VI. PHYSICAL CHEMISTRY_Encyclopaedia Britannica, 11th Edition, "Châtelet" to "Chicago" Volume 6, Slice 1

We have seen how chemistry may be regarded as having for its province

the investigation of the composition of matter, and the changes in

composition which matter or energy may effect on matter, while physics

is concerned with the general properties of matter. A physicist,

however, does more than merely quantitatively determine specific

properties of matter; he endeavours to establish mathematical laws which

co-ordinate his observations, and in many cases the equations expressing

such laws contain functions or terms which pertain solely to the

chemical composition of matter. One example will suffice here. The

limiting law expressing the behaviour of gases under varying temperature

and pressure assumes the form pv = RT; so stated, this law is

independent of chemical composition and may be regarded as a true

physical law, just as much as the law of universal gravitation is a true

law of physics. But this relation is not rigorously true; in fact, it

does not accurately express the behaviour of any gas. A more accurate

expression (see CONDENSATION OF GASES and MOLECULE) is (p + a/v^2)(v - b)

= RT, in which a and b are quantities which depend on the composition of

the gas, and vary from one gas to another.

It may be surmised that the quantitative measures of most physical

properties will be found to be connected with the chemical nature of

substances. In the investigation of these relations the physicist and

chemist meet on common ground; this union has been attended by fruitful

and far-reaching results, and the correlation of physical properties and

chemical composition is one of the most important ramifications of

physical chemistry. This branch receives treatment below. Of

considerable importance, also, are the properties of solids, liquids and

gases in solution. This subject has occupied a dominant position in

physico-chemical research since the investigations of van't Hoff and

Arrhenius. This subject is treated in the article SOLUTION; for the

properties of liquid mixtures reference should also be made to the

article DISTILLATION.

Another branch of physical chemistry has for its purpose the

quantitative study of chemical action, a subject which has brought out

in clear detail the analogies of chemical and physical equilibrium (see

CHEMICAL ACTION). Another branch, related to energetics (q.v.), is

concerned with the transformation of chemical energy into other forms of

energy--heat, light, electricity. Combustion is a familiar example of

the transformation of chemical energy into heat and light; the

quantitative measures of heat evolution or absorption (heat of

combustion or combination), and the deductions therefrom, are treated in

the article THERMOCHEMISTRY. Photography (q.v.) is based on chemical

action induced by luminous rays; apart from this practical application

there are many other cases in which actinic rays occasion chemical

actions; these are treated in the article PHOTOCHEMISTRY.

Transformations of electrical into chemical energy are witnessed in the

processes of electrolysis (q.v.; see also ELECTROCHEMISTRY and

ELECTROMETALLURGY). The converse is presented in the common electric

cell.

_Physical Properties and Composition._

For the complete determination of the chemical structure of any

compound, three sets of data are necessary: (1) the empirical chemical

composition of the molecule; (2) the constitution, i.e. the manner in

which the atoms are linked together; and (3) the configuration of the

molecule, i.e. the arrangement of the atoms in space. Identity in

composition, but difference in constitution, is generally known as

"isomerism" (q.v.), and compounds satisfying this relation differ in

many of their physical properties. If, however, two compounds only

differ with regard to the spatial arrangement of the atoms, the physical

properties may be (1) for the most part identical, differences, however,

being apparent with regard to the action of the molecules on polarized

light, as is the case when the configuration is due to the presence of

an asymmetric atom (optical isomerism); or (2) both chemical and

physical properties may be different when the configuration is

determined by the disposition of the atoms or groups attached to a pair

of doubly-linked atoms, or to two members of a ring system (geometrical

isomerism or allo-isomerism). Three sets of physical properties may

therefore be looked for: (1) depending on composition, (2) depending on

constitution, and (3) depending on configuration. The first set provides

evidence as to the molecular weight of a substance: these are termed

"colligative properties." The second and third sets elucidate the actual

structure of the molecule: these are known as "constitutional

properties."

In any attempts to gain an insight into the relations between the

physical properties and chemical composition of substances, the fact

must never be ignored that a comparison can only be made when the

particular property under consideration is determined under strictly

comparable conditions, in other words, when the molecular states of the

substances experimented upon are identical. This is readily illustrated

by considering the properties of gases--the simplest state of

aggregation. According to the law of Avogadro, equal volumes of

different gases under the same conditions of temperature and pressure

contain equal numbers of molecules; therefore, since the density depends

upon the number of molecules present in unit volume, it follows that for

a comparison of the densities of gases, the determinations must be made

under coincident conditions, or the observations reduced or re-computed

for coincident conditions. When this is done, such densities are

measures of the molecular weights of the substances in question.

_Volume Relations._[17]--When dealing with colligative properties of

liquids it is equally necessary to ensure comparability of conditions.

In the article CONDENSATION OF GASES (see also MOLECULE) it is shown

that the characteristic equation of gases and liquids is conveniently

expressed in the form (p + a/v^2)(v - b) = RT. This equation, which is

mathematically deducible from the kinetic theory of gases, expresses the

behaviour of gases, the phenomena of the critical state, and the

behaviour of liquids; solids are not accounted for. If we denote the

critical volume, pressure and temperature by V_k, P_k and T_k, then it

may be shown, either by considering the characteristic equation as a

perfect cube in v or by using the relations that dp/dv = 0, d^2p/dv^2 = 0

at the critical point, that V_k = 3b, P_k = a/27b^2, T^k = 8a/27b.

Eliminating a and b between these relations, we derive P_kV_k/T_k =

(3/8)R, a relation which should hold between the critical constants of

any substance. Experiment, however, showed that while the quotient on

the left hand of this equation was fairly constant for a great number of

substances, yet its value was not (3/8)R but (1/3.7)R; this means that

the critical density is, as a general rule, 3.7 times the theoretical

density. Deviation from this rule indicates molecular dissociation or

association. By actual observations it has been shown that ether,

alcohol, many esters of the normal alcohols and fatty acids, benzene,

and its halogen substitution products, have critical constants agreeing

with this originally empirical law, due to Sydney Young and Thomas;

acetic acid behaves abnormally, pointing to associated molecules at the

critical point.

Volume at critical point and at absolute zero.

The critical volume provides data which may be tested for additive

relations. Theoretically the critical volume is three times the volume

at absolute zero, i.e. the actual volume of the molecules; this is

obvious by considering the result of making T zero in the

characteristic equation. Experimentally (by extrapolation from the

"law of the rectilinear diameter") the critical volume is four times

the volume at absolute zero (see CONDENSATION or GASES). The most

direct manner in which to test any property for additive relations is

to determine the property for a number of elements, and then

investigate whether these values hold for the elements in combination.

