CHEM101: General Chemistry I

The idea of the atom – at one time a theory, but now directly observable – is the basic concept that unites all aspects of chemistry, so this is where we begin. This lesson introduces you to these building-blocks of matter, and explains how they are characterized.

Scientists at the IBM Almaden laboratory in California teamed up with artists to create these striking images.

Carbon monoxide man was made by selectively removing CO molecules from a platinum surface.

The Kanji characters for atom consist of iron atoms deposited on a copper surface.

1. The Idea of the Atom

Atoms and Elements: What is the Difference?

The parallel concepts of the element and the atom constitute the very foundations of chemical science.

An element is an actual physical substance that cannot be broken down into a simpler form, and is capable of an independent existence as observable matter. As such, the concept of the element is a macroscopic one that relates to the world that we can observe with our senses.

The atom is the microscopic realization of this concept; that is, it is the actual physical particle that is unique to each chemical element. Their very small size has long prevented atoms from being observable by direct means, so their existence was not universally accepted until the late 19th century. The fact that we still hear mention of the atomic theory of matter should not imply that there is now any doubt about the existence of atoms; few theories in the history of science have been as thoroughly validated and are as well understood.

Although the word atom usually refers to a specific kind of particle (an atom of magnesium, for example), our everyday use of element tends to be more general, referring not only to a substance composed of a particular type of atom (bromine is one of the few elements that are liquids at room temperature), but also to atoms in a collective sense (magnesium is one of the elements having two electrons in its outer shell).

Atoms are Not New!

Image source: Great Indians

The underlying concept of atoms as the basic building blocks of matter has been around for a long time.

As early as 600 BCE, the Gujarati (Indian) philosopher Acharya Kanad wrote that "Every object of creation is made of atoms which in turn connect with each other to form molecules".

A couple of centuries later in 460 BCE, the Greek philosopher Democritus reasoned that if you keep breaking a piece of matter into smaller and smaller fragments, there will be some point at which the pieces cannot be made any smaller. He called these basic matter particles – in other words, atoms. But this was just philosophy; it would not become science until 1800 when John Dalton showed how the atomic concept followed naturally from the results of quantitative experiments based on weight measurements.

2. The Chemical Elements

The Philosopher's Elements

The four elements of western alchemy: This figure shows how the four elements were imagined to combine in various pairs to produce the qualities of hot, cold, wetness and dryness.

The concept of the element is an ancient one which developed in many different civilizations in an attempt to rationalize the variety of the world and to understand the nature of change, such as that which occurs when a piece of wood rots, or is burnt to produce charcoal or ash. Most well known to us are the four elements – earth, air, fire, and water – that were popularized by Greek philosophers (principally Empedocoles and Aristotle) in the period 500-400 BCE.

To these, Vedic (Hindu) philosophers of India added space, while the ancient Chinese concept of Wu Xing regarded earth, metal, wood, fire and water as fundamental. These basic elements were not generally considered to exist as the actual materials we know as earth, water, etc., but rather represented the principles or essences that these elements contributed to the various kinds of matter we encounter in the world.

The Chemist's Elements

Antoine Lavoisier (1743-1794)

Eventually, practical experience (largely connected with the extraction of metals from ores) and the beginnings of scientific experimentation in the 18th Century led to our modern concept of the chemical element.

Simplest, in the context of experimentation at the time, was defined in terms of weight; cinnabar (mercuric sulfide) can be broken down into two substances, mercury and sulfur, which themselves cannot be reduced to any lighter forms.

The first textbook of chemistry, Traitè Èlèmentaire de Chemie, published by Antoine Lavoisier, the so-called father of chemistry, in 1789, contained the table of elements shown here.

Although Lavoisier got many of these right, he did manage to include a few things that do not quite fit into our modern idea of what constitutes a chemical element. There are two such mistakes in the top section of the table that you should be able to identify even if your French is less than tip-top – can you find them?

Lavoisier's other mis-assignment of the elements in the bottom section was not really his fault. Chalk, magnesia, barytes, alumina and silica are highly stable oxygen-containing compounds; the high temperatures required to break them down could not be achieved in Lavoisier's time. (Magnesia, after all, is what fire brick is made of!) The proper classification of these substances was delayed until further experimentation revealed their true nature.

Ten of the chemical elements have been known since ancient times. Five more were discovered through the 17th Century.

Some Frequently-Asked Questions about Elements

How Many Elements are There?

