

sents roughly the shape of an electron cloud in the molecule. This model is called

“space-ﬁlled.” They are exempliﬁed by those in Fig. 19.8c. We will often use the

simplest structural expression shown in Fig. 19.8a, because it shows clearly the con-

nections (bonds) between atoms. The “ball and stick” representation is good at show-

ing the three-dimensional relative arrangement of atoms, but often it can obscure

some of the atoms, unless we rotate the molecule. The “space-ﬁlled” model is good

at showing the overall three-dimensional structure.

How about C

, ethylene (“ethane” in technical terms)? It has two less hydro-

gen atoms than ethane. If you remove two hydrogen atoms from ethane, you would

end up with C

but you would also see that one electron on each carbon will be

left. The carbon–carbon bond present in the original ethane remains as such, but the

odd electron on each carbon that is created in the process is distributed in the direc-

tion perpendicular to the C–C bond. The direction of the C–C bond is deﬁned as

z-axis. The odd electron, technically speaking, is in a p

atomic orbital. The elec-

trons in p

orbital, one in each carbon atom, also contribute to binding the two car-

bon atoms. The distribution of electrons in this type of bonding, which is designated

as p-bond is shown in Fig. 19.9. Hence the connection C-to-C in C

molecule is

made of two different types of bond, s and p and is contributed by two pairs of

electron. The bond between the carbons in this molecule is said to be a “double”

bond and is indicated by a double line (see Fig. 19.8). You may guess that the bind-

ing through the p-bond is weaker than that of s-type bond, as the electrons in p-bond

are less effectively overlapped.

Let us consider another simple-looking molecule, N

, dinitrogen, the major com-

ponent of the air. Each nitrogen atom has ﬁve valence electrons available. One

electron on each nitrogen atom binds in the manner of the s-bond. The second elec-

tron in p

(refer above) forms a p-bond. The third electron is in another p-orbital, say p

which is perpendicular to the s-bond as well as p

. This electron in an N atom

forms another p-bond together with that on another N atom. So now, there are one

s-bond, and two p-bonds that are perpendicular to one another (Fig. 19.10). Two

more electrons still remain in each nitrogen atom. These are not involved in the

bonding and instead remain in an orbital in each N-atom. The two electrons have to

have opposite spins as they occupy the same space, and they are called “lone pair”

(not involved in bonding). The entire situation of electrons in an N

molecule is

shown in Fig. 19.10. So we may say that the bond in N

is a “triple” bond.

How about the next molecule, O

? We now have two more electrons. These will

remain as independent electrons with the same spin, but have an effect to cancel the

Fig. 19.9 Sigma (s)-bond

and pi (p)-bonds

24319.7 Molecules/Compounds: Chemical Bonding and Structures

boding effect of a p-orbital. Then the bonding in O

is equivalent to a “double” bond.

The resulting electron conﬁguration of O

is shown in Fig. 19.10b. Note that the O

molecule has two unpaired electrons; that is, those two electrons have the same spin

(either both are +1/2 or both are −1/2). This has an interesting consequence. As a spin

behaves like a magnet, O

behaves as if it is a magnet. This is called “paramagnet-

ism.” All the other molecules mentioned above in this section are, instead, diamag-

netic. The electrons in these compounds are all paired up as in s- and p-bonds or lone

pairs. One of a pair of electrons has a spin of +1/2 and the other −1/2, so that their

magnetic characters cancel one another and, as a result, these compounds behave as

non-magnets. The fact that O

is paramagnetic can readily be demonstrated by pour-

ing liquid O

through a magnet. The liquid stream will be attracted toward and sticks

to the poles of the magnet (Fig. 19.11). Nothing like this happens with liquid nitro-

gen N

; it will stream through the poles of magnet without sticking.

The overall shape of these molecules, i.e., N

and O

, would be similar to that of

shown in Fig. 19.7; that is, a dumb bell, two balls bound together. Indeed such a

shape has been observed for O

recently by a technique called scanning tunneling

microscope. It is shown in Fig. 21.7.

Fig. 19.10 Bondings

in (a) dinitrogen and

(b) dioxygen; each of p

, p

and sp-orbitals (one on each

nitrogen atom) contains a pair

of (coupled) electrons; in the

case of dioxygen molecule,

each of the two additional

electrons occupies the p

and p

* orbitals, respectively.

