Alchemy Through Chemistry

 

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1.1 CHEMISTRY

Certain forms of matter such as wood and glass, water and gasoline, salt and sugar, coal and granite, and iron and gold differ markedly from each other in many properties. These differences in properties occur as a result of differences in the composition and structure of matter. This is not the whole story of chemistry for matter is not static. Much of our existence depends upon the changes which occur in matter. What changes in matter are taking place when iron rusts, milk sours, a storage battery produces an electric current, and food is digested and assimilated by the body? As we proceed in our study of chemistry, we shall look into the changes in the composition and structure of matter, the causes which produce these changes, the changes in energy which accompany them, and the principles and laws involved in these changes. The science of CHEMISTRY IS THE STUDY OF THE COMPOSITION, STRUCTURE, AND PROPERTIES OF MATTER, AND THE CHANGES WHICH IT UNDERGOES.

The science of chemistry has grown so vast that chemists usually specialize in one of several principal branches. ANALYTICAL CHEMISTRY is concerned with the identification, separation, and quantitative determination of the composition of different substances. PHYSICAL CHEMISTRY is primarily concerned with the structure of matter, energy changes, and the laws, principles, and theories which explain the transformations of one form of matter into another. ORGANIC CHEMISTRY is the branch dealing with the compounds of carbon. INORGANIC CHEMISTRY is concerned with the chemistry of the elements other than carbon and their compounds. BIOCHEMISTRY is the chemistry of the substances comprising living organisms, plants and animals. A course in GENERAL CHEMISTRY is a study of all the branches of chemistry and introduces the student to the entire field of the science.

It should be mentioned that the boundaries between the branches of chemistry are arbitrarily defined.

1.2 MATTER AND ENERGY

All objects in the universe are composed of matter. MATTER is that which occupies space and has mass. The property of occupying space is easily perceived by our senses of sight and touch. The property of MASS pertains to the quantity of matter that a body contains. The force required to give a body of matter a given acceleration, or the resistance of the body to being moved (inertia), is a measure of its mass.

ENERGY may be defined as the capacity for doing work. We are familiar with such forms of energy as heat, light, and electricity. A body of matter may possess POTENTIAL ENERGY by virtue of its position, condition, or composition. Water at the top of a waterfall possesses potential energy as a result of its position; by falling it can do work on an electric generator, causing it to turn. Coal possesses chemical energy, a kind of potential energy, because of its characteristic composition. As the coal burns it produces heat energy which can be used to drive a steam turbine.

If a body is allowed to fall, its potential energy is transformed into energy of motion. This energy of motion is the form of energy which can be transferred most readily to other systems and represents the ability to do work. The energy which a body possesses because of its motion is called KINETIC ENERGY.

1.3 LAW OF CONSERVATION OF MATTER

When a piece of magnesium metal is burned and the product (magnesium oxide) is collected and weighed, it is found that the product weighs more than the original piece of metal. If, however, the weight of the oxygen of the air that combines with the burning metal is taken into consideration, it may be shown that the weight of the product is equal to the weight of the reactants (magnesium and oxygen). This behavior of matter is in accord with what is called the LAW OF CONSERVATION OF MATTER: THERE IS NO DETECTABLE INCREASE OR DECREASE IN THE QUANTITY OF MATTER DURING A CHEMICAL CHANGE.

1.4 LAW OF CONSERVATION OF ENERGY

Chemical changes are always accompanied by the conversion of chemical energy into other forms of energy, or of other forms of energy into chemical energy. Usually, heat energy is evolved or absorbed, but sometimes the conversion involves light or electrical energy instead of or in addition to heat energy. Many transformations of energy do not involve chemical changes. Electrical energy may be converted into mechanical, light, heat, or chemical energy. Mechanical energy is converted into electrical energy in the dynamo. Potential and kinetic energy can be converted into one another. Many other conversions are possible. All the energy involved in any change appears in some form after the change is completed. This leads to the LAW OF CONSERVATION OF ENERGY: ENERGY CANNOT BE CREATED NOR DESTROYED, ALTHOUGH IT CAN BE CHANGED IN FORM.