Ninety-two elements have been found in nature. Around 25 more have been made artificially, but all of these decay into lighter elements, with some of them disappearing in minutes or even seconds.

Where do the Elements Come From?

The processes by which elements (or more properly, their nuclei) are formed is known as nucleosynthesis. The very first nuclei of the lightest elements (hydrogen and helium) formed about 20 minutes after the big bang 13.8 billion years ago.

The next 23 elements (up through iron) are continuously being formed mostly by nuclear fusion processes within stars, in which lighter nuclei combine into successively heavier elements. Elements heavier than iron cannot be formed in this way, and are produced only during the catastrophic collapse of massive stars (supernovae explosions).

How do the Elements Vary in Abundance?

Quite markedly, and very differently in different bodies in the cosmos. Most of the atoms in the universe still consist of hydrogen, with helium being a distant second. On Earth, oxygen, silicon, and aluminum are most abundant. These profiles serve as useful guides for constructing models for the formation of the earth and other planetary bodies.

Elemental abundances in the lithosphere (Earth's crust) and in the universe.

Note that the vertical axis is logarithmic, which has the effect of greatly reducing the visual impression of the differences between the various elements.

How Did the Elements Get Their Names?

This is too big a subject to cover here in detail, especially since most elements have different names in different languages.

• By far the most interesting resource is the Elementymology Multidictionary which covers the history, discovery and naming of each element and a list of the names of each element in 97 languages.
  
  
• The naming of some of the more recently-discovered (artificially-created) elements has been a matter of some controversy.

How Did the Element Symbols Originate?

In 1814, the Swedish chemist Jöns Jacob Berzelius devised the one- and two-letter symbols for the elements known at that time. Prior to that time, graphic alchemical symbols were used, which were later modified and popularized by John Dalton. Fortunately for English speakers, the symbols of most of the elements serve as mnemonics for their names, but this is not true for the seven metals known from antiquity, whose symbols Berzelius based on their Latin names. The other exception is tungsten (a name derived from Swedish), whose symbol W reflects the German name which is more widely used.

How are the Elements Organized?

Two general organizing principles developed in the 19th Century: one was based on the increasing relative weights (atomic weights) of the elements, yielding a list that begins this way:

H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca...

The other principle took note of the similarities of the properties of the elements, organizing them into groups with similar properties (Döbereiner's triads, 1829). It was later noted that groups of elements (Chancourtois's telluric helix, Newlands octaves, 1864, and Meyer, 1869) having changing properties tended to repeat themselves within the atomic weight sequence, giving rise to the idea of periodic sequences of properties. These concepts were finally integrated into the periodic table published by Mendeleev in 1869, which evolved into the various forms of the periodic table in use today.

3. Atoms Become Real

The word atom comes from the Greek a-tomos, meaning un-cut-able

Throughout most of history the idea that matter is composed of minute particles had languished as a philosophical abstraction known as atomism, and no clear relation between these atoms and the chemical elements had been established. This began to change in the early 1800's when the development of balances that permitted reasonably precise measurements of the weight changes associated with chemical reactions ushered in a new and fruitful era of experimental chemistry. This resulted in the recognition of several laws of chemical change that laid the groundwork for the atomic theory of matter: conservation of mass, definite proportions, and multiple proportions.

Laws of Chemical Change

Conservation of Mass-Energy in Chemistry

Mass conservation is usually considered the most fundamental of law of nature. It is also a good example of a law that had to be modified; it was known simply as Conservation of Mass until Einstein showed that energy and mass are interchangeable. However, the older term is perfectly acceptable within the field of ordinary chemistry in which energy changes are too small to have a measurable effect on mass relations.

Mass conservation had special significance in understanding chemical changes involving gases, which were for some time not always regarded as real matter at all. (Owing to their very small densities, carrying out actual weight measurements on gases is quite difficult to do, and was far beyond the capabilities of the early experimenters.) Thus when magnesium metal is burned in air, the weight of the solid product always exceeds that of the original metal, implying that the process is one in which the metal combines with what might have been thought to be a weightless component of the air, which we now know to be oxygen.

More importantly, as we will see later, this experimental result tells us something very important about the mass of the oxygen atom relative to that of the magnesium atom.

Law of Definite Proportions

This law is also known as the law of constant composition. It states that the proportion by weight of the element present in any pure substance is always the same. This enables us to generalize the relationship we illustrated above.