Want of data for the elements, however, restricts this method to

narrow limits, and hence an indirect method is necessary. It is found

that isomers have nearly the same critical volume, and that equal

differences in molecular content occasion equal differences in

critical volume. For example, the difference due to an increment of

CH2 is about 56.6, as is shown in the following table:--

+--------------------+--------------+--------------+--------------+

+--------------------+--------------+--------------+--------------+

| Ethyl formate | H.C02C2H5 | 228 \ | 56.5 |

| Propyl formate | H.CO2C3H7 | 284 \ | 55.8 |

| Ethyl acetate | CH3.C02C2H5 | 285 } 283.3 | |

| Propyl acetate | CH3.CO2C3H7 | 343 \ | 57.4 |

| Ethyl propionate | C2H5.CO2C2H5 | 343 } 340.7 | |

| Methyl n-butyrate | }C3H7.CO2CH3 | 339 } | |

| Methyl isobutyrate | } | 337 / | |

+--------------------+--------------+--------------+--------------+

Since the critical volume of normal pentane C5H12 is 307.2, we have H2

= C5H12 - 5CH2 = 307.2 - 5 X 56.6 = 24.2, and C = CH2 - H2 = 32.4. The

critical volume of oxygen can be deduced from the data of the above

table, and is found to be 29, whereas the experimental value is 25.

Volume at boiling-point.

The researches of H. Kopp, begun in 1842, on the molecular volumes,

i.e. the volume occupied by one gramme molecular weight of a

substance, of liquids measured at their boiling-point under

atmospheric pressure, brought to light a series of additive relations

which, in the case of carbon compounds, render it possible to predict,

in some measure, the composition of the substance. In practice it is

generally more convenient to determine the density, the molecular

volume being then obtained by dividing the molecular weight of the

substance by the density. By the indirect method Kopp derived the

following atomic volumes:

C. O. H. Cl. Br. I. S.

11 12.2 5.5 22.8 27.8 37.5 22.6.

These values hold fairly well when compared with the experimental

values determined from other compounds, and also with the molecular

volumes of the elements themselves. Thus the actually observed

densities of liquid chlorine and bromine at the boiling-points are

1.56 and 2.96, leading to atomic volumes 22.7 and 26.9, which closely

correspond to Kopp's values deduced from organic compounds.

These values, however, require modification in certain cases, for

discrepancies occur which can be reconciled in some cases by assuming

that the atomic value of a polyvalent element varies according to the

distribution of its valencies. Thus a double bond of oxygen, as in the

carbonyl group CO, requires a larger volume than a single bond, as in

the hydroxyl group -OH, being about 12.2 in the first case and 7.8 in

the second. Similarly, an increase of volume is associated with doubly

and trebly linked carbon atoms.

Recent researches have shown that the law originally proposed by

Kopp--"That the specific volume of a liquid compound (molecular

volume) at its boiling-point is equal to the sum of the specific

volumes of its constituents (atomic volumes), and that every element

has a definite atomic value in its compounds"--is by no means exact,

for isomers have different specific volumes, and the volume for an

increment of CH2 in different homologous series is by no means

constant; for example, the difference among the esters of the fatty

acids is about 57, whereas for the aliphatic aldehydes it is 49. We

may therefore conclude that the molecular volume depends more upon the

internal structure of the molecule than its empirical content. W.

Ostwald (_Lehr. der allg. Chem._), after an exhaustive review of the

material at hand, concluded that simple additive relations did exist

but with considerable deviations, which he ascribed to differences in

structure. In this connexion we may notice W. Stadel's determinations:

CH3CCl3 108 CHClBr.CH3 96.5

CH2Cl.CHCl2 102.8 CH2Br.CH2Cl 88

These differences do not disappear at the critical point, and hence

the critical volumes are not strictly additive.

Theoretical considerations as to how far Kopp was justified in

choosing the boiling-points under atmospheric pressure as being

comparable states for different substances now claim our attention.

Van der Waal's equation (p+a/v^2)(v-b) = RT contains two constants a

and b determined by each particular substance. If we express the

pressure, volume and temperature as fractions of the critical

constants, then, calling these fractions the "reduced" pressure,

volume and temperature, and denoting them by [pi], [phi] and [theta]

respectively, the characteristic equation becomes

([pi]+3/[phi]^2)(3[phi]-1) = 8[theta]; which has the same form for all

substances. Obviously, therefore, liquids are comparable when the

pressures, volumes and temperatures are equal fractions of the

critical constants. In view of the extremely slight compressibility of

liquids, atmospheric pressure may be regarded as a coincident

condition; also C.M. Guldberg pointed out that for the most diverse

substances the absolute boiling-point is about two-thirds of the

critical temperature. Hence within narrow limits Kopp's determinations

were carried out under coincident conditions, and therefore any

regularities presented by the critical volumes should be revealed in

the specific volumes at the boiling-point.

Volume relations of solids.

The connexion between the density and chemical composition of solids

has not been investigated with the same completeness as in the case of

gases and liquids. The relation between the atomic volumes and the

atomic weights of the solid elements exhibits the periodicity which

generally characterizes the elements. The molecular volume is additive

in certain cases, in particular of analogous compounds of simple

constitution. For instance, constant differences are found between the

chlorides, bromides and iodides of sodium and potassium:--

+------------+------+-------------+------+----------------+

+------------+------+-------------+------+----------------+

| KCl = 37.4 | 6.9 | NaCl = 27.1 | 6.7 | 10.3 |

| KBr = 44.3 | 9.7 | NaBr = 33.8 | 9.7 | 10.5 |

| KI = 54.0 | | NaI = 43.5 | | 10.5 |

+------------+------+-------------+------+----------------+

According to H. Schroeder the silver salts of the fatty acids exhibit

additive relations; an increase in the molecule of CH2 causes an

increase in the molecular volume of about 15.3.

_Thermal Relations._

_Specific Heat and Composition._---The nature and experimental

determination of specific heats are discussed in the article

CALORIMETRY; here will be discussed the relations existing between the

heat capacities of elements and compounds.

Specific heat of gases.

In the article THERMODYNAMICS it is shown that the amount of heat

required to raise a given weight of a gas through a certain range of

temperature is different according as the gas is maintained at constant

pressure, the volume increasing, or at constant volume, the pressure

increasing. A gas, therefore, has two specific heats, generally denoted

by C_p and C_v, when the quantity of gas taken as a unit is one gramme

molecular weight, the range of temperature being 1 deg. C. It may be

shown that C_p - C_v = R, where R is the gas-constant, i.e. R in the

equation PV = RT. From the ratio C_p/C_v conclusions may be drawn as to

the molecular condition of the gas. By considerations based on the

kinetic theory of gases (see MOLECULE) it may be shown that when no

energy is utilized in separating the atoms of a molecule, this ratio is

5/3 = 1.67. If, however, an amount of energy a is taken up in separating

atoms, the ratio is expressible as C_p/C_v = (5+a)/(3+a), which is

obviously smaller than 5/3, and decreases with increasing values of a.