Recently it has become necessary to regard matter and energy, not as distinct realities, but as different forms of a single reality. In the decomposition of atoms, such as takes place in the explosion of the atomic bomb, matter can be converted into energy. The reverse conversion of energy into matter can be demonstrated. This inter-conversion of mass and energy takes place in ordinary chemical reactions, but the mass of the products differs so slightly from the mass of the reactants that it is impossible to measure the difference experimentally. The magnitude of mass change is so small that we may say, for practical purposes, that the law of conservation of matter holds in ordinary chemical reactions. However, to be more precise, we may combine the separate laws of conservation of matter and energy: THE TOTAL QUANTITY OF MATTER AND ENERGY AVAILABLE IN THE UNIVERSE IS FIXED.

1.5 PHYSICAL STATES OF MATTER

Matter can exist in three different states, designated as solid, liquid, and gas, which can be distinguished by certain qualities. A substance in the SOLID STATE is rigid, possesses a definite shape, and has a volume which is very nearly independent of changes in conditions such as temperature and pressure.

A LIQUID, such as water, possesses the property of flowing to take the shape of its container except that it assumes a horizontal surface. Liquids are only slightly compressible and so for practical purposes have definite volumes.

A substance in the GASEOUS STATE takes both the shape and volume of its container. Gases are readily compressible and capable of infinite expansion.

1.6 CHEMICAL AND PHYSICAL PROPERTIES

The characteristics which enable us to distinguish one substance from another are known as PROPERTIES. Those properties involved in a transformation of one substance into another are known as CHEMICAL PROPERTIES. Thus, a chemical property is exhibited when wood burns, for the constituents of the wood combine with oxygen to form different substances. As iron corrodes, it shows a chemical property as the iron combines with oxygen to form the reddish-brown iron oxide which we know as iron rust.

In addition, to chemical properties, every substance also possesses definite PHYSICAL PROPERTIES. The physical properties of a substance are those which do not involve a change in composition of the material. Some familiar physical properties of matter are color, hardness, crystalline form, ductility, malleability, physical state, melting point, boiling point, density, electrical and thermal conductivity, and specific heat. Taste and odor are often classed as physical properties, but these sensations actually involve chemical changes. Changes in physical conditions such as temperature or pressure may modify the physical properties of a substance. For example, a substance in the gaseous state has certain physical properties such as density, specific heat, and thermal conductivity. By decreasing the temperature of the substance or compressing it to a smaller volume, it may be changed from the gaseous to the liquid state, in which condition it has an entirely different density, specific heat, and thermal conductivity.

A substance may be identified by its chemical and physical properties and by its composition because no two substances are alike in all respects.

1.7 CHEMICAL AND PHYSICAL CHANGES

When carbon, a black solid substance, burns in air, an invisible gas consisting of both carbon and oxygen (carbon dioxide) is formed. When milk sours, the sugar in the milk is converted into an acid, and the composition and the properties of the acid differ greatly from those of the sugar. Iron rust formed by the corrosion of iron metal contains oxygen as well as iron, and it is therefore a different substance with different properties than iron metal. All such changes are called CHEMICAL CHANGES. A chemical change always produces at least one substance entirely different in composition and properties from those that existed before the change occurred. In addition, all chemical changes are accompanied by either the production or absorption of some form of energy.

Changes that do not alter the composition of a substance are known as PHYSICAL CHANGES. The melting of ice, the freezing of water, the dissolving of sugar in water, and the heating of iron to redness, are all examples of physical changes. In each of these there is a change in properties but there is no alteration of the chemical composition of the substances involved. Water, whether in the solid, liquid, or gaseous state, retain the same chemical composition. Sugar is the same chemical substance in solution in water as it is in the solid state and can readily be recovered as crystals by evaporation of the water. Iron, an emitter of light when red hot, is still the same substance that reflects light when cold.

