The Periodic Table
Click here for Periodic Table Terminology
The periodic table is a classification system for the elements. It helps give an order to a seemingly unrelated bunch of elements. Scientists had discovered, through experience, that every so many elements certain properties recurred. This "periodicity" in getting to a similarly behaving element gave the periodic table its name.
Press here for a brief overview of the history of this fundamental tool of chemical understanding.
In the periodic table, the elements are positioned in sequence, right to left, top to bottom, by their proton count. The number of protons is called the 'Atomic Number'. The top left element, hydrogen, is atomic number 1, and continues through to Ununbium, number 112. Doesn't seem so special. Yet a deeper look reveals some astonishing things. The first two elements right to left horizontally, are in what's called the first period. They are hydrogen (H) and helium (He). The next eight are in the second period, Li, Be, B, C, N, O, F, Ne. Each horizontal row is numbered in sequence, top to bottom, 1 through 7. Across a period, left to right, there is a gradual change in the properties of the elements.
The second division of the periodic table is into columns. Columns are the vertical groups of elements. There are two different numbering systems, an older, and a newer system. The older system is more complicated to remember, but has usefulness since it predicts certain properties by its number. The newer system is simple in that the columns are simply numbered 1 to 18, from left to right. Column 1 has Li, Na, K, Rb, Cs, Fr. The elements in a given column have many similar properties even though their atomic number is increasing substantially from one to the next. The reason for this was determined long after the periodic table was composed. It is astonishing that the periodic table was put together so correctly and remains structurally the same even though it was put together in its final form before the discovery of valence shells or Quantum Mechanics.
It was remarkable too in that it predicted the properties of elements not yet discovered.
Also, by knowing what column an element occupied, many things could be known about it.
Things like what element it might bond with, whether it would conduct electrical impulses etc.
Some light was shed on the reasons for the periodic table's predictive powers with the discovery of the tendency of atoms to move towards having eight electrons in their outer or valence shells.
This is called the octet rule or rule of eight.
It says that atoms tend to lose or gain electrons when bonding to give eight electrons in their valence shells.
The American chemist Gilbert Newton Lewis in 1916 put forth explanations as to why non-metal atoms might bond and also devised electron dot symbols to show the bonding that occurs.
This set the stage for molecular modeling and the understanding of why certain bonds form and how strong they are. It served the practical purpose of allowing useful bonding diagrams to be drawn up. Yet a true understanding of the nature of element attraction was not achieved until the development of Quantum Mechanics in the late 1920"s.
Quantum Mechanics was able to explain the the reasons behind the number of elements in each period of the periodic table and even more amazingly, for all of us wishing to understand, explained why molecules bond as they do, particularly those very very long chains of elements that are the building blocks of life.
The reason that carbon is the cornerstone of all living things can be found in the study of covalent bonding.
With these concepts and principles in our back pocket, we can now see why elements are 'reactive or non-reactive', 'electron donors or electron accepters', 'inclined to form simple molecules or complex molecules' and
all kinds or other stuff!
This is all about why nitroglycerine explodes and helium doesn't.
Why the iron in hemoglobin transports oxygen readily to the cells and just as happily, drops it off.
It also tells us why the periodic table is divided as it is.
(insert pic of periodic table with orbital divisions) The basic principal here is that an element is inclined to like to have a completed outer or valence shell.
This is the shell that other atoms are coming into contact with and they too like to have a completed outer shell.
For an element with few electrons this is more important than to
one
with very many, so as we get to elements with d and f orbitals filling,
these tend to be more lackadaisical about how many electrons they will
trade
off or accept.
This desire to have a completed valence shell is why hydrogen (atomic number 1), with only one electron in its 1s obital, will readily bind with almost anything, including itself, to have a completed pair of electons in its s orbital.
And why helium (atomic number 2) its unwilling to accept or lose electrons and so is very inert (or unreactive).
It's fine as is, why react with anything?
The same is true for all the noble gases, the elements on the far right hand side of the periodic table, in column VIIIA or 18.
Neon (atomic number 10) has two electrons in its 1s orbital, two electrons in its 2s orbital, and two electrons in each of its three p orbitals for a total of 10 electrons.
It has completed its n=3D1 and n=3D2 quantum shells and is completely content, with no desire to gain, lose, or share electrons or any of these type of stategies that elements with incomplete electron shells use.
Flourine (atomic number 9) on the other hand, is one electron shy of completion, and so is extremely reactive.
It really wants to be like neon and with only one more electron it could be fully satisfied. It forms a molecular combintion with itself just for solace, as does chlorine, bromine, iodine and astatine beneath it in column VIIA or 17.
These also are just one electron shy of completion.
Becoming molecules is one of the main stategies elements that are a few electrons short of filling their valence shells use for a form of completion.
Then there are the elements on the far left side of the table. They are just one electon over the contented noble gas stage.