In both of these, the sigma

bond between the two atoms

are omitted

Fig. 19.11

Paramagnetism

of dioxygen molecule (from

P. Atkins and L. Jones,

“Chemical Principles”, 3rd

edn (2005, W.H. Freeman))

244

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

One more interesting mode of bonding needs to be mentioned before we leave

this topic. It is illustrated by benzene, C

. The six carbon atoms form a skele-

ton of hexagon (see Fig. 11.3). This is formed by s-bonds between two adjacent

carbons. Each carbon has used three electrons for these bondings, and has still

one more electron left. These electrons are in orbitals (p) that are sticking up and

below of the hexagonal plane. These electrons can be combined to form three

p-bonds between adjacent carbons (see Fig. 11.3). But as Fig. 11.3 shows, there

are two ways of doing this, and neither of them should be preferred over the

other. In fact the two are equivalent, and the reality may be said to be a sort of

mixture of both forms. This situation is called (chemical) “resonance.”

Unfortunately, “resonance” in everyday use or physics has a different connota-

tion. It may be better described as a six-electron cloud being distributed equally

along the hexagon perimeter above and below the plane, like a sandwich of

doughnuts (see Fig. 19.12). The atomic structure of benzene has recently been

imaged also by a technique called scanning tunneling microscope. The image

that is shown in Fig. 21.8 indeed shows that this description of the benzene’s

structure is valid.

19.7.4 Polarity of Covalent Bond

The electrons in H

, N

, or O

molecule will be distributed equally between the two

atoms, because the electron-attracting strength of the two atoms is the same. Thus

the bonding in these molecules can be said to be “purely covalent,” and such a bond

Fig. 19.12 Structure of

benzene: (a) emphasizes the

p-clouds (above and below

The hexagonal plane);

(b) represents each atom

by a space-ﬁlling sphere,

hydrogen atoms bound with

each of 6 carbon atoms are

omitted in (a), and the white

sphere represents hydrogen

atom in (b)

24519.7 Molecules/Compounds: Chemical Bonding and Structures

is called “non-polar.” That is because each atom in say N

would share ten valence

electrons equally and hence would have ﬁve electrons about it and would appear to

be electrically neutral, no net electric charge on it.

The electron-attracting strength of an atom (or rather element) is called “electro-

negativity.” Let’s consider methane, CH

. There are four C–H bonds (see Fig. 19.8).

The electronegativities of elements carbon and hydrogen happen to be about the

same (though C is slightly more electronegative than H). Therefore, the electrons

are almost equally distributed between C and H, and hence this bond is considered

to be non-polar.

What happens if we replace one of the hydrogen atoms of methane with chlo-

rine, for example, i.e., CH

Cl (monochloromethane)? As chlorine is more electro-

negative (attracts electrons more strongly) than carbon, the two electrons located

in between carbon and chlorine are not equally distanced from the two atoms.

They are located closer to the chlorine atom. So it looks that the chlorine is nega-

tively charged and the carbon is positively charged. It is expressed as: C

−Cl

d−

The value of d here is called partial charge and less than 1. And the bond C–Cl is

said to be “polar.” If the value of d is 1, the two electrons are now completely

located in the chlorine atoms and are not shared by carbon and chlorine atoms.

The bonding then would be essentially an ionic bond. The polarity of a bond is

perhaps the most important character, which will dictate the behavior and reactiv-

ity of a compound (molecule). Polar bonds will form between atoms of different

electronegativity, but important ones among carbon compounds (i.e., organic

compounds) are C–X, where X is F (ﬂuorine), Cl (chlorine), Br (bromine), and I

(iodine). By the way these four elements are collectively called “halogen.” The

other important ones are C–O, as in CH

–OH (methanol), and C–N and C–S. And,

of course, O–H bond itself in methanol is also polar. [Which carries a small posi-

tive charge, O or H here?] On the contrary, when a carbon atom binds with a

metallic element such as Mg, the carbon atom will carry a negative charge and the

metal (Mg in this case) will be positive, because metallic elements are much less

electronegative than carbon.