1.8 SUBSTANCES

A SUBSTANCE IS ANY VARIETY OF MATTER ALL SPECIMENS OF WHICH HAVE IDENTICAL PROPERTIES AND COMPOSITION. Pure water is an example of a substance. All samples of pure water, regardless of their source, have exactly the same composition, 2.01594 parts by weight of hydrogen to 15.9994 parts of oxygen, and are identical in melting point, boiling point, and all other properties. Pure iron, pure aluminum, pure carbon, pure sugar, pure oxygen, and pure carbon dioxide are representative substances. Two substances may be distinguished from each other by a study of their characteristic properties. Sugar and salt may be distinguished by taste, iron and gold by color, and silver and mercury by physical state.

1.9 MIXTURES

A MIXTURE IS COMPOSED OF TWO OR MORE SUBSTANCES EACH OF WHICH RETAINS ITS IDENTITY AND SPECIFIC PROPERTIES. This composition of a mixture can be varied continuously. Black gunpowder is a mixture of carbon, sulfur, and potassium nitrate; granite is a mixture of quartz, feldspar, and mica; plumber's solder is a mixture of crystals of tin and lead; a solution of sugar and water is a mixture; and air is a mixture of nitrogen, oxygen, carbon dioxide, water vapor, and other gases. Milk, butter, cement, flour, and gasoline are other examples of mixtures. In fact, most naturally occurring materials are mixtures.

Since each component of a mixture possesses its own set of characteristic properties, the various components can be separated by physical methods. The heterogeneous character of black gunpowder is readily detected by examining it under a microscope. Treating a sample of it with water causes the potassium nitrate to dissolve, leaving the sulfur and charcoal as solid particles; subsequent treatment of the residue with carbon disulfide causes the sulfur to dissolve, leaving only the carbon. The potassium nitrate and sulfur may be reclaimed as crystalline particles by evaporating their respective solutions to dryness. An intimate mixture of iron fillings and sulfur may be separated either by dissolving the sulfur in carbon disulfide, leaving the iron, or by removing the iron with a magnet, leaving the sulfur.

1.10 ELEMENTS

There are two classes of substances-elements and compounds. ELEMENTS ARE PURE SUBSTANCES, WHICH CANNOT BE DECOMPOSED BY A CHEMICAL CHANGE. One hundred and four elements are known at the present time. Eighty-eight elements have been found in nature and the other sixteen have been synthesized. Only about one-fourth of the elements ever occur in nature in the free state; the others are found only in combination with other elements. Eleven of the elements make up about 99 per cent of the earth's crust and atmosphere. Oxygen constitutes nearly one-half and silicon one-fourth of the total quantity of these elements in the atmosphere and the earth's crust.

1.11 COMPOUNDS

COMPOUNDS ARE SUBSTANCES WHICH CAN BE DECOMPOSED BY CHEMICAL CHANGES AND ARE COMPOSED OF TWO OR MORE DIFFERENT ELEMENTS. Elements in combination are different from elements in the free or uncombined state. The term "element" is used to designate an elemental substance whether free or present in compounds. For example, white crystalline sugar is a compound consisting of the element carbon, which is a black solid when free, and the two elements of hydrogen and oxygen, which are colorless gases when uncombined. Water, a compound, can be decomposed by an electric current into its two constituent elements hydrogen and oxygen. Table salt can be broken down by an electric current into sodium and chlorine.

Whereas there are only 104 known elements, there are hundreds of thousands of chemical compounds representing different combinations of these elements. Each of these compounds possesses definite chemical and physical properties by which chemists can distinguish it from all other compounds.

1.12 MOLECULES

Water has a definite composition and set of chemical and physical properties which enable us to recognize it as a distinct substance. One might ask the question, "To what extent could a drop of water be subdivided before the smallest particle would be obtained which would still have the chemical properties of water?" The limit to which the subdivision could be carried is a particle called a MOLECULE of water. Subdivision of a molecule of water would result in the formation of two substances, hydrogen and oxygen, each with a set of properties quite different from water and each other. THE SMALLEST PARTICLE OF AN ELEMENT OR COMPOUND THAT CAN HAVE A STABLE, INDEPENDENT EXISTENCE IS CALLED A MOLECULE. Molecules are too small to be seen even with very powerful optical microscopes. An idea of the minute size of molecules can be appreciated from the fact that if a drop of water were to be magnified to the size of the earth, its constituent molecules would appear to be about the size of baseballs. One hundred million molecules of water laid side by side make a row about one inch long.

