Atoms, elements, and chemical reactions

French chemist Antoine Lavoisier began a series of experiments in 1770 that led to his discovery of oxygen. Lavoisier is shown here with his wife, who acted as his laboratory assistant.

Fire is one of the oldest civilizing influences on earth. Since civilization began, fire has been near at hand, inspiring human beings to search for ways to create and use new substances. Fire helped humans survive the Ice Age. The chemical energy released by the combustion (burning) of wood warmed Stone Age people, cooked their food, and lit up the cold, dark caves in which they sought shelter.

When prehistoric people began to settle down and raise crops, the drawbacks of their stone tools soon became apparent. Metal is much more versatile. It can be shaped into tools that work more efficiently. Metal tools are thinner and lighter. They take a sharp cutting edge and can be easily resharpened. And since copper is virtually the only metal found in usable form in nature, it became the first metal used by our prehistoric ancestors.

Copper is softer than stone, so these ancient metalworkers found that copper cutting tools had to be constantly resharpened. Then, in the fires of a primitive furnace fueled by burning charcoal, the early metallurgists discovered they could release copper from several blue and green mineral compounds. Later, they found that combining the hot, molten copper with another metal—tin—created an alloy—a harder, more useful substance called bronze. And, about 1000 B.C., ancient chemists learned howto use fire to free iron from its mineral compounds, creating tools and weapons far superior to any others of that time.

These ornamental horses’ heads date from the Bronze Age. They were made in Denmark in about 1000 B.C. Bronze was one of the first alloys made and used by humans. It was considerably harder than copper and other metals previously known.

Analyzing elemental burning

In the 1700’s, after several thousand years of trial and error, most of the metals used today were available. A host of other useful materials, including glass and ceramics, were also produced by chemical changes brought about by the heat of combustion—but how combustion caused these changes was poorly understood until 1770. During that year, French chemist Antoine Lavoisier (1743-1794) began a series of experiments designed to learn what happens to elemental substances when they burn.

By careful observation and analysis,

Lavoisier found that chemical substances like sulfur and phosphorus produce substances of greater weight when they burn. He theorized that these substances combine with a gaseous element in the air in the process of combustion. He named this element oxygene, which means “acid former,” because he observed that when nonmetallic substances such as sulfur, phosphorus, and carbon burn and combine with oxygen, the products of combustion (called oxides whether solid or gas) react with water to form acids. Lavoisier concluded that all acids must contain oxygen. [Hydrochloric acid, HCI, which consists of just hydrogen and chlorine, was discovered later and forced a revision of Lavoisier’s theory of acidity.]

Another newly discovered gas—one that reacts with oxygen to produce water—was given the name hydrogene, which means “water former,” and recognized as an element. Hydrogen had been discovered in 1766 by English chemist Henry Cavendish (1731-1810), who prepared the gas in his laboratory by reacting iron with sulfuric acid. The chemical formula for the reaction is:

Fe (iron) + H2S04 (sulfuric acid) —> H, T + FeS04 (iron sulfate)

When Cavendish mixed two parts of this new gas with one part oxygen and ignited it, the mixture exploded, producing droplets of water. He collected the water and found that it weighed exactly as much as the combined weight of the quantity of the two gases used in the reaction. Cavendish’s work prompted Lavoisier to conclude that water was a substance made from the chemical combination of oxygen and hydrogen.

Lavoisier believed that air is mainly a mixture of two elemental gases: oxygen—the gas essential to life and required for combustion, and a gas he named azote, which by itself is inert, quite unreactive. Today, azote is known as nitrogen, which means “niter former,” from the old name for potassium nitrate (KN03).

By the end of the 1700’s, due to the work of experimenters like Lavoisier, scientists had come to understand that air is not just one gas but a blend of gaseous elements. They also knew that when substances burn, they combine with one of the elements in this blend-oxygen. Discoveries like these were essential to the development of modern chemistry and served as the foundation for all later progress.

