Reading the Periodic Table
The periodic table of elements is a fundamental tool in chemistry that organizes all known chemical elements based on their atomic number and electron configuration. Each element is represented by a unique symbol, often derived from its name or its historical nomenclature, such as "H" for hydrogen and "Fe" for iron, reflecting the rich history of scientific discovery. The table is structured in rows called periods and columns known as groups, where elements in the same group share similar chemical properties due to having the same number of valence electrons. For instance, alkali metals in the first column exhibit similar reactivity patterns.
The atomic number, indicating the number of protons in an atom, is prominently displayed above each element's symbol, while atomic mass is found below, typically represented as a decimal due to the presence of isotopes. Notably, the periodic table not only serves as a classification system but also reveals trends in elemental behavior, such as ionization energy and atomic radius. The concept of periodicity highlights the recurring properties of elements, which became clearer as scientists, including Dmitri Mendeleev, developed more comprehensive models in the late 19th century. The periodic table remains essential for understanding chemical interactions and behaviors, making it a cornerstone of chemistry education.
Reading the Periodic Table
FIELDS OF STUDY: Inorganic Chemistry; Organic Chemistry
ABSTRACT
The characteristics of the periodic table of the elements are discussed. The periodic table organizes the different elements in periods according to their number of electron shells and in groups according to the configuration of their valence electrons.
Introducing the Periodic Table
The periodic table is the first tool that students encounter in the study of chemistry. It is also the single most important depiction of the chemical elements, because each frame shows the principles underlying every one of the element’s behaviors. The information presented in the periodic table frame begins with the element symbol, the familiar letter or letter pair used to represent an individual atom in chemical formulas. Generally, each symbol is taken from the universally recognized element name assigned by the International Union of Pure and Applied Chemistry (IUPAC). For example, H stands for hydrogen, Cl for chlorine, and so on. Some symbols seem to defy this logic, however. The element symbol for tungsten is W, which certainly does not appear anywhere in the word; the symbol comes from the element’s earlier name, wolfram. Similarly, the symbols for copper (Cu) and iron (Fe) come from their original Latin names, cuprum and ferrum, respectively. Such symbols demonstrate a certain respect for the long tradition of science in the study of materials and the history of chemistry.
Structure of Periodic Table
The periodic table presents the known elements ordered by their atomic number. The atomic number, as defined by the modern atomic theory, is the precise number of protons that are contained in the nucleus of any given atom of the element. Hydrogen, atomic number 1, has only a single proton in its nucleus; helium, atomic number 2, has two; and so on. This number generally appears above the element symbol.
The atomic mass of each atom is almost exactly equal to the number of protons and neutrons in its nucleus and thus should be a whole number, as subatomic particles cannot be divided and still maintain their identity. Yet almost all atomic weights, which are displayed below the element symbols, are stated as decimals. This fractional value is due to the natural presence of isotopes of different elements, which are atoms that have the same number of protons in their nucleus but a different number of neutrons, thus changing the atomic mass of that particular atom. The atomic weight of an element is a proportional average of the atomic masses of the different isotopes, weighted according to the natural occurrence of each one. The element chlorine, for example, occurs naturally as two isotopes, one with a mass of 35 atomic mass units (u) and abundance of about 24 percent and the other with a mass of 37 u and abundance of 76 percent, giving it an atomic weight of 35.45. The same principle applies to all other naturally occurring elements.
According to modern atomic theory, electrons can possess only very specific amounts of energy, and only a maximum specified number of electrons can occupy each energy level, or electron shell, which corresponds to a defined area around the atomic nucleus. Each horizontal row, or period, represents the shell that contains the valence electrons of the elements within that period. The progression of the principal energy levels—that is, the outermost electron shells—is typically shown on the extreme right-hand side of the periodic table, at the end of each period. Thus, the first period, containing hydrogen (H) and helium (He), corresponds to the first energy level; the second period, containing lithium (Li), beryllium (Be), boron (B) carbon (C), nitrogen (N), oxygen (O), fluorine (F), and neon (Ne), corresponds to the second level; and so on. At the bottom of the periodic table are two sets of elements, the lanthanides and the actinides, which would normally be part of the sixth and seventh periods but are instead shown at the bottom to save space.
The elements in the periodic table are also arranged in vertical columns, or groups. The elements in each group all have the same number of valence electrons and therefore are similarly reactive. The first column, for example, comprises the alkali metals: hydrogen, lithium, sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). Each has just one electron in its outermost electron shell, which they all readily give up to form an ion with single positive charge. Accordingly, the alkali metals all form very similar types of compounds, undergo very similar reactions, and so on. Similarly, the group 2 elements—beryllium, magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra)—all have two valence electrons, which they readily give up to form ions with two positive charges.
As can be seen from this, the groups and periods of the periodic table contain essentially all of the information needed to understand the behavior of the chemical elements.
History of the Periodic Table
As the science of chemistry developed, many scientists recognized that certain elements exhibited similar chemical behaviors, and attempts were made to organize them in some way. However, because so many of the chemical elements were still unknown, these scientists had little success in establishing a meaningful system. As more materials came to be recognized as actual chemical elements, the periodicity of their behavior became more apparent. In the late nineteenth century, Russian chemist Dmitri Mendeleev (1834–1907) produced the first correctly ordered periodic table.
PRINCIPAL TERMS
- atomic mass: the total mass of the protons, neutrons, and electrons in an individual atom.
- atomic number: the number of protons in the nucleus of an atom, used to uniquely identify each element.
- element name: the official name by which each element is known, according to the standards of the International Union of Pure and Applied Chemistry (IUPAC).
- element symbol: a one- or two-letter abbreviation uniquely assigned to each element, usually derived from the internationally recognized name of the element.
- periodicity: the tendency of elements with similar electron distributions in their valence shells to exhibit similar chemical properties, such as ionization energy, atomic radius, and electronegativity.
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