1.13 ATOMS

WE MAY DEFINE AN ATOM AS THE SMALLEST PARTICLE OF AN ELEMENT WHICH CAN ENTER INTO A CHEMICAL COMBINATION. For example, one atom of carbon can combine with two atoms of oxygen to form one molecule of carbon dioxide. An atom of an element may or may not be capable of an independent existence. In some cases, an atom of an element and a molecule of it are identical. When the molecule of an element contains only one atom it is said to be monoatomic. Examples of elements which are composed of monoatomic molecules are helium, neon, and xenon. Hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine, and iodine consist of diatomic molecules (two atoms per molecule). Molecules of phosphorus and sulfur normally contain four and eight atoms, respectively. It follows, then, that the term "molecule" applies to small particles of either elements or compounds, whereas the term "atom" always applies to the smallest particle of an element.

The word "atom", from the Greek word atomos, means indivisible. The early Greek philosophers were the originators of the conception of atoms, and they taught that matter, being composed of atoms, is therefore finitely divisible. Although John Dalton, an English chemist and physicist, did not devise the atomic theory, it was he who revived the old Greek atomic hypothesis and put it on a quantitative basis.

For this reason Dalton is generally credited with being the father of the atomic theory (1808), the most important of all chemical theories. Dalton made no distinction between atoms and molecules as we do today, and he applied the name "atom" to particles of both elements and compounds.

1.14 THE SCIENTIFIC METHOD

The SCIENTIFIC METHOD furnishes a logical way of finding the answer to questions which can be subjected to inquiry and investigation. The first step in applying the scientific method to the solution of a problem involves the carrying out of carefully planned experiments to gain facts which give information about all phases of the problem. The second step consists of an attempt to formulate a simple generalization which will correlate a number of these facts. If this attempt is successful, the simple generalization becomes a LAW. Usually, how-ever, no general law can be formulated to correlate the facts, and instead a provisional conjecture, known as a HYPOTHESIS, is advanced to explain the data. A hypothesis is then tested by further experiments, if it is capable of explaining a large body of facts in a given field, it is dignified by the name THEORY. Theories serve as guides for further work by serving as the basis for predicting new information or the direction in which additional information must be sought. Finally, a theory must be established or modified in such a manner that it can be accepted as a general truth, which is often referred to as a LAW.

1.15 UNITS OF MEASUREMENT

Chemistry is a quantitative science which is concerned with the measurement of quantities of matter and energy. In any kind of quantitative work it is necessary to have a system of units of measurement, and it is of great advantage to have a convenient system. In scientific work the METRIC SYSTEM of weights and measures is used throughout the world. Calculations in the metric system are simple because the units are based upon the decimal system. However, in the ENGLISH SYSTEM the units are based on arbitrary standards of measurement and no simple relationship exists between them. Five elementary units of measurement are: MASS, LENGTH, VOLUME, TIME, and TEMPERATURE.

The following prefixes will be useful when using the metric system.

Table: Prefixes of the International System

Prefixes for SMALL quantities:

Number Prefix Sysmbol
10-1 deci d
10-2 centi c
10-3 milli m
10-6 micro u (mu)
10-9 nano n
10-12 pico p
10-15 femto f
10-18 atto a
10-21 zepto z
10-24 yocto y

Prefixes for LARGE quantities:

Number Prefix Symbol
101 deca da
102 hecto h
103 kilo k
106 mega M
109 giga G
1012 tera T
1015 peta P
1018 exa E
1021 yotta Y
1024 zetta Z

1.16 MASS AND WEIGHT

We know that MASS is the quantity of matter which a body contains and that the force required to give the body a given acceleration is a measure of its mass. The mass of a body of matter is an invariable quantity. On the other hand, the WEIGHT, of a body pertains to the force of attraction of the earth for the body and is dependent upon its distance from the earth's center. If this book were taken to the top of a mountain it would weigh less than it does at sea level. If it were to be taken far out into space, its weight would become negligible. Astronauts experience weightlessness while in outer space. Scientists have come to measure quantities of matter in terms of mass rather than weight because the mass of a body remains constant whereas the weight of a body is an "accident of its environment". It should be noted, however, that in common terminology the term "weight" is often used for the mass of a substance.