These relations may be readily tested, for the ratio C_p/C_v is capable

of easy experimental determination. It is found that mercury vapour,

helium, argon and its associates (neon, krypton, &c.) have the value

1.67; hence we conclude that these gases exist as monatomic molecules.

Oxygen, nitrogen, hydrogen and carbon monoxide have the value 1.4; these

gases have diatomic molecules, a fact capable of demonstration by other

means. Hence it may be inferred that this value is typical for diatomic

molecules. Similarly, greater atomic complexity is reflected in a

further decrease in the ratio C_p/C_v. The following table gives a

comparative view of the specific heats and the ratio for molecules of

variable atomic content.

The abnormal specific heats of the halogen elements may be due to a

loosening of the atoms, a preliminary to the dissociation into

monatomic molecules which occurs at high temperatures. In the more

complex gases the specific heat varies considerably with temperature;

only in the case of monatomic gases does it remain constant. Le

Chatelier (_Zeit. f. phys. Chem._ i. 456) has given the formula C_p

= 6.5 + aT, where a is a constant depending on the complexity of the

molecule, as an expression for the molecular heat at constant pressure

at any temperature T (reckoned on the absolute scale). For a further

discussion of the ratio of the specific heats see MOLECULE.

+-----------+-------------------------------+-------+-------+---------+

+-----------+-------------------------------+-------+-------+---------+

| Monatomic | Hg, Zn, Cd, He, Ar, &c. | 5 | 3 | 1.66 |

+-----------+-------------------------------+-------+-------+---------+

| | H2, 02, N2 (0 deg.-200 deg.) | 6.83 | 4.83 | 1.41 |

| Diatomic | Br2, I2 (0 deg.-200 deg.) | 8.6 | 6.6 | 1.30 |

| | Cl2,HCl, HBr, HI, NO, CO | | | 1.41 |

+-----------+-------------------------------+-------+-------+---------+

| Triatomic | H2O, H2S, N2O, CO2 | 9.2 | 7.2 | 1.28 |

+-----------+-------------------------------+-------+-------+---------+

| Tetratomic| As4, P4 | 13.4 | 11.4 | 1.175 |

| | NH3, C2H2 | 11.6 | 9.6 | 1.21 |

+-----------+-------------------------------+-------+-------+---------+

| Pentatomic| CHCl3 | 14 | 12 | 1.17 |

+-----------+-------------------------------+-------+-------+---------+

| Hexatomic | C2H4, C2H3Br | 16.4 | 14.4 | 1.14 |

+-----------+-------------------------------+-------+-------+---------+

_Specific Heats of Solids._--The development of the atomic theory and

the subsequent determination of atomic weights in the opening decades of

the 19th century inspired A.T. Petit and P.L. Dulong to investigate

relations (if any) existing between specific heats and the atomic

weight. Their observations on the solid elements led to a remarkable

generalization, now known as Dulong and Petit's law. This states that

"the atomic heat (the product of the atomic weight and specific heat) of

all elements is a constant quantity." The value of this constant when H

= 1 is about 6.4; Dulong and Petit, using O = 1, gave the value .38, the

specific heat of water being unity in both cases. This law--purely

empirical in origin--was strengthened by Berzelius, who redetermined

many specific heats, and applied the law to determine the true atomic

weight from the equivalent weight. At the same time he perceived that

specific heats varied with temperature and also with allotropes, e.g.

graphite and diamond. The results of Berzelius were greatly extended by

Hermann Kopp, who recognized that carbon, boron and silicon were

exceptions to the law. He regarded these anomalies as solely due to the

chemical nature of the elements, and ignored or regarded as

insignificant such factors as the state of aggregation and change of

specific heat with temperature.

The specific heats of carbon, boron and silicon subsequently formed

the subject of elaborate investigations by H.F. Weber, who showed that

with rise of temperature the specific (and atomic) heat increases,

finally attaining a fairly constant value; diamond, graphite and the

various amorphous forms of carbon having the value about 5.6 at 1000

deg., and silicon 5.68 at 232 deg.; while he concluded that boron

attained a constant value of 5.5. Niison and Pettersson's observations

on beryllium and germanium have shown that the atomic heats of these

metals increase with rise of temperature, finally becoming constant

with a value 5.6. W.A. Tilden (_Phil. Trans._, 1900, p. 233)

investigated nickel and cobalt over a wide range of temperature (from

-182.5 deg. to 100 deg.); his results are:--

+--------------------------------+---------+---------+

| | Cobalt. | Nickel. |

+--------------------------------+---------+---------+

| From -182.5 deg. to -78.4 deg. | 4.1687 | 4.1874 |

| -78.4 deg. to 15 deg. | 5.4978 | 5.6784 |

| 15 deg. to 100 deg. | 6.0324 | 6.3143 |

+--------------------------------+---------+---------+

It is evident that the atomic heats of these intimately associated

elements approach nearer and nearer as we descend in temperature,

approximating to the value 4. Other metals were tested in order to

determine if their atomic heats approximated to this value at low

temperatures, but with negative results.

It is apparent that the law of Dulong and Petit is not rigorously

true, and that deviations are observed which invalidate the law as

originally framed. Since the atomic heat of the same element varies

with its state of aggregation, it must be concluded that some factor

taking this into account must be introduced; moreover, the variation

of specific heat with temperature introduces another factor.

We now proceed to discuss molecular heats of compounds, that is, the

product of the molecular weight into the specific heat. The earliest

generalization in this direction is associated with F.E. Neumann, who,

in 1831, deduced from observations on many carbonates (calcium,

magnesium, ferrous, zinc, barium and lead) that stoichiometric

quantities (equimolecular weights) of compounds possess the same heat

capacity. This is spoken of as "Neumann's law." Regnault confirmed

Neumann's observations, and showed that the molecular heat depended on

the number of atoms present, equiatomic compounds having the same

molecular heat. Kopp systematized the earlier observations, and, having

made many others, he was able to show that the molecular heat was an

additive property, i.e. each element retains the same heat capacity when

in combination as in the free state. This has received confirmation by

the researches of W.A. Tilden (_Phil. Trans._, 1904, 203 A, p.139) for

those elements whose atomic heats vary considerably with temperature.

The specific heat of a compound may, in general, be calculated from

the specific heats of its constituent elements. Conversely, if the

specific heats of a compound and its constituent elements, except one,

be known, then the unknown atomic heat is readily deducible.

Similarly, by taking the difference of the molecular heats of

compounds differing by one constituent, the molecular (or atomic) heat

of this constituent is directly obtained. By this method it is shown

that water, when present as "water of crystallization," behaves as if

it were ice.