Applying math to chemical phenomena

In 1661, Robert Boyle (1627-1691), an Irish scientist, gave the world the first modern definition of element. He stated that an element was something primitive and simple, pure and unmixed, not made of anything else. His opinion differed from other scientists of the time, who thought that water was an element. Boyle believed that water was a compound, consisting of elements in combination, though he didn’t know what those elements were.

Boyle was among the first to apply mathematics to chemical phenomena. He formulated the law that describes how the volume of a gas is related to the gas’s pressure. When the pressure of a gas increases, the volume decreases proportionally and vice versa, provided that the temperature of the gas remains the same. Mathematically, Boyle’s Law is expressed as the following:

P (pressure) X V (volume) = k (a constant quantity) or V = k/P

Boyle also believed in the existence of atoms, which he called “corpuscles,” but they played no meaningful role in his concept of chemical change. It took the work of later scientists, like John Dalton (1766-1844), an English contemporary of Lavoisier, to put chemistry on a firm, atomic footing.

In 1803, Dalton became the first to clearly state that all matter consists of atoms, that an element consists of atoms with the same size and weight (except for isotopes), and that the atoms of one element are different from the atoms of another. Dalton’s atomic theory included two other important propositions: in chemical reactions, atoms combine or separate from each other; and when atoms combine to form a compound, they do so in their entirety, rather than in fractional parts. Therefore, elements always combine in the same proportions and in simple ratios such as 1:1,

1:2, 3:4. An example of these proportions and ratios, embodied in the Law of Definite Proportions, is water.

When water is created by the chemical changes resulting from the union of oxygen and hydrogen, 8 times as much oxygen as hydrogen (by weight) is consumed. This means that 1 gram of hydrogen will always combine with 8 grams of oxygen to yield exactly 9 grams of water, the simple ratio here being 1:8. This may seem obvious now, but as late as 1800, some scientists still believed that a chemical compound could have a variety of compositions. For example, they believed that water contained different percentages of hydrogen depending on its source.

Using his atomic theory, Dalton made the first calculations of the relative weight of atoms of different elements; in other words, how many times heavier an atom of one element is than an atom of another element. Fie chose hydrogen as a standard since it is the lightest element, assigning to it an atomic weight equal to 1. Dalton knew that a unit weight of hydrogen would combine with 8 unit weights of oxygen, so he assigned oxygen a relative atomic weight of 8. Flowever, he was unaware that hydrogen is made up of H2 molecules, not single FI atoms. Therefore, each volume of hydrogen gas contained double the number of atoms he supposed. As a result, his calculation for the atomic weight of oxygen was only half the actual value, and his equation for the reaction that produces water was:

H + O HO

Earth s atmosphere contains about 21 per cent (by volume) of oxygen at sea level, at the bottom of the troposphere. The amount of oxygen decreases rapidly with altitude. At orbital heights, there is virtually no oxygen at all. Astronauts have to carry their own oxygen supply in tanks strapped to their backs (far left). Within the stratosphere, ultraviolet radiation (left)irom the sun converts some oxygen into its al-lotrope ozone (below). Ozone acts as a filter that blocks most ultraviolet radiation.

This equation, although disputed by some scientists, persisted until 1860, when chemists became generally convinced of the diatomic nature of gases. When they recognized that the common gaseous elements occur as atomic twins or diatomic molecules, they wrote the equation that is known today for water reaction:

2H2 + 02 2H20

The scientist whose ideas made this change in theory possible was an Italian chemist named Amadeo Avogadro (1776-1856). In 1811, Avogadro proposed that the common elemental gases consist of atoms in combination-particles he called molecules— and that when two such gases react with each other, the molecules split and the atoms change partners, recombining to form a new compound. For example, when hydrogen reacts with chlorine, hydrogen molecules split, chlorine molecules split, and each atom of one combines with an atom of the other to form a molecule of the colorless gas hydrogen chloride, as follows:

