Atomic Radius
Since in an atom there is no clear boundary beyond which the electron never strays, the way atomic radius is measured is by calculating the distance between the two nuclei of atoms when they are involved in a chemical bond. If the two bonded atoms are of the same element, you can divide the distance by 2 to get the atom’s radius. That said, one of the two important things you’ll need to know about atomic radii for the SAT II Chemistry exam is that
atomic radii decrease (
)
moving across a period from left to right. But why? It seems as though the more protons you add, the more space the atom should take up, but this is not the case. The reason for this lies in the basic concept that opposite charges attract each other and like charges repel each other. As you increase the number of protons in the nucleus of the atom, you increase the
effective nuclear charge of the atom (
Zeff), and the nucleus pulls more strongly on the entire electron cloud. This makes the atomic radius decrease in size. The second thing you’ll need to know is that
atomic radii increase moving down a group or family. This is easier to understand if you refer to the Bohr model. As you move down the table, the value of
n increases as we add another shell. Remember that the principal quantum number,
n, determines the size of the atom. As we move down a family, the attractive force of the nucleus dissipates as the electrons spend more time farther from the nucleus.
One more thing about atomic size. As you know, when an atom loses an electron, a cation, or positive ion, is formed. When we compare the neutral atomic radius to the cationic radius, we see that the cationic radius is smaller. Why? The protons in the nucleus hold the remaining electrons more strongly. As you might expect, for negatively charged ions, or anions, the nuclear attractive force decreases (and there is enhanced electron-electron repulsion), so the electrons are less tightly held by the nucleus. The result is that the anion has a larger radius than the neutral atom.
The SAT II Chemistry test might ask you to compare the sizes of two atoms that are isoelectronic, meaning that they have the same number of electrons. In this case, you would then consider the number of protons the two atoms possess.
Example
Which ion is larger, F– or O2-?
Explanation
Since these two atoms are isoelectronic and in the same period, the atom with more protons in its nucleus will hold its electrons more tightly and be smaller. Fluoride will be smaller since it has more protons (9, compared to oxide’s 8).
Ionization Energy (IE)
The ionization energy of an atom is the energy required to remove an electron from the atom in the gas phase. Although removing the first electron from an atom requires energy, the removal of each subsequent electron requires even more energy. This means that the second IE is usually greater than the first, the third IE is greater than the second, and so on. The reason it becomes more difficult to remove additional electrons is that they’re closer to the nucleus and thus held more strongly by the positive charge of the protons.