_Deductions from Dulong and Petit's Law._--Denoting the atomic weight by

W and the specific heat by s, Dulong and Petit's law states that 6.4 =

Ws. Thus if s be known, an approximate value of W is determinate. In the

determination of the atomic weight of an element two factors must be

considered: (1) its equivalent weight, i.e. the amount which is

equivalent to one part of hydrogen; and (2) a factor which denotes the

number of atoms of hydrogen which combines with or is equivalent to one

atom of the particular element. This factor is termed the valency. The

equivalent weight is capable of fairly ready determination, but the

settlement of the second factor is somewhat more complex, and in this

direction the law of atomic heats is of service. To take an example: 38

parts of indium combine with 35.4 parts of chlorine; hence, if the

formula of the chloride be InCl, InCl2 or InCl3, indium has the atomic

weights 38, 76 or 114. The specific heat of indium is 0.057; and the

atomic heats corresponding to the atomic weights 38, 76 and 114 are 3.2,

4.3, 6.5. Dulong and Petit's law thus points to the value 114, which is

also supported by the position occupied by this element in the periodic

classification. C. Winkler decided the atomic weight of germanium by

similar reasoning.

_Boiling-Point and Composition._--From the relation between the critical

constants P_kV_k/T_k = (1/3.7)R or T_k/P_k = 3.7V_k/R, and since V_k is

proportional to the volume at absolute zero, the ratio T_k/P_k should

exhibit additive relations. This ratio, termed by Guye the critical

coefficient, has the following approximate values:--

Double Triple

C. H. Cl. -O-. =O. N. N=. P. linkage. linkage.

1.35 0.57 2.66 0.87 1.27 1.6 1.86 3.01 0.88 1.03

Since at the boiling-point under atmospheric pressure liquids are in

corresponding states, the additive nature of the critical coefficient

should also be presented by boiling-points. It may be shown

theoretically that the absolute boiling-point is proportional to the

molecular volume, and, since this property is additive, the

boiling-point should also be additive.

These relations have been more thoroughly tested in the case of

organic compounds, and the results obtained agree in some measure with

the deductions from molecular volumes. In general, isomers boil at

about the same temperature, as is shown by the isomeric esters

C9H18O2:--

Methyl octoate 192.9 deg. Amyl butyrate 184.8 deg.

Ethyl heptoate 187.1 deg. Heptyl acetate 191.3 deg.

Propyl hexoate 185.5 deg. Octyl formate 198.1 deg.

Butyl pentoate 185.8

Equal increments in the molecule are associated with an equal rise in

the boiling-point, but this increment varies in different homologous

series. Thus in the normal fatty alcohols, acids, esters, nitriles and

ketones, the increment per CH2 is 19 deg.-21 deg.; in the aldehydes it

is 26 deg.-27 deg.. In the aromatic compounds there is no regularity

between the increments due to the introduction of methyl groups into

the benzene nucleus or side chains; the normal value of 20 deg.-21

deg. is exhibited, however, by pyridine and its derivatives. The

substitution of a hydrogen atom by the hydroxyl group generally

occasions a rise in boiling-point at about 100 deg.. The same increase

accompanies the introduction of the amino group into aromatic nuclei.

Constitutive influences.

While certain additive relations hold between some homologous series,

yet differences occur which must be referred to the constitution of

the molecule. As a general rule, compounds formed with a great

evolution of heat have high boiling-points, and vice versa. The

introduction of negative groups into a molecule alters the

boiling-point according to the number of negative groups already

present. This is shown in the case of the chloracetic acids:

Diff.

CH3CO2H = 118 deg.

67 deg.

ClCH2.CO2H = 185 deg.

10 deg.

Cl2CH.CO2H = 195 deg.

3 deg.

Cl3C.CO2H = 195 deg.-200 deg.

According to van 't Hoff the substitution of chlorine atoms into a

methyl group occasions the following increments:--

Cl in CH3 66 deg.

Cl " CH2Cl 39 deg.

Cl " CHCl2 13 deg.

The introduction of chlorine, however, may involve a fall in the

boiling-point, as is recorded by Henry in the case of the chlorinated

acetonitriles:--

NC.CH3. NC.CH2Cl. NC.CHC12. NC.CC13.

81 deg. 123 deg. 112 deg. 83 deg.

42 deg. -11 deg. -29 deg.

The replacement of one negative group by another is accompanied by a

change in the boiling-point, which is independent of the compound in

which the substitution is effected, and solely conditioned by the

nature of the replaced and replacing groups. Thus bromine and iodine

replace chlorine with increments of about 22 deg. and 50 deg.

respectively.

A factor of considerable importance in determining boiling-points of

isomers is the symmetry of the molecule. Referring to the esters

C9H18O2 previously mentioned, it is seen that the highest

boiling-points belong to methyl octoate and octyl formate, the least

symmetrical, while the minimum belongs to amyl butyrate, the most

symmetrical. The isomeric pentanes also exhibit a similar relation

CH3(CH2)4CH3 = 38 deg., (CH3)2CHC2H5 = 30 deg., (CH3)4C = 9.5 deg..

For a similar reason secondary alcohols boil at a lower temperature

than the corresponding primary, the difference being about 19 deg..

A.E. Earp (_Phil. Mag._, 1893 [5], 35, p. 458) has shown that, while

an increase in molecular weight is generally associated with a rise in

the boiling-point, yet the symmetry of the resulting molecule may

exert such a lowering effect that the final result is a diminution in

the boiling-point. The series H2S = -61 deg., CH3SH = 21 deg., (CH3)2S

= 41 deg. is an example; in the first case, the molecular weight is

increased and the symmetry diminished, the increase of boiling-point

being 82 deg.; in the second case the molecular weight is again

increased but the molecule assumes a more symmetrical configuration,

hence the comparatively slight increase of 20 deg.. A similar

depression is presented by methyl alcohol (67 deg.) and methyl ether

(-23 deg.).

Among the aromatic di-substitution derivatives the _ortho_ compounds

have the highest boiling-point, and the _meta_ boil at a higher, or

about the same temperature as the _para_ compounds. Of the

tri-derivatives the symmetrical compounds boil at the lowest

temperature, the asymmetric next, and the vicinal at the highest.

An ethylenic or double carbon union in the aliphatic hydrocarbons has,

apparently, the same effect on the boiling-point as two hydrogen

atoms, since the compounds C_{n}H_{2n+2} and C_{n}H_{2n} boil at about

the same temperature. An acetylenic or triple linkage is associated

with a rise in the boiling-point; for example, propargyl compounds

boil about 19.5 deg. higher than the corresponding propyl compound.

Certain regularities attend the corresponding property of the

melting-point. A rule applicable to organic compounds, due to Adolf v.

Baeyer and supported by F.S. Kipping (_Jour. Chem. Soc._, 1893, 63,

p.465) states, that the melting-point of any odd member of a

homologous series is lower than the melting-point of the even member

containing one carbon atom less. This is true of the fatty acid

series, and the corresponding ketones and alcohols, and also of the

succinic acid series. Other regularities exist, but generally with

many exceptions. It is to be noted that although the correlation of

melting-point with constitution has not been developed to such an

extent as the chemical significance of other physical properties, the

melting-point is the most valuable test of the purity of a substance,

a circumstance due in considerable measure to the fact that impurities

always tend to lower the melting-point.

