Chapter 5: The Periodic Law Coach Kelsoe Chemistry Pages 132-169

Section 5-1: History of the Periodic Table Coach Kelsoe Chemistry Pages 133-137

Section Objectives Explain the roles of Mendeleev and Moseley in the development of the periodic table. Describe the modern periodic table. Explain how the periodic law can be used to predict the physical and chemical properties of elements. Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number.

Origins of the Periodic Table In September of 1860, chemists assembled at the First International Congress of Chemists in Karlsruhe, Germany. Italian chemist Stanislao Cannizzaro presented a method for accurately measuring atomic masses.

Mendeleev and Chemical Periodicity Russian chemist Dmitri Mendeleev started writing a chemistry book, using the values agreed upon at the Congress of Chemists and organized all this information. He noticed that when the discovered elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals.

Mendeleev’s Periodic Table

Mendeleev and Chemical Periodicity Mendeleev created a table in which elements with similar properties were grouped together. His original table was arranged in order of increasing atomic mass, not atomic number. This “periodic table” was first published in 1869. Mendeleev originally set up his table horizontally and left gaps in the table where he predicted the existence of new elements.

Mendeleev and His Table Mendeleev was credited as the discoverer of the periodic law. “Periodic” means “having a repeating pattern.” There were two questions remaining about the periodic table though: Why could most of the elements be arranged in the order of increasing atomic mass but a few could not? What was the reason for chemical periodicity? Mendeleev never won a Nobel Prize!

Moseley and the Periodic Law The question of why elements were not always listed in order of atomic mass was unanswered for another 40 years. In 1911, English scientist Henry Moseley discovered that elements in the table fit better when arranged in order of increasing nuclear charge (atomic number).

Moseley and the Periodic Law Moseley’s discovery was consistent with Mendeleev’s ordering of the periodic table by properties instead of strictly atomic mass. Today, Mendeleev’s principle of chemical periodicity is correctly stated as the periodic law.

Periodic Law The periodic law states “the physical and chemical properties of the elements are periodic functions of their atomic numbers.” In our terms, when the elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals.

The Modern Periodic Table Today’s periodic table is a lot different from Mendeleev’s original table. The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.

The Noble Gases English physicist John William Strutt and Scottish chemist Sir William Ramsay discovered argon, only 26 years after discovering Helium. Ramsey proposed Group 18; gases that are generally unreactive. We call them Noble Gases.

The Lanthanides The next addition to the periodic table was the lanthanides. The lanthanides are the 14 atomic elements with atomic numbers from 58 (Cerium, Ce) to 71 (Lutetium, Lu).

The Lanthanides The lanthanides belong in period 6. Because these elements are so similar in chemical and physical properties, the process of separating them from the rest was a long process.

The Actinides The next major development of the periodic table came with the discovery of the actinides. The actinides are the 14 elements with atomic numbers from 90 (Thorium, Th) to 103 (Lawrencium, Lr).

The Actinides The actinides belong in period 7. The actinides as well as the lanthanides are usually set below the main portions of the periodic table. Both actinides and lanthanides fall between Groups 3 and 4.

Periodicity Periodicity with respect to atomic number can be observed in any group of elements. Group 1 elements are very similar with respect to their properties, but are different from each other and different from those in other groups. Groups 1-2, and 13-18 follow a pattern of 8, 8, 18, 18, and 32 with the exception of hydrogen.