It's human nature to organize things. Cooks painstakingly organize their spices into various groupings, whether alphabetically or according to how often they're used. Kids dump out their piggy banks and sort their riches into piles of pennies, nickels, dimes and quarters. Even the items in a grocery store are grouped a certain way. Head down the international aisle, and you'll find packages of Chinese egg noodles sitting next to boxes of taco shells.
Chemists, as it turns out, are organizational junkies, too. They look for similar physical and chemical properties among the elements, the basic forms of matter, and then try to fit them into similar groups.
Scientists began attempting to organize the elements in the late 1800s when they knew of about 60. Their efforts, however, were premature since they were missing a key piece of information: the structure of the atom. While initial efforts failed, one attempt by a Russian chemist named Dmitry Mendeleyev showed much promise. Although Mendeleyev wasn't 100 percent correct, his approach laid the groundwork for what is now the modern periodic table of the elements.
Today, the periodic table organizes 112 named elements and acknowledges several more unnamed ones. It has become one of the most useful tools in chemistry, not only for students, but for working chemists as well. It classifies the elements according to their atomic number (more on that soon), tells us about the nuclear composition of any given element, describes how electrons are arranged around a given element and allows us to predict how one element will react with another.
So, exactly what is this feat of organization? Keep reading as we examine the history, organization and uses of this most handy chemical tool.
Getting Organized: Origins of the Periodic Table
In 1829, a German chemist by the name of J. W. Dobereiner noticed that certain groups of three elements had similar properties. He called these groups triads and published a system of classification based on them. For example, chlorine, bromine and iodine formed a triad, based on the fact that the atomic weight of bromine (79.904) was close to the average of the atomic weights of chlorine (35.453) and iodine (126.904). Unfortunately for Dobereiner and his scientific legacy, not all of the elements could be grouped into triads, so his efforts failed. Another classification system unsuccessfully attempted to group the elements into octaves, like musical notes.
In 1869, Russian chemist Dmitry Mendeleyev published the first periodic table of elements, writing the chemical properties and masses of each element on cards. He arranged the cards according to increasing atomic mass and found that elements of similar properties appeared at regular intervals. But he took some liberties with his table. In some cases, he violated his order of increasing atomic masses to keep elements with similar properties together. For example, he placed tellurium (atomic weight 128) before iodine (atomic weight 127), so that iodine could be grouped with chlorine, bromine and fluorine, all of which have properties similar to iodine. He also reasoned that if elements had to be reversed to preserve the periodic pattern, then the atomic mass values must be wrong. Lastly, he left gaps in his table for elements that he reasoned should exist, but hadn't been discovered.
Mendeleyev's periodic table predicted three elements of atomic weights 45, 68 and 70. He was proven right when these elements were later discovered and identified as scandium, gallium and germanium, respectively. The atomic weights listed in modern periodic tables are slightly different than those in Mendeleyev's time because methods for measuring atomic weights were improved during the 20th century. These discoveries demonstrated the usefulness of Mendeleyev's approach, even if it wasn't without problems. Explanations would have to wait until the early 20th century, when the structure of the atom began to be revealed.
In 1911, English chemist Henry Moseley studied the frequencies of X-rays given off by various elements when high-energy electrons bombarded each. The X-rays each element emitted had a unique frequency that increased with increasing atomic mass. Moseley arranged the elements in order of increasing frequency and assigned each one a number called the atomic number (Z). He realized that the atomic number was equal to the number of protons or electrons. When the elements were arranged by increasing atomic number, the periodic pattern was observed without having to switch some elements (as Mendeleyev did), and "holes" in the periodic table led to the discovery of new elements. Moseley's discovery was summarized as the periodic law: When elements are arranged in order of increasing atomic number, there's a periodic pattern in their chemical and physical properties. That law led to the modern periodic table.
Building the Periodic Table Block by Block
Each block of the periodic table houses an element, along with a few standard facts about that element:
- Atomic number: integer equal to the number of protons or electrons in the element. Gold's atomic number is 79.