Ie. sodium. It has its 1s and 2s shells complete and one extra pesky electron in its 2s shell.
(This is also true of the other elements beneath it in colunm IA or 1; they are all very reactive. Elements in column IIA or 2 want to lose two electrons.)
What to do?
One answer is to lose the electron(s) and become an ion.
An element that has two or three spare electrons will also tend to lose them because its still far from filling all the orbitals in its valence shell.
Becoming an ion is the strategy of most elements classified as metals.
After argon (atomic number 18), the 4s orbital is filled (in
potassium-atomic number 19, and calcium-atomic number 20) and then the 3d
orbitals begins to be filled.
This is due to the fact that the d orbitals have a slightly higher energy level than the s orbitals of the next up shell, so they are filled after the next up "s" is filled.
The 4p orbitals are then filled, resulting in another noble gas, krypton at the far right.
This pattern continues.
The far left block (made up of the alkali metals and alkaline earth metals) are elements filling the s orbitals.
The middle block (transition metals) is composed of elements filling their d orbitals.
The far right block (mainly non-metals) are all elements filling their p orbitals.
The part pulled out (the inner transition metals, the Lanthanides and Actinides), are elements filling their f orbitals.
Click here for a more detailied history of the development of the periodic table.
Although Dmitri Mendeleev is often considered the "father" of the periodic table, the work of many scientists contributed to its present form.
During the next 200 years, a vast body of knowledge concerning the properties of elements and their compounds was acquired by chemists (view a 1790 article on the elements). By 1869, a total of 63 elements had been discovered.
As the number of known elements grew, scientists began to recognize patterns in properties and began to develop classification schemes.
In 1817 Johann Dobereiner noticed that the atomic weight of strontium fell midway between the weights of calcium and barium, elements possessing similar chemical properties. In 1829, after discovering the halogen triad composed of chlorine, bromine, and iodine and the alkali metal triad of lithium, sodium and potassium he proposed that nature contained triads of elements the middle element had properties that were an average of the other two members when ordered by the atomic weight (the Law of Triads).
This new idea of triads became a popular area of study. Between 1829 and 1858 a number of scientists (Jean Baptiste Dumas, Leopold Gmelin, Ernst Lenssen, Max von Pettenkofer, and J.P. Cooke) found that these types of chemical relationships extended beyond the triad. During this time fluorine was added to the halogen group; oxygen, sulfur,selenium and tellurium were grouped into a family while nitrogen, phosphorus, arsenic, antimony, and bismuth were classified as another. Unfortunately, research in this area was hampered by the fact that accurate values of were not always available.
If a periodic table is regarded as an ordering of the chemical elements demonstrating the periodicity of chemical and physical properties, credit for the first periodic table (published in 1862) probably should be given to a French geologist, A.E.Beguyer de Chancourtois. De Chancourtois transcribed a list of the elements positioned on a cylinder in terms of increasing atomic weight. When the cylinder was constructed so that 16 mass units could be written on the cylinder per turn, closely related elements were lined up vertically.
This led de Chancourtois to propose that "the properties of the elements are the properties of numbers." De Chancourtois was first to recognize that elemental properties reoccur every seven elements, and using this chart, he was able to predict the the stoichiometry of several metallic oxides. Unfortunately, his chart included some ions and compounds in addition to elements.
John Newlands, an English chemist, wrote a paper in 1863 which classified the 56 established elements into 11 groups based on similar physical properties, noting that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. In 1864 Newlands published his version of the periodic table and proposed the Law of Octaves (by analogy with the seven intervals of the musical scale). This law stated that any given element will exhibit analogous behavior to the eighth element following it in the table.
There has been some disagreement about who deserves credit for being the "father" of the periodic table, the German Lothar Meyer (see a picture) or the Russian Dmitri Mendeleev. Both chemists produced remarkably similar results at the same time working independently of one another. Meyer's 1864 textbook included a rather abbreviated version of a periodic table used to classify the elements.
This consisted of about half of the known elements listed in order of their atomic weight and demonstrated periodic valence chages as a function of atomic weight. In 1868, Meyer constructed an extended table which he gave to a colleague for evaluation. Unfortunately for Meyer, Mendeleev's table became available to the scientific community via publication (1869) before Meyer's appeared (1870).
Dmitri Ivanovich Mendeleev (1834-1907), the youngest of 17 children was born in the Siberian town of Tobol'sk where his father was a teacher of Russian literature and philosophy (see a picture). Mendeleev was not considered an outstanding student in his early education partly due to his dislike of the classical languages that were an important educational requirement at the time even though he showed prowess in mathematics and science.
After his father's death, he and his mother moved to St. Petersburg to pursue a university education. After being denied admission to both the University of Moscow and St. Petersburg University because of his provincial background and unexceptional academic background, he finally earned a place at the Main Pedagogical Institute (St. Petersburg Institute).