Law of Multiple Proportions

Many combinations of elements can react to form more than one compound. In such cases, this law states that the weights of one element that combine with a fixed weight of another of these elements are integer multiples of one another. It is easy to say this, but please take the time to make sure that you understand how it works!

Consider, for example, the five compounds of nitrogen and oxygen depicted in the chart shown here.

• Lineshows the ratio of the relative weights of the two elements in each compound. These ratios were calculated by simply taking the molar mass of each element, and multiplying by the number of atoms of that element per mole of the compound. Thus for NO₂, we have (1 × 14) : (2 × 16) = 14:32. (These numbers were not known in the early days of chemistry because atomic weights (i.e., molar masses) of most elements were not reliably known.)
  
  
• The numbers in Lineare just the mass ratios of O:N, found by dividing the corresponding ratios in line 1. But someone who depends solely on experiment would work these out by finding the mass of O that combines with unit mass (1 g) of nitrogen.
  
  
• Line Line is obtained by dividing the figures the previous line by the smallest O:N ratio in the line above, which is the one for N₂O. Note that just as the law of multiple proportions says, the weight of oxygen that combines with unit weight of nitrogen work out to small integers.
  
  
• Of course we just as easily could have illustrated the law by considering the mass of nitrogen that combines with one gram of oxygen; it works both ways!

In order to make sure you understand this problem, note carefully that:

• The masses of hydrogen that react with 1.00 g of nitrogen to form the three compounds:  .1428 g., .0714 g, and .2857 g  are sometimes referred to as combining weights.
  
  
• These combining weights would be different if the mass of nitrogen that was used had a value different from 1.00 g.
  
  
• In an actual experiment, the mass of nitrogen would almost certainly be different from 1.00 g, but the combining weights would differ in exact proportion. This means that if n grams of nitrogen were used, the mass of hydrogen that reacts to form urea, for example, would be .1428 × n.

The most systematic way of solving this problem is to construct a table whose three leftmost columns summarize the information given in the problem:

Next, identify the smallest weight ratio (.0714 here), and, in the rightmost column, divide each of the values in column three by this smallest ratio. What you get are the small whole numbers that illustrate the law of simple multiple proportions.

Note that, in some cases, the values in column four will not all initially work out to integer numbers. Suppose, for example, that the bottom row refers to some compound other than ammonium chloride, and that the numbers in column four (reading from top to bottom), work out to {2 - 1- 1.33}. If you recognize that the last number corresponds to 4/3, then multiplying each value by 3 yields the whole numbers {6 3 4} – not quite as small, but at least whole as the law requires.

How Dalton's Interpretation of the Laws of Chemical Change Established the Atomic Theory

The idea that matter is composed of tiny atoms of some kind had been around for at least 2,000 years. Dalton's accomplishment was to identify atoms with actual chemical elements.

If Nobel prizes had existed in the early 1800's, John Dalton (1766-1844), the English schoolteacher, meteorologist and chemist, would certainly have won one for showing how the experimental information available at that time, as embodied in the laws of chemical change that we have just described, are fully consistent with the hypothesis that atoms are the smallest units of chemical identity. These points of Dalton's atomic theory provided satisfactory explanations of all the laws of chemical change noted above:

Explanation of the Law of Conservation of Mass

This is really a consequence of conservation of atoms which are presumed to be indestructible by chemical means. In chemical reactions, the atoms are simply rearranged, but never destroyed.

Explanation of the Law of Constant Composition

If compounds are made up of definite numbers of atoms, each of which has its own characteristic mass, then the relative mass of each element in a compound must always be the same. Thus the elements must always be present in a pure sample of a compound in the same proportions by mass.

Explanation of the Law of Multiple Proportions

A given set of elements can usually form two or more compounds in which the numbers of atoms of some of the elements are different. Because these numbers must be integers (you cannot have half an atom!), the mass of one element combined with a fixed mass of any other elements in any two such compounds can differ only by integer numbers. Thus, for the series of nitrogen-hydrogen compounds cited in the problem example above, we have the following relations:

Seeing is Believing: the Atomic Force Microscope

These wooden balls have evolved into computer-generated images derived from the atomic force microscope (AFM), an exquisitely sensitive electromechanical device in which the distance between the tip of a submicroscopic wire probe and the surface directly below it is recorded as the probe moves along a surface to which atoms are adsorbed.