Elements on the periodic table are arranged in groups according to their electron configuration. All elements in a vertical column have the same number of electrons in their outer energy level. Their distance from the nucleus increases as you go down a column. Electrons in the outer energy level are the electrons involved in chemical reactions. Since all the elements in a column have the same number of electrons available for chemical change, their properties are similar and are often referred to as families. The horizontal rows or periods also have predictable trends in characteristics because as you move left to right in a row only one electron is added changing the atomic number by one. Variations of properties are based on electron configuration. Predictable trends that occur repeatedly in a definite pattern are periodic properties or periodicity.Dmitry Ivanovich Mendeleev arranged the known elements according to mass in the development of the first table and found elements with similar properties were grouped together. This was prior to the knowledge of electron configuration. Properties and position of an element on the periodic table are based on their electron configuration. The most predictable and regular patterns are found in the s and p block elements (groups 1,2, 13 through 17). The d block elements do not show the same regularity in properties as the main group elements because the electrons in both the s and d sublevels can be involved in chemical reactions. The periodicity of the following properties can be predicted; valence electrons, atomic and ionic radii, ionization energy, electron affinity and electronegativity.
Valence electrons are the number of electrons available that can be lost, gained or shared in the formation of a compound. Since elements are grouped in vertical columns according to the number of outer shell electrons their group number represents the valence electrons. Group one and two have one and two outer shell electrons. Groups 13-18 have ten less electrons than the group number. For example, Group 14 elements have four valence electrons or four electrons in their outer shell available to react.
The atomic radius is half the distance between the nuclei of the same atom in an element or a compound. Atomic size, generally increases down a group because electrons are being added to higher energy levels at a greater distance from the nucleus increasing the size of the electron cloud. Atomic size decreases from left to right in a row or period because electrons are being added to the same energy level. As a result, the attraction between the positively charged nucleus and the number of electrons increases, pulling the electrons toward the center of the atom.
The ionic radius refers to the radius of an ion in an ionic compound. Atoms gain or lose electrons to form ions. Group one metals lose one electron and group two metals lose two electrons to become cations. Group 15 through 17 elements form anions more frequently by gaining electrons. Anions, therefore have more electrons than their neutral atom and are larger. Positively charged cations are smaller than their neutral atoms because they lose electrons. Going left to right across a period cations decrease in size as more electrons are lost. Beginning with group 15, anions decrease in size because fewer electrons are being added. Going down in a group electrons are being added to higher levels, increasing the size of the ions.
Ionization energy or ionization potential can also be predicted. The ionization energy is the energy needed to remove an electron from a gaseous atom. The first ionization energy refers to the energy needed to remove the most loosely held outer shell electron. Ionization energy decreases as you go down a column or group because the outer shell electrons are farther and farther away from the nucleus and are less influenced by its positive charge. The electrons are shielded (shielding effect) by inner shell electrons. Therefore, less energy is required to remove an electron. The element at the top of the column is held more tightly because the outer shell electron is closer to the positively charged nucleus. Ionization energies generally tend to increase across a row or period because electrons are being added to the same energy level and the attraction between the nucleus and the increasing number of negatively charged electrons in the shell holds them more tightly. Second, third and higher ionization energies progressively increase because each successive electron is attracted to a larger positive charge. A large jump in ionization energy generally indicates removal of the electron from a filled level or sublevel. Group 18, the noble gases have the highest first ionization energies because their energy level is full.
Electron affinity is the energy released when an electron is added to a neutral atom. Electron affinity can have positive or negative values. The addition of one electron is an exothermic reaction and the electron affinity value will be negative. Positive values indicate energy is needed (endothermic) to add another electron. An alternate definition of electron affinity is the attraction an atom has for an electron. Electron affinity is affected by the same factors as ionization energy. An element's valence shell that loses electrons easily will have little desire to attract electrons. Electron affinity increases from left to right in a row indicating a greater ease in acquiring electrons and decreases from top to bottom in a group. Metals give up electrons easily and have low electron affinities, nonmetals gain electrons more easily and have high electron affinities.
Electronegativity refers to the ability of an atom in a compound to attract electrons. This differs from ionization energy and electron affinity which describes the properties of single atoms. Electronegativity increases from left to right since metals tend to give up electrons and non metals tend to gain electrons. In the formation of a compound such as magnesium chloride, chlorine attracts electrons in the compound.Fluorine is the most electronegative element because it has a high ionization energy and a small radius.Cesium and Francium have the lowest electronegative values. Electronegativity decreases from top to bottom because larger atoms ionize easily and do not attract electrons. Smaller atoms which do not ionize readily attract electrons. The electronegativity scale ranges from 0.0 to 4.0 with 4.0 being the highest value.Linus Pauling developed this scale and in 1954 received a Nobel prize in chemistry for his work on chemical bonding which included electronegativity.
Trends in atomic and ionic radii are influenced most by the number of electrons in the element. Ionization energy, electron affinity, and electronegativity are directly related to nuclear charge and the distance of the outer shell electrons from the nucleus. Metallic and nonmetallic characteristics show a pattern also. Metallic characteristics increase down a column and decrease left to right. Periodicity is also observed in atomic numbers. Going down a column on the periodic table the atomic numbers differ by eight between rows one and two and two and three. Between rows three and four and four and five the difference in atomic numbers are eighteen. The difference in atomic numbers between rows five and six is thirty two. For example, calcium is in row four and has an atomic number of 20. The element below it is strontium in row five with an atomic number of 36. Periodicity continues to be an important scientific tool in understanding the behavior and interaction of elements.
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