_Heat of Combustion and Constitution._--In the article THERMOCHEMISTRY a

general account of heats of formation of chemical compounds is given,

and it is there shown that this constant measures the stability of the

compound. In organic chemistry it is more customary to deal with the

"heat of combustion," i.e. the heat evolved when an organic compound is

completely burned in oxygen; the heat of formation is deduced from the

fact that it is equal to the heats of formation of the products of

combustion less the observed heat of combustion. The researches of

Julius Thomsen and others have shown that in many cases definite

conclusions regarding constitution can be drawn from quantitative

measurements of the heats of combustion; and in this article a summary

of the chief results will be given.

The identity of the four valencies of the carbon atom follows from the

fact that the heats of combustion of methane, ethane, propane,

trimethyl methane, and tetramethyl methane, have a constant difference

in the order given, viz. 158.6 calories; this means that the

replacement of a hydrogen atom by a methyl group is attended by a

constant increase in the heat of combustion. The same difference

attends the introduction of the methyl group into many classes of

compounds, for example, the paraffins, olefines, acetylenes, aromatic

hydrocarbons, alcohols, aldehydes, ketones and esters, while a

slightly lower value (157.1) is found in the case of the halogen

compounds, nitriles, amines, acids, ethers, sulphides and nitro

compounds. It therefore appears that the difference between the heats

of combustion of two adjacent members of a series of homologous

compounds is practically a constant, and that this constant has two

average values, viz. 158.6 and 157.1.

An important connexion between heats of combustion and constitution is

found in the investigation of the effect of single, double and triple

carbon linkages on the thermochemical constants. If twelve grammes of

amorphous carbon be burnt to carbon dioxide under constant volume, the

heat evolved (96.96 cal.) does not measure the entire thermal effect,

but the difference between this and the heat required to break down

the carbon molecule into atoms. If the number of atoms in the carbon

molecule be denoted by n, and the heat required to split off each atom

from the molecule by d, then the total heat required to break down a

carbon molecule completely into atoms is nd. It follows that the true

heat of combustion of carbon, i.e. the heat of combustion of one

gramme-atom, is 96.96 + d. The value of d can be evaluated by

considering the combustion of amorphous carbon to carbon monoxide and

carbon dioxide. In the first case the thermal effect of 58.58 calories

actually observed must be increased by 2d to allow for the heat

absorbed in splitting off two gramme-atoms of carbon; in the second

case the thermal effect of 96.96 must be increased by d as above. Now

in both cases one gramme-molecule of oxygen is decomposed, and the two

oxygen atoms thus formed are combined with two carbon valencies. It

follows that the thermal effects stated above must be equal, i.e.

58.58 + 2d = 96.96 + d, and therefore d = 38.38. The absolute heat of

combustion of a carbon _atom_ is therefore 135.34 calories, and this

is independent of the form of the carbon burned.

Consider now the combustion of a hydrocarbon of the general formula

C_{n}H_{2m}. We assume that each carbon atom and each hydrogen atom

contributes equally to the thermal effect. If [alpha] be the heat

evolved by each carbon atom, and [beta] that by each hydrogen atom,

the thermal effect may be expressed as H = n[alpha] + 2m[beta] - A,

where A is the heat required to break the molecule into its

constituent atoms. If the hydrocarbon be saturated, i.e. only contain

single carbon linkages, then the number of such linkages is 2n - m,

and if the thermal effect of such a linkage be X, then the term A is

obviously equal to (2n - m)X. The value of H then becomes H = n[alpha]

+ 2m[beta] - (2n - m)X or n[xi] + m[nu], where [xi] and [nu] are

constants. Let double bonds be present, in number p, and let the

energy due to such a bond be Y. Then the number of single bonds is 2n

- m - 2p, and the heat of combustion becomes H1 = n[xi] + m[nu] +

p(2X - Y). If triple bonds, q in number, occur also, and the energy of

such a bond be Z, the equation for H becomes

H = n[xi] + m[nu] + p(2X - Y) + q(3X - Z).

This is the general equation for calculating the heat of combustion of

a hydrocarbon. It contains four independent constants; two of these

may be calculated from the heats of combustion of saturated

hydrocarbons, and the other two from the combustion of hydrocarbons

containing double and triple linkages. By experiment it is found that

the thermal effect of a double bond is much less than the effect of

two single bonds, while a triple bond has a much smaller effect than

three single bonds. J. Thomsen deduces the actual values of X, Y, Z to

be 14.71, 13.27 and _zero_; the last value he considers to be in

agreement with the labile equilibrium of acetylenic compounds. One of

the most important applications of these values is found in the case

of the constitution of benzene, where Thomsen decides in favour of the

Claus formula, involving nine single carbon linkages, and rejects the

Kekule formula, which has three single and three double bonds (see

section IV.).

The thermal effects of the common organic substituents have also been

investigated. The thermal effect of the "alcohol" group C.OH may be

determined by finding the heat of formation of the alcohol and

subtracting the thermal effects of the remaining linkages in the

molecule. The average value for primary alcohols is 44.67 cal., but

many large differences from this value obtain in certain cases. The

thermal effects increase as one passes from primary to tertiary

alcohols, the values deduced from propyl and isopropyl alcohols and

trimethyl carbinol being:--primary = 45.08, secondary = 50.39,

tertiary = 60.98. The thermal effect of the aldehyde group has the

average value 64.88 calories, i.e. considerably greater than the

alcohol group. The ketone group corresponds to a thermal effect of

53.52 calories. It is remarkable that the difference in the heats of

formation of ketones and the paraffin containing one carbon atom less

is 67.94 calories, which is the heat of formation of carbon monoxide

at constant volume. It follows therefore that two hydrocarbon radicals

are bound to the carbon monoxide residue with the same strength as

they combine to form a paraffin. The average value for the carboxyl

group is 119.75 calories, i.e. it is equal to the sum of the thermal

effects of the aldehyde and carbonyl groups.

The thermal effects of the halogens are: chlorine = l5.13 calories,

bromine = 7.68; iodine = -4.25 calories. It is remarkable that the

position of the halogen in the molecule has no effect on the heat of

formation; for example, chlorpropylene and allylchloride, and also

ethylene dichloride and ethylidene dichloride, have equal heats of

formation. The thermal effect of the ether group has an average value

of 34.31 calories. This value does not hold in the case of methylene

oxide if we assign to it the formula H2C.[O.CH2], but if the formula

H2C.O.CH2 (which assumes the presence of two free valencies) be

accepted, the calculated and observed heats of formation are in

agreement.