- Element symbol: one or two letters. In the case of two letters, the first one is always capitalized. Hydrogen's symbol is just H, while helium's is He. Symbols can be tricky because some are based on the first letter(s) of the element's common name, as hydrogen's is, while other symbols are based on the Latin names of the element, such as Au for gold (or aurum in Latin).
- Element name
- Atomic weight: usually a decimal value, such as 196.966 569(4) for gold
Some periodic tables include the electron configuration (arrangement of electrons) in a corner of the block or below the name of the element. In addition, some periodic tables thoughtfully include colored symbols to indicate whether the element is a solid, liquid or gas at standard temperature (25 degrees C or 77 degrees F) and colored backgrounds to indicate the type of element (alkali metals, alkaline earth metals, nonmetals, noble gases and so on).
Within the table, the elements are arranged by increasing atomic number, as you'll recall. The elements stretch across seven rows. Each row is called a period and indicates the energy levels or shells occupied by the electrons around the nucleus of that element (see How Atoms Work). For example, the first energy level can only hold two electrons max, so hydrogen and helium occupy period 1. In period 2, the second energy level begins to fill. The pattern continues. The elements in period 7 have enough electrons to start filling the seventh energy level. No known element yet has eight energy levels.
Each energy level above the first one has sublevels or orbitals. The orbitals are s (sharp), p (principal), d (diffuse) and f (fundamental). But electrons don't fill directly in the order of s then p then d then f. That would be too easy. There's some overlap between the orbitals of one energy level and those of the one below it. For example, electrons in the fourth energy level fill in this order: 4s then 3d then 4p. (If you can't quite picture it, the American Chemical Society has a periodic table that allows you to see how the various electron configurations work here.)
As the atomic number increases and one energy level fills, a new period begins. If you placed all of the elements in order of increasing atomic number, the periodic table would span more than one tidy sheet of standard paper. That's why chemist Glenn Seaborg suggested pulling out the lanthanoids and actinoids and placing them below the table to make it more compact.
The electrons of the outermost energy levels are the restless ones that participate in chemical reactions. So as each new period begins, there are elements that have similar chemical properties -- those with one outer electron, those with two, three and so on. Mendeleyev couldn't have predicted this periodic nature because he didn't know about atomic structure. But what about the columns?
The Columns Supporting the Periodic Table
The columns that comprise the periodic table are called groups -- 18 in total. Groups indicate elements with similar chemical and physical properties. About 80 percent of the elements are metals (shiny elements that conduct heat and electricity well), and 15 percent of the elements are nonmetals (poor conductors of heat and electricity). The remaining elements are metalloids, which share properties of both metals and nonmetals. Let's look at some of these element cliques and remember, sometimes group members are spread around the table, not necessarily in one neat column. For example, hydrogen looks like it should belong to group 1, the alkali metals, but it actually prefers the company of nonmetals.
Alkali metals (group 1 or IA) such as lithium, sodium and potassium, are highly reactive and aren't usually found freely in nature. They get their name from their chemical reactions with water in which they produce highly alkaline substances such as sodium hydroxide or lye. They have one valence electron (or outermost electron that's farthest from the nucleus), which they give up in chemical reactions. Sodium gas fills streetlights, while sodium liquid is used to transfer heat in certain types of nuclear reactors.
Alkaline earth metals (group 2 or IIA) include magnesium, calcium and barium among others. These elements have two valence electrons, which they yield in chemical reactions. Although they're less reactive than alkali metals, they're not usually found alone in nature. For example, calcium combines with carbon to make calcium carbonate, which makes up limestone, marble and seashells. Teeth and bones are also made of calcium compounds. Beryllium contributes to the bling found in the gemstones aquamarine and emerald.
Lathanoids and actinoids (group 3 or IIIB) include shiny metals (lanthanide series or rare earth elements) and radioactive elements (actinide series). Lanthanoids are abundant in the Earth's crust, but difficult to separate from their compounds. All of the actinoids are radioactive, but only actinium, thorium, protractinium and uranium are found naturally. The other actinoids are made in nuclear reactors and particle accelerators.