Let’s now take a look at carbon dioxide: O=C=O. The carbon–oxygen bond here

is polar (negative on oxygen and positive on carbon). The polar character, polarity,

is deﬁned in terms of dipole moment which is directed from the positive toward the

negative and deﬁned as “the partial charge x the distance between the charges.”

In this case, one bond (say, the right-hand C=O) has a dipole moment directed to the

right and the other bond (the left-hand O=C) has the dipole moment of the same

magnitude but directed to the left. Therefore, the dipole moments cancel one another,

and the molecule CO

does not have a dipole moment, and hence it is non-polar as

a molecule. The polarity of a molecule is thus determined by the polarities of indi-

vidual bonds and the entire structure of the molecule. The former, i.e., bond polarity

is important in determining the chemical reactivity of a molecule, and the polarity

of the entire molecule is important in determining its physical behavior. Polar mol-

ecules, because of having separate electric charges in themselves, attract each other

more strongly than non-polar comparable molecules. And this will be reﬂected in

their physical state, as will be discussed in the next section.

246

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

19.8 Energy: How and Why Changes Occur

19.8.1 Energies

Chemicals, whether atoms, molecular compounds, or ionic compounds, are made of

particles: nuclei and electrons, atoms, ions, and molecules. The particles affect each

other; i.e., they interact. Energy is associated with any interaction. For example,

a positive ion and a negative ion attract one another, and the magnitude of interaction

can be expressed in terms of energy that will be required to separate these two ionic

species. We have already talked about covalent bonding between atoms; this is

another example of interaction. The energy that will be required to separate two

atoms bound by a covalent bond is called “bond energy.” Different bonds would

have different bond energies.

A bond or rather atoms bound by a bond are not static. For example, H-atoms in

an H–H molecule do not sit still. They are moving all the time. Rather, they are

vibrating about a certain distance. This distance is a time-averaged value and called

“bond length (or bond distance).” Vibration requires a certain amount of energy.

If you heat up (i.e., add energy to) a molecule, the vibration will become more vigor-

ous. A molecule is also tumbling (rotation), and this motion is also associated with

energy. The motions mentioned so far are those within a molecule; we can say, these

are intra-molecular motions.

We do not, however, deal with a single molecule ordinarily. We will be dealing with

an ensemble of a very large number of molecules (and ions). So we have to consider the

interactions between molecules as well. Again the force is “electrostatic.” Then if the

interacting entities are ions, the interaction energy will be large and depends on their

electric charges and the distance between them. Polar molecules will interact through

the partial charges as well in them. The partial charges in a molecule are represented by

a value called “dipole moment,” which depends on both the magnitude of the partial

charges and the distance between the charges in a molecule. The interaction between

polar molecules depends on the dipole moment and the distance between molecules.

The interaction energy will be small between neutral non-polar entities. There is

a fundamental electrostatic interaction, however, even between neutral non-polar

entities. This is called “London dispersion” force, and it depends on the size (hence

the molar mass) of the molecule. The elements in the last column of the periodic

chart are called inert gas elements; they are neutral and non-polar (spherical).

Therefore, the smallest, helium He, has the lowest boiling point, because the interac-

tion energy between helium atoms is minimal. By the way, these inert gas elements

are literally inert; that is, they can neither form a molecule with itself (like He–He)

nor with other elements (though Xe-containing compounds have been made). The

relationship between the inter-particle interaction and the boiling point (and others)

will be discussed in the next section. As you go down the column, Ne–Ar–Kr–Xe,

their boiling points go up; the boiling point of Xe is −107°C as compared with that

of He, −269°C. This trend reﬂects the increase of London dispersion force as you

go down the column, i.e., the elements become larger and heavier, and hence the

inter-particle energy (London force) becomes larger, as you go down.

24719.8 Energy: How and Why Changes Occur

Ordinarily, the interaction energy is largest between ions, then between polar

molecules, and weakest between non-polar molecules. However, the non-polar

interaction (London dispersion) can become fairly large when it comes to larger mol-

ecules, even if they are non-polar.

