These can typically be explained by their electron configuration. For example, Magnesium has a higher ionization energy than Aluminum. Magnesium has an electron configuration of [Ne]3s2. Magnesium has a high ionization energy because it has a filled 3s orbital and it requires a higher amount of energy to take an electron from the filled orbital. Electron Affinity Electron affinity (E.A.) is the energy change that occurs when an electron is added to a gaseous atom. Electron affinity can further be defined as the enthalpy change that results from the addition of an electron to a gaseous atom. It can be either positive or negative value. The greater the negative value, the more stable the anion is. - (Exothermic) The electron affinity is positive \[\ce{X(g) + e^{-} -> X^{-} + Energy} \nonumber \] - (Endothermic) The electron affinity is negative \[\ce{X(g) + e^{-} + Energy -> X^{-}} \nonumber \] It is more difficult to come up with trends that describe the electron affinity. Generally, the elements on the right side of the periodic table will have large negative electron affinity. All of these elements display several other trends and we can use the periodic law and table formation to predict their chemical, physical, and atomic properties. Understanding these trends is done by analyzing the elements electron configuration; all elements prefer an octet formation and will gain or lose electrons to form that stable configuration. Atomic Radius We can never determine the atomic radius of an atom because there is never a zero probability of finding an electron, and thus never a distinct boundary to the atom. All that we can measure is the distance between two nuclei (internuclear distance). In order to comprehend the extent of screening and penetration within an atom, scientists came up with the effective nuclear charge, \(Z_{eff}\). The equation for calculating the effective nuclear charge is shown below. \[Z_{eff}= Z - S \nonumber \] In the equation S represents the number of inner electrons that screen the outer electrons. Students can easily find S by using the atomic number of the noble gas that is one period above the element. For example, the S we would use for Chlorine would be 10 (the atomic number of Neon). Z is the total number of electrons in the atom. Cations have a smaller radius than the atom that they were formed from. With the loss of an electron, the positive nuclear charge out Electron Affinity Electron affinity (E.A.) is the energy change that occurs when an electron is added to a gaseous atom. Electron affinity can further be defined as the enthalpy change that results from the addition of an electron to a gaseous atom. It can be either positive or negative value. The greater the negative value, the more stable the anion is. (Exothermic) The electron affinity is positive \[\ce{X(g) + e^{-} -> X^{-} + Energy} \nonumber \] (Endothermic) The electron affinity is negative \[\ce{X(g) + e^{-} + Energy -> X^{-}} \nonumber \] It is more difficult to come up with trends that describe the electron affinity. Generally, the elements on the right side of the periodic table will have large negative electron affinity. The electron affinities will become less negative as you go from the top to the bottom of the periodic table. However, Nitrogen, Oxygen, and Fluorine do not follow this trend. The noble gas electron configuration will be close to zero because they will not easily gain electrons. Figure \(\PageIndex{7}\): Courtesy of Jessica Thornton (UCD) Electronegativity Electronegativity is the measurement of an atom to compete for electrons in a bond. The higher the electronegativity, the greater its ability to gain electrons in a bond. Electronegativity will be important when we later determine polar and nonpolar molecules. Electronegativity is related with ionization energy and electron affinity. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted by the positive nucleus on the negative electrons. Therefore the electronegativity increases from bottom to top and from left to right. Figure \(\PageIndex{8}\): Courtesy of Jesscia Thornton (UCD) Metallic Character The metallic character is used to define the chemical properties that metallic elements present. Generally, metals tend to lose electrons to form cations. Nonmetals tend to gain electrons to form anions. They also have a high oxidation potential therefore they are easily oxidized and are strong reducing agents. Metals also form basic oxides; the more basic the oxide, the higher the metallic character. They are also very nonreactive as they already have a full valence shell with 8 electrons. Therefore, the noble gases have little tendency to lose or gain electrons. Useful Relationships from the Periodic Table The periodic table of elements is useful in determining the charges on simple monoatomic ions. For main-group elements, those categorized in groups 1, 2, and 13-18, form ions they lose the same number of electrons as the corresponding group number to which they fall under. For example, K atoms (group 1) lose one electron to become K + and Mg atoms (group 2) lose two electrons to form Mg 2 + . The other main-group elements found in group 13 and higher form more than one possible ion. An element that is an example of a metalloid is (a) S; (b) Zn; (c) Ge; (d) Re; (e) none of these Answer: (c) Ge 2. In the periodic table, the vertical (up and down) columns are called (a) periods; (b) transitions; (c) families/groups; (d) metalloids; (e) none of these. Answer: (c) families/groups 3. Why are noble gases inert (nonreactive)? Answer: Noble gases are inert because they already have a full valence electron shell and have little tendency to gain or lose electrons. 4. What are compounds that contain a halogen called? Answer: Salts 5.