Transition metals (groups 4-12 or IB, IIB and IVB-VIIIB) are all shiny metals that are found naturally, but are less reactive than groups 1 and 2. Electrons of the outermost s orbital and the inner d orbital can participate in chemical reactions. They include elements that we usually think of as metals, like iron, nickel, chromium and precious metals such as gold, copper, silver and platinum.
Metals are located mostly in group 13 (IIIA) and some in groups 14-16 (IVA - VIA). Metals include aluminum, tin, lead and bismuth. Metals are harder and denser than those in groups 1 and 2, but softer and less dense than the transition metals. Most of them are found as compounds in nature, but can exist freely once refined, as aluminum does.
Noble gases (group 18 or VIIIA) include helium, neon, argon, krypton, xenon and radon. Helium, of course, fills balloons and blimps. Neon, argon and xenon are used in lights. Radon is a product of radioactive decay from the Earth and comes up through the soil into your home. The noble gases are also called inert gases because they don't react chemically with other elements. Why not? The orbitals of their highest energy level are filled with electrons. Thus sated, they tend not to take or share their valence electrons with other elements.
You're not quite done yet. Metalloids and nonmetals round out the groups. Nonmetals can form compounds by sharing valence electrons with each other or swiping them from metals. One group of nonmetals (17 or VIIA) are highly reactive and called halogens (fluorine, chlorine, bromine, iodine and astatine).
How can all of this information help you detect some trends among Earth's elements?
Trends in the Periodic Table
It's handy to know about what group a particular element resides in and what its atomic structure is like, but that's not all the periodic table has to tell you. If you're looking at it, you're casually taking in work that scientists have spent lifetimes struggling with. And if you look at the table as a whole, some big trends start to emerge that tell us how one element will react with another.
Before we can see these trends, a quick chemistry recap might be good. First, metals react with nonmetals to form ionic compounds. The nonmetal atom takes one or more valence electrons from the metal atom. When an atom gains or loses a valence electron, it forms an ion. An ion with more protons than electrons is positively charged and called a cation (comes from the metal). An ion with more electrons than protons is negatively charged and is called an anion (comes from the nonmetal). In the end, both ions have a full outer energy level. Second, nonmetals tend to share electrons so that both atoms have full outer energy levels; they form covalent compounds. But how do you know which element will react with which to produce an ionic or a covalent compound? That depends on a few factors:
- Ionization energy: the amount of energy it takes to strip away the first valence electron
- Electronegativity: a measure of how tightly an atom holds onto its valence electrons
- Nuclear charge: the attractive force between the positive protons in the nucleus and the negative electrons in the energy levels. The more protons, the greater the nuclear charge.
- Shielding: inner electrons tend to shield the outer electrons from the attractive force of the nucleus. The more energy levels between the valence electrons and the nucleus, the more shielding.
Let's see how these factors can help predict what type of chemical reactions any two elements will make.
If you look at the periodic table, ionization energy tends to decrease as you move down a column and increase as you move across a period from left to right. When you compare elements in groups 1 and 2 (on the left) with those in 16 and 17 (on the right), you'll find that the elements in the first groups have lower ionization energies, won't hold on to their valence electrons as tightly and will tend to form cations. So, elements in groups 1 and 2 will tend to form ionic compounds.
Like ionization energy, electronegativity decreases as you go down a column and increases as you go across a period from left to right. So, fluorine is more likely to take electrons from another element than lithium. The difference in electronegativity between two elements will determine whether they exchange electrons (ionic compounds) or share electrons (covalent compounds). You can use trends in ionization energy and electronegativity to predict whether two elements will form ionic or covalent compounds.
Lastly, the nuclear charge increases as you go across and down the table, while the shielding stays constant across the periods, but increases as you go down the columns. These tendencies tell you about atom size. Atoms and ions get bigger as you go down the columns because the shielding effect outweighs the effects of the nuclear charge, so the attraction between the nucleus and electrons is weaker and the atom expands in size. In contrast, atoms get smaller as you go across the periods because the nuclear charge effect outweighs the shielding effect, so the attraction between the nucleus and the electron is greater and the atom shrinks in size.
It's hard to believe that one measly sheet of paper can contain that much information.