There is a special type of interaction between an electronegative atom (such as

F and O) in a molecule and a hydrogen atom bound to an electronegative atom in

another molecule. Such an interaction is called “hydrogen bond.” For example,

a hydrogen atom of a water molecule interacts with an oxygen atom of another water

molecule through a hydrogen bond. The nature of the hydrogen bond has not yet

been understood fully but it is certainly a mixture of ionic interaction and covalent

interaction. Hydrogen bonds are crucial in the biological systems, i.e., our lives, as

we discuss in several other chapters.

So an ensemble of molecules or compounds (let’s call such an ensemble a “system”)

would have a certain amount of energy under a certain condition. This energy is

called “internal energy,” which depends among other things on the pressure under

which the system exists. It will be more convenient to use energy value under a

constant pressure, because we normally operate things under the ambient constant

pressure. Such an energy is technically called “enthalpy.” We mean “enthalpy”

when we say “energy” in this book, unless otherwise stated. If a system changes

(in a widest sense), the accompanying energy would also change, because the inter-

actions among all particles inside molecules and/or between molecules in the sys-

tem would change.

19.8.2 Phase Changes and Energies

Suppose we have water at 25°C. It is a liquid, in which water molecules are

rapidly moving around but are always in close contact with each other. If we want

to separate a molecule from another, we need to put some energy, perhaps in the

form of heat. Heat would encourage more vigorous movements of molecules and

may overcome the attractive interactions between them. If this happens, water

molecules get out of the conﬁne of liquid, and move far apart from each other

(except for occasional collisions). Now they are in the form of “gas” (vapor). This

is a change of state, from liquid state to gas state. Since you added heat (a form of

energy) in the process, you can imagine that the energy content (enthalpy) of a gas

is larger than that of the corresponding liquid. This is an example of the so-called

phase change. It is accompanied by a change in energy content. As a liquid

changes to a gas, the system would gain energy; we say that the energy change (of

the system) in this case is positive. This means that the system’s energy will

increase (i.e., positive) as it turns from liquid to gas. The opposite process, i.e.,

vapor turning into liquid (condensation), would then release energy which goes to

the surrounding; hence the system loses energy and hence the energy change in

the system accompanying the condensation process is negative. The change from

liquid water to solid ice is another example of phase change, and accompanied by

a negative energy change.

248

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

How high temperature you have to bring a liquid to in order to vaporize would

depend on how strongly molecules are interacting in the liquid. Technically speak-

ing, the temperature at which the vapor pressure (gas pressure above liquid) becomes

1 atmospheric pressure is deﬁned as boiling point. So the boiling point depends on

how strongly molecules interact. In the case of water, the boiling point is 100°C,

reﬂecting a relatively strong interaction between water molecules. Water (H

O) is

polar, resulting in a strong dipole–dipole intermolecular interaction, and in addition

there is a strong hydrogen-bond interaction. On the contrary, comparable molecules

such as CH

(methane) and NH

(ammonia) have much lower boiling points,

−164 and −33°C, respectively. Ammonia is polar but less so than water and the

hydrogen-bond interaction is not as strong. Methane is a non-polar molecule and

the only possible interaction between methane molecules is the weak London

dispersion force.

Melting point, the temperature at which a compound melts, is also dependent on

the strength of intermolecular (inter-particle in general) interactions. Naphthalene,

the smelly stuff used as an insect repellant (mothballs) in the chest of drawers, is a

non-polar neutral molecule and hence the intermolecular interactions are relatively

weak. Therefore, its melting point is rather low, 80.5°C, and thus it relatively easily

turns into gas (hence the smell). In contrast, table salt, sodium chloride, melts at a

high temperature, 801°C, because the components, sodium ions (Na

) and chloride

ions (Cl

−

), attract each other very strongly. In general, ionic compounds have high

melting and boiling points, while molecular compounds melt and boil at relatively

low temperatures. The strength of interaction between molecules is also a determin-

ing factor for the other physical properties such as surface tension and viscosity. The

main principles and concerns of chemistry up to this point (i.e., except for the issue

of chemical reaction which is discussed below) can be summarized as in Fig. 19.13.

water

Cocept of Mole

made of a large

number of

Water molecules

water

hydrogen

atom

nucleus

oxygen

atom

electron

Oxygen

nucleus

proton

neutron

molecules

Avogadro No.