The combination of nitrogen with carbon may result in the formation of

nitriles, cyanides, or primary, secondary or tertiary amines. Thomsen

deduced that a single bond between a carbon and a nitrogen gramme-atom

corresponds to a thermal effect of 2.77 calories, a double bond to

5.44, and a treble bond to 8.31. From this he infers that cyanogen is

C:N.N:C and not N:C-C:N, that hydrocyanic acid is HC.N, and

acetonitrile CH3.C:N. In the case of the amines he decides in favour

of the formulae

H2C:NH3 H2C H2C

\\ \\

NH2 NH.CH3

/ /

H3C H3C

primary, secondary, tertiary.

These involve pentavalent nitrogen. These formulae, however, only

apply to aliphatic amines; the results obtained in the aromatic series

are in accordance with the usual formulae.

_Optical Relations_.

_Refraction and Composition._--Reference should be made to the article

REFRACTION for the general discussion of the phenomenon known as the

refraction of light. It is there shown that every substance, transparent

to light, has a definite refractive index, which is the ratio of the

velocity of light _in vacuo_ to its velocity in the medium to which the

refractive index refers. The refractive index of any substance varies

with (1) the wave-length of the light; (2) with temperature; and (3)

with the state of aggregation. The first cause of variation may be at

present ignored; its significance will become apparent when we consider

dispersion (_vide infra_).The second and third causes, however, are of

greater importance, since they are associated with the molecular

condition of the substance; hence, it is obvious that it is only from

some function of the refractive index which is independent of

temperature variations and changes of state (i.e. it must remain

constant for the same substance at any temperature and in any form) that

quantitative relations between refractivity and chemical composition can

be derived.

The pioneer work in this field, now frequently denominated

"spectro-chemistry," was done by Sir Isaac Newton, who, from theoretical

considerations based on his corpuscular theory of light, determined the

function (n^2-1), where n is the refractive index, to be the expression

for the refractive power; dividing this expression by the density (d), he

obtained (n^2-1)/d, which he named the "absolute refractive power." To

P.S. Laplace is due the theoretical proof that this function is

independent of temperature and pressure, and apparent experimental

confirmation was provided by Biot and Arago's, and by Dulong's

observations on gases and vapours. The theoretical basis upon which this

formula was devised (the corpuscular theory) was shattered early in the

19th century, and in its place there arose the modern wave theory which

theoretically invalidates Newton's formula. The question of the

dependence of refractive index on temperature was investigated in 1858

by J.H. Gladstone and the Rev. T.P. Dale; the more simple formula

(n-1)/d, which remained constant for gases and vapours, but exhibited

slight discrepancies when liquids were examined over a wide range of

temperature, being adopted. The subject was next taken up by Hans

Landolt, who, from an immense number of observations, supported in a

general way the formula of Gladstone and Dale. He introduced the idea of

comparing the refractivity of equimolecular quantities of different

substances by multiplying the function (n-1)/d by the molecular weight

(M) of the substance, and investigated the relations of chemical

grouping to refractivity. Although establishing certain general

relations between atomic and molecular refractions, the results were

somewhat vitiated by the inadequacy of the empirical function which he

employed, since it was by no means a constant which depended only on the

actual composition of the substance and was independent of its physical

condition. A more accurate expression (n^2-1)/(n^2+2)d was suggested in

1880 independently and almost simultaneously by L.V. Lorenz of

Copenhagen and H.A. Lorentz of Leiden, from considerations based on the

Clausius-Mossotti theory of dielectrics.

Assuming that the molecules are spherical, R.J.E. Clausius and O.F.

Mossotti found a relation between the dielectric constant and the

space actually occupied by the molecules, viz. K = (1 + 2a)/(1 - a),

or a = (K - 1)/(K + 2), where K is the dielectric constant and a the

fraction of the total volume actually occupied by matter. According to

the electromagnetic theory of light K = N^2, where N is the refractive

index for rays of infinite wave-length. Making this substitution, and

dividing by d, the density of the substance, we obtain a/d = (N^2 -

1)/(N^2 + 2 )d. Since a/d is the real specific volume of the molecule,

it is therefore a constant; hence (N^2 - 1)/(N^2 + 2)d is also a

constant and is independent of all changes of temperature, pressure,

and of the state of aggregation. To determine N recourse must be made

to Cauchy's formula of dispersion (q.v.), n = A + B/[lambda]^2 +

C/[lambda]^4 + ... from which, by extrapolation, [lambda] becoming

infinite, we obtain N = A. In the case of substances possessing

anomalous dispersion, the direct measurement of the refractive index

for Hertzian waves of very long wave-length may be employed.

It is found experimentally that the Lorenz and Lorentz function holds

fairly well, and better than the Gladstone and Dale formula. This is

shown by the following observations of Ruhlmann on water, the light used

being the D line of the spectrum:--

+------+------------+-----------------------+

| t. | (n - 1)/d. | (n^2 - 1)/(n^2 + 2)d. |

+------+------------+-----------------------+

| 0 | 0.3338 | 0.2061 |

| 10 | 0.3338 | 0.2061 |

| 20 | 0.3336 | 0.2061 |

| 90 | 0.3321 | 0.2059 |

| 100 | 0.3323 | 0.2061 |

+------+------------+-----------------------+

Eykmann's observations also support the approximate constancy of the

Lorenz-Lorentz formula over wide temperature differences, but in some

cases the deviation exceeds the errors of observation. The values are

for the H[alpha] line:--

+----------------------------+--------------+-----------------------+

| Substance. | Temp. | (n^2 - 1)/(n^2 + 2)d. |

+----------------------------+--------------+-----------------------+

| Isosafrol, C10H10O2 | / 17.6 deg. | 0.2925 |

| | \ 141.2 deg. | 0.2962 |

| Diphenyl ethylene, C14H12 | / 22 deg. | 0.3339 |

| | \ 143.4 deg. | 0.3382 |

| Quinoline, C9H7N | / 16.2 deg. | 0.3187 |

| | \ 141 deg. | 0.3225 |

+----------------------------+--------------+-----------------------+

The empirical formula (n^2 - 1)/(n^2 + 0.4)d apparently gives more

constant values with change of temperature than the Lorenz-Lorentz form.

The superiority of the Lorenz-Lorentz formula over the Gladstone and

Dale formula for changes of state is shown by the following observations

of Bruhl (_Zeit. f. phys. Chem._, 1891, 71, p. 4). The values are for

the D line:--

+-------------------+---------+---------------------+---------------------+

| Substance. | Temp. +----------+----------+---------+-----------+

+-------------------+---------+----------+----------+---------+-----------+

| Water | 10 deg. | 0.3101 | 0.3338 | 0.2068 | 0.2061 |

| Carbon disulphide | 10 deg. | 0.4347 | 0.4977 | 0.2898 | 0.2805 |

| Chloroform | 10 deg. | 0.2694 | 0.3000 | 0.1796 | 0.1790 |

+-------------------+---------+----------+----------+---------+-----------+

[Sidenote: Additive relations.]

