Could you please explain in detail the concept and trends of the following? I had a hard time understanding in class and from the textbook: electronegativity, electron affinity, atomic radius, ionic radius, and ionization energy.
EDIT: oops typo about atomic radius trends
Electronegativity: measure of how tightly an element will hold onto an electron; the higher the value, the more tight the hold. Thus, the difference in electronegativity between two elements allows one to examine the nature of their bonding (whether it is ionic, polar covalent, or pure covalent). For example, the extremely large difference in electronegativity between Na and Cl, with Cl having the larger value, means that Cl will pull the electron cloud of the Na-Cl bond greatly towards itself, allows one to confirm that the bond in rock salt, NaCl, is ionic. SIDENOTE: The descending order of nonmetal electronegativities goes roughly FONClBrISCP, which kinda sounds like a funny word, so it’s easy to remember!
TREND: Electronegativity increases across a period, and decreases down a group. Fluorine is the most electronegative element. These trends do not include noble gases.
Electron Affinity (EA):The amount of energy released when the atom/molecule recieves an electron. NOTE: This means that when you see a negative EA value, it implies that the atom will absorb energy when taking in the electron. For instance, the 2nd EA value of oxygen is negative, which makes sense since an electron is being added to an already negative ion O-, which is unfavorable (making the ion more negative), thus raising the energy of the product O2-.
TREND: Electron affinity really does not follow the trends too well, but overall tends to increase going across a period. The values going down a group are rather variable. Frankly, a table of values is probably your best bet.
Atomic Radius: Measures the mean distance from the electron cloud to the nucleus. The main thing to know about the atomic radius is that the larger the value, the farther away electrons are from the nucleus and consequently less tightly held they are. Atomic radius actually explains pretty much all the other periodic trends. The larger the radius, the more easy it is to take an electron, and the lower the ionization energy and electronegativity values. The smaller the radius, the more tightly the nucleus holds onto electrons and would hold onto any incoming electrons, thus demonstrating the higher ionization energy and electronegativity.
TRENDS: Atomic radius decreases across a period, which makes sense as effective nuclear charge increases and thus attracts the electrons greatly. Atomic radius increases down a group, which also makes sense as we go up an electron shell and consequently the ends of the electron cloud move further from the nucleus. However, atomic radius doesn’t increase significantly across the transition metals because new electrons are being added to the d orbitals, which are on a lower shell than the s and p orbitals, and do not increase the size of the electron cloud.
Ionic Radius: Measure of the distance from the electron cloud to the nucleus in an ion. This works the same way atomic radius does: a higher effective nuclear charge will decrease the ionic radius. The important thing to know for ionic radius is that the ions of many elements will often have the same electron configuration. For example, the oxygen ion is most commonly -2, while the fluorine ion is most commonly -1. After gaining their respective electrons, these two ions have the same number of electrons. So, to determine which one has a smaller ionic radius, look at their nuclei. The fluorine nucleus has more protons than the oxygen nucleus does, so it will attract electrons more strongly and decrease the ionic radius to become smaller than oxygen.
Ionization energy: a measure of the energy needed to remove an electron from an atom. The removal of an electron to ionize, or make an ion, of an element will always require energy. This goes hand in hand with electronegativity, which should make sense; an atom that holds electrons tightly will require more energy to ionize.
TRENDS: Since ionization energy and electronegativity are so closely related, it makes sense that their trends are similar too: generally, ionization energy increases across a period and decreases down a group. However, there are some exceptions. In period 1, the ionization energy for boron is less than beryllium, and oxygen’s is less than nitrogen’s. This is because, in the case of beryllium, the removed electron will be in a p orbital, which is of higher energy than boron’s s orbital electron, meaning that less energy is neaded to remove it. In oxygen, the electron to be removed is found in a full orbital with 2 electrons, which repel each other and make it easier for the electron to be removed than nitrogen’s electron, which is the only one in its orbital and experiences less electrostatic repulsion.
This is another long answer, so feel free to message us with any questions!