Upon graduation, Mendeleev took a position teaching science in a gymnasium. After a time as a teacher, he was admitted to graduate work at St. Petersburg University where he earned a Master's degree in 1856. Mendeleev so impressed his instructors that he was retained to lecture in chemistry.
After spending 1859 and 1860 in Germany furthering his chemical studies, he secured a position as professor of chemistry at St. Petersburg University, a position he retained until 1890. While writing a textbook on systematic inorganic chemistry, Principles of Chemistry, which appeared in thirteen editions the last being in 1947, Mendeleev organized his material in terms of the families of the known elements which displayed similar properties. The first part of the text was devoted to the well known chemistry of the halogens.
Next, he chose to cover the chemistry of the metallic elements in order of combining power -- alkali metals first (combining power of one), alkaline earths (two), etc.
However, it was difficult to classify metals such as copper and mercury which had multiple combining powers, sometimes one and other times two. While tryuing to sort out this dilema, Mendeleev noticed patterns in the properties and atomic weights of halogens, alkali metals and alkaline metals. He observed similarities between the series Cl-K-Ca , Br-/Rb-Sr and I-Cs-Ba.
In an effort to extend this pattern to other elements, he created a card for each of the 63 known elements. Each card contained the element's symbol, atomic weight and its characteristic chemical and physical properties. When Mendeleev arranged the cards on a table in order of ascending atomic weight grouping elements of similar properties together in a manner not unlike the card arrangement in his favorite solitare card game, patience, the periodic table was formed.
From this table, Mendeleev developed his statement of the periodic law and published his work on the Relationship of the Properties of the Elements to their Atomic Weights in 1869 (view a copy of Mendeleev's table as published in Annalen suppl. VIII, 133 (1871).
The advantage of Mendeleev's table over previous attempts was that it exhibited similarities not only in small units such as the triads, but showed similarities in an entire network of vertical, horizontal, and diagonal relationships.
In 1906, Mendeleev came within one vote of being awarded the Nobel Prize for his work.
At the time that Mendeleev developed his periodic table since the
experimentally determined atomic masses were not always accurate, he
reordered elements despite their accepted masses. For example, he
changed
the weight of beryllium from 14 to 9. This placed beryllium into Group
2
above magnesium whose properties it more closely resembled than where
it had
been located above nitrogen.
In all Mendeleev found that 17 elements had to be moved to new positions from those indicated strictly by atomic weight for their properties to correlate with other elements. These changes indicated that there were errors in the accepted atomic weights of some elements.
From the gaps present in his table, Mendeleev predicted the existence and properties of unknown elements which he called eka-aluminum, eka-boron, and eka-silicon. The elements gallium, scandium and germanium were found later to fit his predictions quite well.
In 1895 Lord Rayleigh reported the discovery of a new gaseous element named argon which proved to be chemically inert. This element did not fit any of the known periodic groups.
In 1898, William Ramsey suggested that argon be placed into the periodic table between chlorine and potassium in a family with helium, despite the fact that argon's atomic weight was greater than that of potassium.
This group was termed the "zero" group due to the zero valency of the elements. Ramsey accurately predicted the future discovery and properties neon.
Although Mendeleev's table demonstrated the periodic nature of the elements, it remained for the discoveries of scientists of the 20th Century to explain why the properties of the elements recur periodically.
In 1911 Ernest Rutherford published studies of the scattering of alpha particles by heavy atom nuclei which led to the determination of nuclear charge. He demonstrated that the nuclear charge on a nucleus was proportional to the atomic weight of the element. Also in 1911, A. van den Broek in a series of two papers (1, 2)proposed that the atomic weight of an element was approximately equal to the charge on an atom.
This charge, later termed the atomic number, could be used to number the elements within the periodic table. In 1913, Henry Moseley published the results of his measurements of the wavelengths of the x-ray spectral lines of a number of elements which showed that the ordering of the wavelengths of the x-ray emissions of the elements coincided with the ordering of the elements by atomic number.
With the discovery of isotopes of the elements, it became apparent that atomic weight was not the significant player in the periodic law as Mendeleev, Meyers and others had proposed, but rather, the properties of the elements varied periodically with atomic number.
The question of why the periodic law exists was answered as scientists
developed an understanding of the electronic structure of the elements
beginning with Niels Bohr's studies of the organization of electrons
into
shells through G.N. Lewis' discoveries of bonding electron pairs.
The last major changes to the periodic table resulted from Glenn Seaborg's work in the middle of the 20th Century. Starting with his discovery of plutonium in 1940, he discovered all the transuranic elements from 94 to 102.
He reconfigured the periodic table by placing the actinide series below the lanthanide series. In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. Element 106 has been named seaborgium (Sg) in his honor.
This history Edited from Western Oregon University Copyright 1997 Western Oregon University