Landolt and Gladstone, and at a later date J.W. Bruhl, have

investigated the relations existing between the refractive power and

composition. To Landolt is due the proof that, in general, isomers,

i.e. compounds having the same composition, have equal molecular

refractions, and that equal differences in composition are associated

with equal differences in refractive power. This is shown in the

following table (the values are for H[alpha]):

+------------------------------+--------+-----------------+--------+-----------+

+------------------------------+--------+-----------------+--------+-----------+

| Ethylene chloride \ C2H4Cl2 | 20.96 | Acetic acid | 12.93 | / 4.49 |

| Ethylidene chloride/ | 21.08 | Propionic acid | 17.42 | \ |

| Fumaric acid \ C4H4O4 | 70.89 | Butyric acid | 22.01 | / 4.59 |

| Maleic acid / | 70.29 | | | \ |

| o-Cresol \ | 32.52 | Acetaldehyde | 11.50 | / 4.43 |

| m-Cresol } C7H8O | 32.56 | Propionaldehyde | 15.93 | \ |

| p-Cresol / | 32.57 | Butylaldehyde | 20.52 | / 4.59 |

+------------------------------+--------+-----------------+--------+-----------+

Additive relations undoubtedly exist, but many discrepancies occur

which may be assigned, as in the case of molecular volumes, to

differences in constitution. Atomic refractions may be obtained either

directly, by investigating the various elements, or indirectly, by

considering differences in the molecular refractions of related

compounds. The first method needs no explanation. The second method

proceeds on the same lines as adopted for atomic volumes. By

subtracting the value for CH2, which may be derived from two

substances belonging to the same homologous series, from the molecular

refraction of methane, CH4, the value of hydrogen is obtained;

subtracting this from CH2, the value of carbon is determined.

Hydroxylic oxygen is obtained by subtracting the molecular refractions

of acetic acid and acetaldehyde. Similarly, by this method of

differences, the atomic refraction of any element may be determined.

It is found, however, that the same element has not always the same

atomic refraction, the difference being due to the nature of the

elements which saturate its valencies. Thus oxygen varies according as

whether it is linked to hydrogen (hydroxylic oxygen), to two atoms of

carbon (ether oxygen), or to one carbon atom (carbonyl oxygen);

similarly, carbon varies according as whether it is singly, doubly, or

trebly bound to carbon atoms.

A table of the atomic refractions and dispersions of the principal

elements is here given:--

+--------------------------+--------+-------+--------+---------------+

+--------------------------+--------+-------+--------+---------------+

| Hydrogen | 1.103 | 1.051 | 1.139 | 0.036 |

| Oxygen, hydroxyl | 1.506 | 1.521 | 1.525 | 0.019 |

| Oxygen, ether | 1.655 | 1.683 | 1.667 | 0.012 |

| Oxygen, carbonyl | 2.328 | 2.287 | 2.414 | 0.086 |

| Chlorine | 6.014 | 5.998 | 6.190 | 0.176 |

| Bromine | 8.863 | 8.927 | 9.211 | 0.348 |

| Iodine |13.808 |14.12 | 14.582 | 0.774 |

| Carbon (singly bound) | 2.365 | 2.501 | 2.404 | 0.039 |

| Double linkage of carbon | 1.836 | 1.707 | 1.859 | 0.23 |

| Triple | 2.22 | | 2.41 | 0.19 |

| and only to carbon | 2.76 | | 2.95 | 0.19 |

+--------------------------+--------+-------+--------+---------------+

_Dispersion and Composition._---In the preceding section we have seen

that substances possess a definite molecular (or atomic) refraction

for light of particular wave-length; the difference between the

refractions for any two rays is known as the molecular (or atomic)

dispersion. Since molecular refractions are independent of temperature

and of the state of aggregation, it follows that molecular dispersions

must be also independent of these conditions; and hence quantitative

measurements should give an indication as to the chemical composition

of substances. This subject has been principally investigated by

Bruhl; he found that molecular dispersions of liquids and gases were

independent of temperature, and fairly independent of the state of

aggregation, but that no simple connexion exists between atomic

refractions and dispersions (see preceding table). He also showed how

changes in constitution effected dispersions to a far greater extent

than they did refractions; thus, while the atomic dispersion of carbon

is 0.039, the dispersions due to a double and treble linkage is 0.23

and 0.19 respectively.

_Colour and Constitution._--In this article a summary of the theories

which have been promoted in order to connect the colour of organic

compounds with their constitution will be given, and the reader is

referred to the article COLOUR for the physical explanation of this

property, and to VISION for the physiological and psychological

bearings. A clear distinction must be drawn between colour and the

property of dyeing; all coloured substances are not dyes, and it is

shown in the article DYEING that the property of entering into chemical

or physical combination with fibres involves properties other than those

essential to colour. At the same time, however, all dyestuffs are

coloured substances.

A survey of coloured substances led O.N. Witt in 1876 to formulate his

"chromophore-auxochrome" theory. On this theory colour is regarded as

due to the presence of a "chromophore," and dyeing power to an

"auxochrome"; the latter by itself cannot produce colour or dyeing

power, but it is only active in the presence of a chromophore, when it

intensifies the colour and confers the property of dyeing. The

principal chromophores are the azo, -N=N-, azoxy, =N2O, nitro, -NO2,

nitroso, -NO, and carbonyl, =CO, groups. The azo-group is particularly

active, both the aliphatic and aromatic compounds being coloured. The

simplest aliphatic compounds, such as diazo-methane, diazo-ethane, and

azo-formic acid, are yellow; the diamide of the latter acid is

orange-red. Of the aromatic compounds azo-benzene is bright

orange-red, and [alpha]-azo-naphthalene forms red needles or small

steel-blue prisms. The azo-group, however, has little or no colouring

effect when present in a ring system, such as in cinnolene,

phthalazine and tolazone. The nitro group has a very important action

mainly on account of the readiness with which it can be introduced

into the molecule, but its effect is much less than that of the azo

group. The colour produced is generally yellow, which, in accordance

with a general rule, is intensified with an increase in the number of

groups; compare, for example, mono-, di-and tri-nitrobenzene. The

nitroso group is less important. The colour produced is generally of a

greenish shade; for example, nitrosobenzene is green when fused or in

solution (when crystalline, it is colourless), and dinitrosoresorcin

has been employed as a dyestuff under the names "solid green" and

"chlorine." The carbonyl group by itself does not produce colour, but

when two adjacent groups occur in the molecule, as for example in the

a-diketones (such as di-acetyl and benzil), a yellow colour is

produced. It also acts as a chromogenic centre when double bonds or

ethylenic linkages are present, as in fluorene ketone or fluorenone. A

more complex chromophoric group is the triple ethylenic grouping

=C\

>C=,

=C/

the introduction of which was rendered necessary by the

discovery of certain coloured hydrocarbons. As a general rule,

hydrocarbons are colourless; the exceptions include the golden yellow

acenaphthylene, the red bidiphenylene-ethylene, and the derivatives of

fulvene

CH:CH \

>CH2,

CH:CH /

which have been discussed by J. Thiele (_Ber._, 1900, 33, p. 666).