H2 + Cl2 2FI Cl

French chemist Joseph Gay-Lussac (1778-1850) had shown that gases always react in simple ratios such as 1:1,1:2. One liter of chlorine, for example, always reacts with one liter of hydrogen to produce two liters of hydrogen chloride. Avogadro had the amazing insight to propose that each of these two equal volumes of gas contains the same number of molecules, despite the fact that the gases are different, the molecules are different, and their weights are different. Different gases contain the same number of molecules in one liter because molecules and atoms are so small, compared with the spaces that separate them, that one type of gaseous molecule does not occupy significantly more space than another type.

Combustion takes many forms, all of which are examples of oxidation. Burning is a comparatively slow reaction in chemical terms, although often extremely destructive.

Weighing and counting atoms

Avogadro’s deduction helped chemists to continue the work that Dalton had begun—determining the relative atomic weights (mass numbers) of the elements. The atomic weight of a gaseous element could be determined by ••• eighing it to see how many times heavier it • as than an equal volume of hydrogen. Since :ne number of molecules (or atoms) of the gas is the same as that in an equal volume of hydrogen (if the two gases are at the same temperature and pressure), the factor by which :ne weight of one exceeds that of the other equals the relative weight of an atom of the ;as. So, weighing a liter of hydrogen and a ter of oxygen, the one containing as many molecules as the other, chemists found that cxygen molecules weigh 16 times as much as hydrogen molecules, making the atomic ••• eight of oxygen 16, instead of the 8 that Dalton calculated.

When more accurate weights could be determined, scientists were able to classify all :ne known elements according to their properties and to list them in order of atomic ••• eight—from the lightest to the heaviest Russian chemist Dmitri Mendeleev (1834-1907) began doing this in 1869. When Mendeleev grouped elements of similar properties together in columns—the forerunner of today’s periodic table—he discovered the natural order that exists among the elements. He found :hat elements with different atomic weights could have quite similar characteristics. Vanadium, niobium, and tantalum, for example,”ave very different atomic weights, but they are close together in the periodic table in a column of “transition metals,” and they tend to occur together in nature.

The table above lists the chemical elements in alphabetical order and gives their symbols. The symbols are used in chemical formulas and in the periodic table, which appears in the next section of this book.

Avogadro’s hypothesis, as it came to be Miown, had another very useful result It encouraged scientists to try to calculate the number of atoms in a given volume of a substance. Scientists knew that there was a huge number of atoms and molecules in the measures of volume used in laboratories, so they decided to give a special name to a very large number of atomic or molecular particles (just as the term dozen is often used to denote the number 12). The unit that chemists adopted is called the mole, and it represents 6.02213 X 1023 (602,213 billion billion), a huge number of atoms or molecules. The number itself is called the Avogadro number, in honor of the Italian chemist.

How was this number determined? By weighing a mass in grams of an element numerically equal to its atomic weight or mass number, scientists learned the molar mass of the element. Then by a variety of experimental methods they calculated that the mass contained approximately 6.02213 X 1023 atoms. The mass of a mole of atoms of an element— the molar mass measured in grams—is thus by definition numerically equal to the element’s atomic weight, or mass number. For example, the atomic weight or mass number of carbon is 12, so a mole of carbon atoms, consisting of 6.02213 X 1023 atoms, has a mass of 12 grams. The molar mass of a compound, such as water, is simply the sum of the atomic weights or mass numbers of the components of the molecule: for water, it is 16 + 1 + 1 = 18 grams.

Knowledge grew as the discoveries of one generation of chemists laid the groundwork for the discoveries of the next.

Niobium, a transition metal, is used in nuclear reactors as a coating for the radioactive fuel rods, for which purpose it is sometimes alloyed with zirconium. Niobium and zirconium allow neutrons to pass through them easily and thus make it possible for the fission reaction to proceed at controllable self-sustaining rates.