= 6.02 x 10

Intermolecular

Interactions between

Nuclear Forces

Stability of a nucleus

depends on the

numbers of protons

and neutrons.

Atoms

determine Bonding

and Reactivity of a

Molecule. This is

governed by Electrons

and their Distributions.

Interactions

determine

Physical

Properties

of a

Compound.

REGULAR CHEMISTRY

Avogadro No has been determined so that a mole of a carbon isotope

C–12 which consists of 6 protons and 6 neutrons (i.e., 12 nucleons)

comes exactly to 12.000 gram/mole.

NUCLEAR CHEMISTRY

Fig. 19.13 Chemistry in a Nutshell

24919.8 Energy: How and Why Changes Occur

19.8.3 Writing Chemical Reaction Equations

A chemical reaction involves any number of chemicals, say, chemical species A, B,

C, D. Let us suppose that A and B react chemically and, as a result, C and D would

form. In other words, a chemical reaction involves change in chemical species.

There is one more thing that needs to be added in order to be able to write this

chemical reaction equation. That is the so-called stoichiometric relationship. It is

about the relationship among the quantities of these chemical species when they

react. For example, two molecules of A react with three molecules of B and form

two molecules of C and one molecule of D. This chemical reaction can then be

represented by the following type of equation, though the equal sign “=” is now

usually replaced by an arrow sign:

2A 3B 2C 1D ( 1 is often omitted) or 2A 3B 2C D+=+ +®+“”

In this particular case, the reaction is supposed to proceed from the left-hand side

to the right-hand side, and the species on the left-hand side (A and B in this case)

are called “reactants” and those on the right-hand side are “products.” The stoichio-

metric relationship implies only the overall (relative) quantitative relationship

among the chemical species involved and would not imply how the reaction takes

place. The latter issue is called “reaction mechanism.” That is, the chemical reaction

equation such as shown above does not necessarily imply the reaction mechanism.

The reaction mechanism is a very difﬁcult issue, and its delineation involves a lot of

detailed studies on the reaction, and currently is one of the most interestingly stud-

ied in the chemical/biological science.

The stoichiometric balance is based on the principle that in a chemical reaction

no change (both quantity and identity) should take place regarding all the individual

atoms involved. In other words, the total number of atoms on the left-hand side

should be equal to that on the right-hand side. And this applies to all the elements

involved.

Let us try a few examples.

1. Sodium metal (represented by the symbol Na) burns (reacts with) in chlorine

gas (Cl

) and forms sodium chloride (NaCl) and nothing else:

Na Cl NaCl+®

Chemical species and their arrangements are correct, but the stoichiometric

relationship is not correctly represented. In other words, the total number of

atoms in chemical compounds should be equal on both sides. You have the same

number of Na on both sides, but there are two Cls on the left-hand side and only

one Cl on the right. To adjust the number you can do the following:

(

)

Na 1/2 Cl NaCl+®

250

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

Now you have the same number (“1”) of Cl atoms on both sides. Therefore, this

is a correct chemical reaction equation. However, the fractional number (in front

of a chemical) should be avoided. So you can multiply the whole quotation by2,

and will get:

2Na Cl 2NaCl+®

This process of adjusting the stoichiometric relationship is called “balancing the

equation,” and the resultant equation is said to be “balanced.”

2. Try this. Magnesium metal (Mg) would be burnt in air; that is, it reacts with

oxygen (O

) in the air and forms only magnesium oxide (MgO). Write a bal-

anced equation.

3. When you dissolve copper sulfate (CuSO

) in water, it forms a light blue solu-

tion, and the solution contains two ions, Cu

and SO

2−

(sulfate). This reaction is

simply expressed as

(

)

(

)

CuSO Cu aq SO aq

®+

“(aq)” implies the ionic species are interacting with water (aquo) molecules, and

“Cu

(aq)” is responsible for the light blue color. Now if you add an aqueous

ammonia solution (i.e., consisting of NH

and H

O) to this light blue solution, a

light blue precipitate, copper hydroxide (Cu(OH)

), and ammonium ion (NH

)

form. This reaction can be expressed as

(

)

(

)

32 4

Cu aq NH H O Cu OH NH

++® +

That is, only Cu

(aq) is involved in this reaction, and sulfate (SO

2−

) is simply

present in the solution. Is this equation balanced? If not, try to balance it. (You

can ignore (aq) in balancing an equation).

