The first ionization energy, I, of an atom is a measure of the energy change associated with removing an electron from the atom to form a cation.3 For example, the first ionization energy of Cl(g), 1251 kJ/mol, is the energy change associated with the process
The second ionization energy, I2, is the energy needed to remove the second electron, and so forth, for successive removals of additional electrons. Thus, I2 for the sodium atom is the energy associated with the process. The greater the ionization energy, the more difficult it is to remove an electron. Ionization energies for a given element increase as successive electrons are removed (I1< I2 < I3). This trend exists because with each successive removal, an electron is being pulled away from an increasingly more positive ion, requiring increasingly more energy.3
Most atoms can also gain electrons to form anions. The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity because it measures the attraction, or affinity, of the atom for the added electron. For most atoms, energy is released when an electron is added. It is important to understand the difference between ionization energy and electron affinity: Ionization energy measures the ease with which an atom loses an electron, whereas electron affinity measures the ease with which an atom gains an electron. The greater the attraction between an atom and an added electron, the more negative the atom’s electron affinity. For some elements, such as the noble gases, the electron affinity has a positive value, meaning that the anion is higher in energy than are the separated atom and electron:
The fact that the electron affinity is positive means that an electron will not attach itself to an Ar atom; the Ar– ion is unstable and does not form. The halogens, which are one electron shy of a filled p subshell, have the most-negative electron affinities. By gaining an electron, a halogen atom forms a stable anion that has a noble-gas configuration. The addition of an electron to a noble gas, however, requires that the electron reside in a higher-energy subshell that is empty in the atom. Because occupying a higher-energy subshell is energetically unfavorable, the electron affinity is highly positive. The electron affinities of Be and Mg are positive for the same reason as the added electron would reside in a previously empty p subshell that is higher in energy. The electron affinities of the group 5A elements are also interesting. Because these elements have half-filled p subshells, the added electron must be put in an orbital that is already occupied, resulting in larger electron-electron repulsions. Consequently, these elements have electron affinities that are either positive (N) or less negative than the electron affinities of their neighbors to the left (P, As, Sb). Recall that in Section 7.4 we saw a discontinuity in the trends for first ionization energy for the same reason. Electron affinities do not change greatly as we move down a group. For flourine, for instance, the added electron goes into a 2p orbital, for chlorine a 3p orbital, for bromine a 4p orbital, and so forth. As we proceed from F to I, therefore, the average distance between the added electron and the nucleus steadily increases, causing the electron-nucleus attraction to decrease. However, the orbital that holds the outermost electron is increasingly spread out, so that as we proceed from F to I, the electron-electron repulsions are also reduced. As a result, the reduction in the electron-nucleus attraction is counterbalanced by the reduction in electron-electron repulsions.
Atoms may share electrons through covalent bonds, but that doesn’t mean they share equally. When the two atoms involved in a bond aren’t the same, the two positively charged nuclei have different attractive forces. They pull on the electron pair to different degrees and the end result is that the electron pair is shifted towards one atom. Electronegativity is the strength an atom has to attract a bonding pair of electrons to itself. The larger the electronegativity value, the greater the atom’s strength to attract a bonding pair of electrons. A bond in which the electron pair is shifted toward one atom is called a polar covalent bond. The atom that more strongly attracts the bonding electron pair is slightly more negative, and the other atom is slightly more positive. The larger the difference in the electronegativities, the more negative and positive the atoms become. Consider hydrogen chloride (HCl). Hydrogen has an electronegativity of 2.1, and chlorine has an electronegativity of 3.0. Because chlorine has a larger electronegativity value, the electron pair that’s bonding HCl together shifts toward the chlorine atom. If the two atoms have extremely different electronegativities, the atoms will probably form ionic, not covalent bonds. For instance, sodium chloride (NaCl) is ionically bonded. An electron has transferred from sodium to chlorine. Sodium has an electronegativity of 1.0, and chlorine has an electronegativity of 3.0. That’s an electronegativity difference of 2.0 (3.0 – 1.0), making the bond between the two atoms very, very polar.
Electron shells are defined by the principle quantum number and going down the periodic table means moving to the next shell. As the next shell is filled (Ne to Na), the effective nuclear charge decreases because the inner shell stands in between the nucleus and the new shell. Filling to the next shell causes a jump in atom size because of decreased effective nuclear charge. As you go down a group (Na to K), the atomic size increases even though the effective nuclear charge stays the same, because higher shells have a larger radius than lower shells. Going across the periodic table means filling up the same shell (by going through subshells). As you fill up a shell, the effective nuclear charge increases because the atomic number (protons) is increasing while the same-shell electrons you add do not shield one another. With increasing effective nuclear charge, the electrostatic attraction between the nucleus and the electrons increases, so the atom becomes more compact. The increasing effective nuclear charge and electrostatic attraction is why going across a periodic table means decreasing atomic size. Atomic size increases as you go down a column and decreases as you go across (to the right of) a row. Atomic sizes may overlap on the periodic table.
Summary of periodic trends
3) Javadi, H. (2011). Effective Nuclear Charge. Retrieved from
4) Lufaso, M. (2016). Electronic Structure of Atoms. Retrieved from University of North Florida: http://www.unf.edu/~michael.lufaso/chem2045/Chapter6.pdf