THE PERIODIC Table

1. CHAPTER 6 THE PERIODIC Table

2. Organizing the elements February 17th, 1869 � University of St. Petersburg (Russia) - Dmitri Ivanovitch Mendeleev was writing a textbook on chemistry. He had the properties of each element written on a separate card, when, while shuffling through the cards he noticed that if he arranged them in order of increasing atomic mass, certain properties were repeated several times. He called this periodicity (repeating pattern). Soon thereafter he arranged them into a table, a predecessor to the modern periodic tables we use today. Mendeleev built the table by lining up the elements in a horizontal row in order of increasing atomic mass. Every time he came to an element with properties similar to one already in the row, he started a new row. The columns contained elements with similar properties. He also left empty spaces in the table, which he predicted the existence and properties of elements that would fill these spaces. In later years, gaps in his table were filled as other predicted elements were discovered. Examples: Scandium, Gallium, and Germanium.

3. Organizing the elements Mendeleev published his first periodic table in 1869. Mendeleev�s periodic table left 2 unanswered questions: Why was it that most elements could be arranged in order of increasing atomic mass, but a few could not? Example: Tellurium and Iodine, were assigned masses of 128 and 127, respectfully, but with chemical similarities meant that he had to place Te in the same group with sulfur and I in the same group as chlorine. 2. What is the reason for chemical periodicity?

4. Organizing the elements 1913 � H.G.J. Moseley (working with Ernest Rutherford) � bombarded many different metals with electrons in a cathode-ray tube and observed the x-rays emitted by the metals. Most importantly, he found that the wavelengths of x-rays emitted by a particular elements are related in a precise way to the atomic number of that element. Based on his experiments, he realized that other atomic properties may be similarly related to atomic number and not to atomic mass. If the elements are arranged in order of increasing atomic number, the defects in the Mendeleev table are corrected.

5. Organizing the elements Mendeleev�s Law of Chemical Periodicity should be restated as �the properties of the elements are periodic functions of atomic number.� Periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic numbers. In other words, when the elements are arranged in order of increasing atomic number, elements with similar properties recur at regular intervals.

6. Organizing the elements The periodic table has undergone extensive change since Mendeleev�s time. J. Thompson � Arranged the periodic table into 7 horizontal rows and 18 (32) vertical columns. Horizontal rows = Periods and/or Series Vertical columns = Families and/or Groups Periods 4, 5, 6, and 7 are called long periods.

7. Organizing the elements Chemists have discovered many new elements, and in recent years, synthesized new ones in the laboratory. Each of the more than 40 new elements can be placed in a group of other elements with similar properties. Periodic Table: An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.

8. Organizing the elements One of the most significant additions to the periodic table was the discovery of the noble gases. In 1894, John William Strutt (Lord Rayleigh) and Sir William Ramsay discovered argon, a gas in the atmosphere that had previously escaped notice because of its total lack of chemical reactivity. In 1868, Helium, had been discovered as a component of the sun, based on the emission spectrum of sunlight. In 1895, Ramsay showed that helium also exists on Earth. In order to fit argon and helium into the periodic table, Ramsay proposed a new group between group 17 and group 1. In 1898, Ramsay discovered two more noble gases to place in his new group (Krypton and Xenon). The final noble gas, Radon, was discovered in 1900 by Friedrich Ernst Dorn.

9. Organizing the elements The next step in the development of the periodic table was completed in the early 1900�s. The 14 elements corresponding to the filling of the 4f orbitals are known as the lanthanide (rare earth) elements (atomic numbers 58 to 71). The lanthanide elements are set below the other elements to avoid making the periodic table unduly wide. The properties of the elements are all quite similar, and they occur together in nature. For many years it was virtually impossible to separate them from one another.

10. Organizing the elements Another major step in the development of the periodic table was the discovery of the actinides (atomic numbers 90 to 103). All these elements are radioactive, and only thorium and uranium occur in nature. The other actinides have been synthesized in the laboratory by nuclear reactions. Their stability decreases rapidly with increasing atomic number. The longest lived isotope of nobelium (102No) has a half-life of about 3 minutes; that is, in 3 minutes half of the sample decomposes. In 1945, Glenn Seaborg made the revolutionary suggestion that the actinides, like the lanthanides, were filling the f sublevel, and that they should be removed from the periodic table and placed below the lanthanides. Element 106 � Seaborgium, was named after him.

11. Classifying the elements The vertical columns, or groups, of the periodic table contain elements having similar chemical and physical properties, and several groups of elements have distinctive names that are useful to know.

12. Classifying the elements Group 1 � H and the Alkali Metals Elements in the leftmost column, Group 1A, are known as the Alkali Metals (Li, Na, K, Rb, Cs, Fr). Except for hydrogen, all are metals and are solids at room temperature. In contrast, Hydrogen, a gas under normal comditions, consists of diatomic, or �two-atom,� molecules. The metals of Group 1 are all very reactive Example: They react with water to produce hydrogen and alkali solutions. Because of their reactivity, these metals are only found in nature combined in compounds, never as the free element. All these metallic elements form compounds with oxygen that have formulas like A2O, where A represents the alkali metal. Examples: Li2O, Na2O, K2O, Rb2O, Cs2O, and even hydrogen follows the same general formula, H2O. Because of their extreme reactivity with air and moisture, they are usually stored under kerosene (except Li). All (except Fr) have a silvery appearance. All are malleable, which means that they are soft enough to cut with a knife. Li, Na, and K are all less dense than water. All atoms in this group contain one electron in their outermost energy level (ns1).

13. Classifying the elements