IUPAC: The Gatekeeper of Elements
The International Union of Pure and Applied Chemistry (IUPAC) oversees the periodic table of elements, which, as of November 2011, consisted of 112 officially named elements, like seaborgium and regular old potassium.
An official element is one that was claimed to have been discovered, the discovery was verified and the element was named. An unofficial element is one that was claimed to have been discovered, but the claim has not been verified, so the element hasn't been named. One of the more recent elements to claim fame on the periodic table was roentgenium, which was unearthed in December 1994 and named after Wilhelm Roentgen, the scientist behind X-rays.
The last elements to be discovered had atomic numbers of 112, 114, 116 and 118. They are unofficially called ununbium (Uub), ununquadium (Uuq), ununhexium (Uuh) and ununoctium (Uuo), respectively -- Greek for the atomic numbers of these elements. There are spots in the periodic table for elements with atomic numbers 115 and 117, but these elements haven't been discovered, much like Mendeleyev left gaps in his table for elements that hadn't turned up yet. Of course, nothing in science is static, so it's always good to check with IUPAC if you're unsure about whether an element is official or not.
Ununquadium doesn't exactly roll off the tongue, so how does an element get a new title and achieve official status? And are there any naming restrictions? Is christening an element after a beloved pet strictly frowned upon, but after a hometown or lab location accepted?
Remember that the new elements are all radioactive ones that are made in particle accelerators and have short lives before they decay into another element. In addition, any new element discovered must have a lifetime greater than 10-14 seconds. Two difficulties exist in confirming these new elements: First, they're not produced in large quantities, and second, they don't last very long. That means it's a long, difficult road to verify the claim that a new element has been discovered. But the procedure for naming an element is as follows:
- The claim that a new element has been found must be published in the scientific literature.
- IUPAC analyzes the claim as to who discovered it (often competing labs claim discoveries of new elements), whether the experiments were valid and whether it meets the criteria for a new element. IUPAC publishes its analysis in its official journal Pure and Applied Chemistry in which it establishes who discovered the element and how it was made.
- The element gets a provisional Greek name and square in the periodic table.
- IUPAC invites the credited discoverers to submit a name and symbol for the new element drawn from a mythological concept, mineral, country or place, property or scientist.
- The proposal is publicly reviewed, usually by neutral scientists.
- IUPAC makes the final decision.
- IUPAC publishes the name in Pure and Applied Chemistry and adds it to the periodic table.
Many people have made different representations of the periodic table, such as spiral forms, 3-D forms, even a humorous "Periodic Table of the Elephants" that features a cartoon elephant with a slightly different take on the elements. For example, the helium block features an elephant balloon filled with helium and the beryllium block features a cold elephant -- Brrr-illium. Get it? Still, none of these twists on the periodic table have yet proved as useful as the standard form that you see in any chemistry textbook today.
The rarity of gold is just one reason why people value the precious metal. Learn why gold inspires men to wage war and how we find and refine gold.
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More Great Links
- Brown, T.L. et al. "Chemistry the Central Science," eighth edition. Prentice-Hall. 2002.
- Chemistry Daily, History of the Periodic Table. Jan. 4, 2007.http://www.chemistrydaily.com/chemistry/History_of_the_periodic_table
- IUPAC. "Naming of New Elements." Pure Appl. Chem. Vol. 74, No. 5, pp. 787-791 (2002).http://old.iupac.org/publications/pac/2002/7405/7405x0787.html
- IUPAC. "Criteria that must be satisfied for the discovery of the new chemical element to be recognized." Pure Appl. Chem., Vol. 63, No. 6, pp. 879-886, 1991. http://old.iupac.org/reports/1991/6306wapstra/index.html
- Kaesz, H. "The synthesis and naming of elements 110 and beyond." Chem Int vol. 24 No. 2 Mar 2002.http://old.iupac.org/publications/ci/2002/2402/elements110.html
- Tzimopoulos, N.D. et al. "Modern Chemistry," third edition. Holt Rinehart and Winston. 1990.
- Wilbraham, A.C. et al. "Chemistry." Prentice-Hall. 2008.