The gravitational attraction between two bodies of matter depends upon their masses and the distance separating them. The instrument used in science for determining the mass of an object is called a BALANCE. In the weighing process, the mass of an object is compared to a standard defined mass. The UNIT OF MASS in the metric system is the GRAM, which is equal to one one-thousandth of the mass of the standard kilogram. The STANDARD KILOGRAM is a cylinder of platinum-iridium alloy which is kept at the International Bureau of Weights and Measures at Sevres, France. The GRAM is very nearly EQUAL TO THE WEIGHT OF ONE CUBIC CENTIMETER OF WATER AT FOUR DEGREES CELSIUS, the temperature of its maximum density.

1 kilogram (Kg) = 1000 Grams (g) = 2.20462 Pounds (lb.)

1 Gram = 1000 Milligrams (mg)

453.6 Grams = 1 lb.

1.17 LENGTH

The second elementary measurement that is of interest to us is that of length. LENGTH is the distance covered by a line segment connecting two points. The standard for measuring length was for many years another piece of metal kept at Sevres, France, with two scratches on it. The distance between these two scratches was defined as 1 METER. In 1960, however, the standard meter was redefined in terms of the wave length of a particular color n the electromagnetic spectrum.

1 Meter (m) = 100 Centimeters (cm) = 1000 Millimeters (mm)

1 Meter = 39.37 Inches (in.)

2.54 cm = 1 in.

1.18 VOLUME

It is often times more convenient to measure the volume of a body of matter than to weigh it; this is particularly true of substances in the liquid or gaseous state. The VOLUME of a body of matter is the space that it occupies. The unit of volume is the LITER, which is the space occupied by one kilogram (1000 g) of water at four degrees celsius.

1 Liter (l) = 1000 Milliliters (ml)

1 cm3 = 1 c.c. = 1 ml

946 ml. = 1 Quart (qt.)

20 drops = approx. 1 ml

1.19 TIME

The fourth elementary measurement with which we are concerned is that of time. TIME is the interval between two occurrences. Our present standard of time, like our standard of length, is now based on an electronic transition in an atom. The older method, which is still adequate for most measurement, is more easily understood. This division of time is 1/86,400 of an average day, called 1 second (sec). The most common device for measuring time is, of course, a watch or clock. A more precise timepiece such as a chronometer or an atomic clock is often used.

The present standard of time is now defined as "the duration of 9,192,631,770 periods of the radiation corresponding to the transition between two hyperfine levels of the fundamental state of the atom Cesium 133".

1.20 TEMPERATURE

The word TEMPERATURE refers to the "hotness" or "coldness" of a body of matter. In temperature measurement some physical property of a substance which varies with temperature must be used. Practically all substances expand with an increase in temperature and contract when the temperature falls. It is this property which is the basis for the thermometer. The mercury or alcohol in our common glass thermometers rises when the temperature increases because its volume expands more than does the column of the glass container. The mercury thermometer has a smaller bore than the alcohol thermometer; hence it is usually more difficult to read. The bore of the mercury thermometer is made smaller because mercury expands about one-sixth as much as alcohol does, making the mercury thermometer inherently less accurate than an alcohol thermometer of the same bore.

1.21 STANDARD REFERENCE TEMPERATURES

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In order that we may agree on a set of temperature values, it is necessary to have fixed temperatures which are readily determined. Two FIXED TEMPERATURES which are commonly used are the freezing and boiling points of water. On the Celsius (sometimes called centigrade) scale the freezing point of water is taken as 0 and the boiling point as 100. The space between these two fixed points is divided into 100 equal intervals or degrees. On the Fahrenheit scale the freezing point of water is 32 and the boiling point is 212. The space between these two fixed points is divided into 180 equal parts or degrees. Thus a degree Fahrenheit is 100/180, or 5/9, of a degree Celsius. The relationships are shown by the following equations:

C = 5/9 (F - 32)

F = (9/5C) + 32

The readings below zero on either scale are treated as negative.