This grouping is not always colour-producing, since diphenyl is

colourless.

The most important auxochromes are the hydroxyl (-OH) and amino (-NH2)

groups. According to the modern theory of auxochromic action, the

introduction of a group into the molecule is accompanied by some

strain, and the alteration in colour produced is connected with the

magnitude of the strain. The amino group is more powerful than the

hydroxyl, and the substituted amino group more powerful still; the

repeated substitution of hydroxyl groups sometimes causes an

intensification and sometimes a diminution of colour.

We may here notice an empirical rule formulated by Nietzski in

1879:--the simplest colouring substances are in the greenish-yellow

and yellow, and with increasing molecular weight the colour passes

into orange, red, violet, blue and green. This rule, however, is by no

means perfect. Examination of the absorption spectra of coloured

compounds shows that certain groupings displace the absorption bands

in one direction, and other groupings in the other. If the bands be

displaced towards the violet, involving a regression through the

colours mentioned above, the group is said to be "hypsochromic"; if

the reverse occurs the group is "bathochromic." It may be generally

inferred that an increase in molecular weight is accompanied by a

change in colour in the direction of the violet.

Auxochromic groups generally aid one another, i.e. the tint deepens as

the number of auxochromes increases. Also the relative position of the

auxochrome to the chromophore influences colour, the ortho-position

being generally the most powerful. Kauffmann (_Ber._, 1906, 39, p.

1959) attempted an evaluation of the effects of auxochromic groups by

means of the magnetic optical constants. The method is based on the

supposition that the magnetic rotation measures the strain produced in

the molecule by an auxochrome, and he arranges the groups in the

following order:--

.OCOCH3 .OCH3 .NHCOCH3 .NH2 .N(CH3)2 .N(C2H5)2

-0.260 1.459 1.949 3.821 8.587 8.816

The phenomena attending the salt formation of coloured and colouring

substances are important. The chromophoric groups are rarely strongly

acid or basic; on the other hand, the auxochromes are strongly acid or

basic and form salts very readily. Notable differences attend the

neutralization of the chromophoric and auxochromic groups. With basic

substances, the chromophoric combination with a colourless acid is

generally attended by a deepening in colour; auxochromic combination,

on the other hand, with a lessening. Examples of the first case are

found among the colourless acridines and quinoxalines which give

coloured salts; of the second case we may notice the colourless

hydrochloride and sulphate of the deep yellow o-aminobenzophenone.

With acid substances, the combination with "colourless" metals, i.e.

metals producing colourless salts with acids, is attended by colour

changes contrary to those given above, auxochromic combination being

accompanied by a deepening, and chromophoric by a lessening of the

tint.

Mention may be made of the phenomenon of halochromism, the name given

to the power of colourless or faintly-coloured substances of combining

with acids to form highly-coloured substances without the necessary

production of a chromophoric group. The researches of Adolf von Baeyer

and Villiger, Kehrmann, Kauffmann and others, show that this property

is possessed by very many and varied substances. In many cases it may

be connected with basic oxygen, and the salt formation is assumed to

involve the passage of divalent into tetravalent oxygen. It seems that

intermolecular change also occurs, but further research is necessary

before a sound theory can be stated.

_Quinone Theory of Colour._--A theory of colour in opposition to the

Witt theory was proposed by Henry Armstrong in 1888 and 1892. This

assumed that all coloured substances were derivatives of ortho- or

para-quinone (see QUINONES), and although at the time of its

promotion little practical proof was given, yet the theory found wide

acceptance on account of the researches of many other chemists. It

follows on this theory that all coloured substances contain either of

the groupings

____ ____

/----\ /----\

/ \ / \

==< >== or < >== ,

\______/ \\ /

\____/ \\___/

the former being a para-quinonoid, the latter an ortho-quinonoid.

While very many coloured substances must obviously contain this

grouping, yet in many cases it is necessary to assume a simple

intermolecular change, while in others a more complex rearrangement of

bonds is necessary. Quinone, which is light yellow in colour, is the

simplest coloured substance on this theory. Hydrocarbons of similar

structure have been prepared by Thiele, for example, the orange-yellow

tetraphenyl-_para_-xylylene, which is obtained by boiling the bromide

C6H4[CBr(C6H5)2]2 with benzene and molecular silver. The quinonoid

structure of many coloured compounds has been proved experimentally,

as, for example, by Hewitt for the benzene-azo-phenols, and Hantzsch

for triaminotriphenyl methane and acridine derivatives; but, at the

same time, many substances cannot be so explained. A notable example

is provided by the phthaleins, which result by the condensation of

phthalic anhydride with phenols. In the free state these substances

are colourless, and were assumed to have the formula shown in 1.

Solution in dilute alkali was supposed to be accompanied by the

rupture of the lactone ring with the formation of the quinonoid salt

shown in 2.

/\ /\ O /\ /\

HO / \ / \ OH \\/ \ / \ OH

| | | | | | | |

\ / \ / \ / \ /\\ / \ /

\/ C \/ \/ C \/

/ \ /

C6H< >O C6H4<

\ / \

CO COON3

(1) (2)

Baeyer (_Ber._, 1905, 38, p. 569) and Silberrad (_Journ. Chem. Soc._,

1906, 89, p. 1787) have disputed the correctness of this explanation,

and the latter has prepared melliteins and pyromelliteins, which are

highly-coloured compounds produced from mellitic and pyromellitic

acids, and which cannot be formulated as quinones. Baeyer has

suggested that the nine carbon atom system of xanthone may act as a

chromophore. An alternative view, due to Green, is that the oxygen

atom of the xanthone ring is tetravalent, a supposition which permits

the formulation of these substances as ortho-quinonoids.

The theories of colour have also been investigated by Hantzsch, who

first considered the nitro-phenols. On the chromophore-auxochrome

theory (the nitro group being the chromophore, and the hydroxyl the

auxochrome) it is necessary in order to explain the high colour of the

metallic salts and the colourless alkyl and aryl derivatives to assume

that the auxochromic action of the hydroxyl group is only brought

strongly into evidence by salt formation. Armstrong, on the other

hand, assumed an intermolecular change, thus:--

/\ /\

HO / \ OH / \ == O

| | ===> | |

| | | |

\ / NO2 \ / == NO2Na.

\/ \/

The proof of this was left for Hantzsch, who traced a connexion with

the nitrolic acids of V. Meyer, which are formed when nitrous acid

acts on primary aliphatic nitro compounds. Meyer formulated these

compounds as nitroximes or nitro-isnitroso derivatives, viz.

R.C(NO2)(NOH). Hantzsch explains the transformation of the colourless

acid into red salts, which on standing yield more stable, colourless

salts, by the following scheme:--

NOH // \ NO2Na

// R.C O //

R.C ===> \ / ===> R.C

\\ N \\

NO2 // \ NO

O ONa

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