If you are interested in learning how the number of valence electrons affects the ionization energy of an element, you have come to the right place. In this article, we will explain what valence electrons and ionization energy are, how they are related, and what are some of the periodic trends that affect them. Let’s get started!
What are Valence Electrons?
Valence electrons are the electrons in the outermost shell of an atom. They are the ones that participate in chemical bonding and reactions, and determine the chemical properties of an element. For example, elements in the same group (column) of the periodic table have the same number of valence electrons, and therefore tend to have similar chemical behavior.
The number of valence electrons can be easily found by looking at the group number of an element. For example, carbon is in group 14, so it has four valence electrons. Sodium is in group 1, so it has one valence electron. The exception is helium, which has only two valence electrons despite being in group 18.
What is Ionization Energy?
Ionization energy is the energy required to remove an electron from an atom. The ionization energy is correlated with the strength of attraction between the positively-charged nucleus and the negatively-charged valence electrons. The higher the ionization energy, the stronger the attractive force between nucleus and valence electrons, and the more energy is required to remove a valence electron.
The energy required to remove the outermost valence electron from a neutral atom is the first ionization energy. The second ionization energy is that required to remove the next electron, and so on. Each succeeding ionization energy is larger than the preceding one, because it becomes harder to remove electrons from a more positively charged species.
How are Valence Electrons and Ionization Energy Related?
The number of valence electrons affects the ionization energy of an element in two ways: by influencing the effective nuclear charge and by influencing the electron-electron repulsion.
- Effective nuclear charge (Z*) is the net positive charge experienced by a valence electron due to the nucleus and the inner electrons. The more protons in the nucleus, the higher the Z*, but the more inner electrons, the lower the Z*, because they shield or screen some of the nuclear charge from reaching the valence electrons. In general, Z* increases across a period (row) as more protons are added to the nucleus without adding more inner shells, but decreases down a group as more inner shells are added that increase shielding.
- Electron-electron repulsion is the force that pushes electrons away from each other due to their negative charges. The more electrons in a shell or subshell, the higher the repulsion, and the easier it is to remove an electron from that shell or subshell.
The combination of these two factors determines how tightly or loosely a valence electron is held by an atom, and thus how much ionization energy is needed to remove it.
What are Some Periodic Trends in Ionization Energy?
There are some general trends in ionization energy that can be observed across the periodic table:
- Across a period: As Z* increases across a period, the ionization energy of the elements generally increases from left to right. However there are breaks or variation in these trends in some cases:
- Ionization energy is especially low when removal of an electron creates a newly empty p or d subshell (examples include I1 of B, Al, Ga)
- Ionization energy is especially low when removal of an electron results in a half-filled p or d subshell (examples include I1 of O, S)
- Ionization energy increases more gradually across the d- and f-subshells compared to s- and p- subshells. This is because d- and f- electrons are weakly penetrating and experience especially low Z*.
- Down a group: As Z* decreases down a group due to increased shielding by inner shells, the ionization energy of elements generally decreases from top to bottom. However there are breaks or variation in these trends in some cases:
- Ionization energy increases slightly when removal of an electron creates a filled p or d subshell (examples include I1 of N, P)
- Ionization energy decreases sharply when removal of an electron creates a new principal quantum level (examples include I1 of Li, Na)
These trends can be seen in Figure 1 below:
Figure 1: Periodic trends in first ionization energies (source: Wikipedia)
We have learned that the number of valence electrons affects the ionization energy of an element by influencing the effective nuclear charge and the electron-electron repulsion. The more valence electrons an element has, the lower its ionization energy, and vice versa. We have also learned some of the periodic trends in ionization energy that can help us predict the relative ease or difficulty of removing electrons from different elements. We hope this article has been helpful and informative for you. Thank you for reading!