1.22 KELVIN OR ABSOLUTE TEMPERATURE

By experimentation it has been found that when 273 ml of gas at 0o is warmed to 1o, its volume increases by 1 ml to 274 ml; at 20o its volume increases to 293 ml; at 273o its volume increases to 546 ml (double that at 0o), and so on, provided that in each case the pressure remains constant. Note that the volume of the gas at 0o increases 1/273 of its volume for each increase of 1o on the Celsius scale. The volume of a gas decreases in the same proportion when the temperature falls. If the temperature of 273 ml of a gas could be lowered from 0o to -273o, then the gas should have no volume at -273o because its volume should decrease at the rate of 1/273 of its volume at 0o for each degree of fall in temperature. Before the temperature of -273o is reached all gases become liquids, to which this rate of change in volume does not apply. The temperature of -273o (or more exactly - 273.16oC) is called ABSOLUTE ZERO.

Sir William Thompson (Lord Kelvin), a British scientist, used the idea of absolute zero to develop a temperature scale which we call by his name today. Each degree on the Kelvin scale (oK), which is also known as the absolute temperature scale (oA), is equal to one degree Celsius. But since 0oK is equal to -273oC, there are no negative temperatures on the Kelvin, or absolute, scale.

To convert oC to oK, simply add 273o to the number of degrees on the Celsius scale.

oK = oC + 273

To convert oK to oC, subtract 273o from the number on the Kelvin scale.

oC = oK - 273

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1.23 THE MEASUREMENT OF HEAT

Chemical reactions are accompanied by either the evolution or the absorption of heat energy. The unit of measurement of heat is the calorie. The SMALL CALORIE is approximately equal to the quantity of heat which will raise the temperature of one gram of pure water one degree Celsius. The amount of heat necessary to raise the temperature of one gram of water one degree Celsius is not quite the same at all temperatures, but for our purposes it is sufficiently accurate to assume that it is the same. The LARGE CALORIE or kilocalorie (kcal.) is equal to 1000 small calories. The HEAT CAPACITY of a body of matter is the number of calories necessary to raise its temperature one degree Celsius. The heat capacity of one gram of water is one calorie, so the heat capacity of any amount of water is numerically equal to its weight in grams. The greater the mass of a substance, the greater its heat capacity.

The specific heat of a substance is a physical property that may be used in describing the substance. The SPECIFIC HEAT of a substance is the number of calories of heat required to raise the temperature of one gram of the substance one degree. Every substance has its own specific heat.

HEAT REQUIRED = MASS X SPECIFIC HEAT X TEMPERATURE CHANGE

1.24 DENSITY

One of the specific properties of a solid, liquid, or gas is density. DENSITY IS DEFINED AS MASS PER UNIT VOLUME. This may be expressed mathematically as:

DENSITY = MASS/VOLUME

D = M/V

Substances may be distinguished by measuring their densities because it is rare that any two substances have identical densities. The density of solids and liquids are expressed in grams/milliliter (g/ml) while gases are expressed in grams/liter (g/l).

1.25 SPECIFIC GRAVITY

The term SPECIFIC GRAVITY (sp. gr.) denotes the ratio of the mass of a substance to the mass of an equal volume of a reference substance. The reference substance for solids and liquids is usually water.

Sp. gr. = MASS OF SOLID OR LIQUID/MASS OF AN EQUAL VOL. OF WATER

Common reference substances used in specifying the specific gravity of a gas is usually air and hydrogen. When measured in the metric system of units, the DENSITY of any substance has practically the same numerical value as the SPECIFIC GRAVITY.


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Copyright May 1987 James R. Fromm (mailto:jfromm@3rd1000.com) Revised February 2000