4. Here is one of the most complicated reactions. MnO

−

(permanganate) reacts

with C

2−

(oxalate) in the presence of acid (H

) and the products are Mn

, CO

(carbon dioxide), and H

O (water). That is:

4 24 2 2

MnO C O H Mn CO H O.

- -+ +

+ +® + +a b cd e f

The stoichiometric coefﬁcients a,…, f are to be determined so that the equation

is balanced. In this case the chemical species are electrically charged (i.e., ions),

and you need to balance the electric charge as well. If you are comfortable with

algebra, you can do something like the following. Because the numbers of Mn

(manganese) on the left- and the right-hand side should be the same, “a” has to

be equal to “d ”; a = d. Likewise, 2b = e with regard to C; 4a + 4 b = 2 e + f for O; and

c = 2f for H

The electric charge on the left-hand side is –a –2b + c and that on the right-

hand side is 2d. These equations constitute a set of simultaneous equations. All

you need to do is solve them. You will realize that you have only ﬁve equations

(relationships) for six unknown values, so that you cannot solve them. It is OK,

25119.8 Energy: How and Why Changes Occur

because all we need is relative values. One thing you can do in this situation is

assume, for example, a = 1.

Then d = 1. Since e = 2b, the balance equation for O will become 4 + 4b = 4b + f

so that f = 4. This means c = 8. From the electrical balance equation you get

c–2b = 3. Therefore, b = 5/2. Now you have obtained: a = 1, b = 5/2, c = 8, d = 1,

e = 5, f = 4. Because b is a fractional number, you multiply everything by 2, and

ﬁnally obtain

4 24 2 2

2MnO 5C O 8H 2Mn 10 CO 8H O

- -+ +

+ +® + +

Balanced chemical reaction equations are necessary when quantities of reac-

tants and products are to be calculated and/or the energy relationship as below is

to be evaluated. Otherwise chemical reactions, even if not balanced, are usually

sufﬁcient to provide the necessary information about the chemical reactions.

19.8.4 Reactions and Energies: Free Energies

Now let’s take a simple reaction: i.e.,

( ) ( )

2H H H O O O 2H O H- += ® --

[This is usually written as

22 2

2H O 2H O+®

]. This reaction involves splitting two

H–H bonds and one O=O bond, and then recombining them to form two molecules

of H–O–H which has two O–H bonds. Bond energies of H–H, O=O, and H–O are

432 kJ/mol, 494, and 459, respectively. Therefore, the energy change accompanying

the reaction will be: 2 × 432 + 494 – 4 × 459 = − 478 kJ. This means that 2 mol of

water is more stable by 478 kJ than the starting material, 2 mol of H

and 1 mol

of O

. [By the way, the end product is assumed to be gaseous water (water vapor) in

this calculation]. Another reaction

C CO 2CO+®

is accompanied by an energy

change of +172 kJ.

Why are we concerned about the energy change accompanying a phase change,

a chemical reaction, or a change of other kinds? Because the energy change is the

driving force for a chemical change. Ordinarily, a change would occur in the direc-

tion in which its energy content decreases or becomes lower, i.e., the system becomes

lower in energy or more stable. In other words, a system will change if the change

brings the system to a lower, more stable state, as water ﬂows from a higher place to

a lower place.

However, it has been demonstrated that a system can turn into a higher-energy

state if a certain condition is met. There is a deeper criterion for the direction of

chemical change. The most basic principle in determining the direction of change is

called “2nd law of thermodynamics.” In order to understand it, we need to introduce

one more concept, entropy. “Entropy” is a measure of randomness and related to the

change in enthalpy divided by the prevailing temperature (as expressed in absolute

temperature scale). The 2nd law of thermodynamics can be expressed in many ways.

One version is: “the entropy of the universe cannot decrease.” That is, the random-

ness of the universe as a whole should either remain the same or increase, and never

decrease. Then, the entropy of the universe eventually will reach a maximum; at that

point no further change can occur. This is called the entropy death of the universe.