The valence electrons are attracted towards the nucleus by the protons. Not always do the outer shell, valence, electrons feel the electrostatic attraction from all the protons in the nucleus. Inner shell electrons tend to shield the valence electrons from the positive nucleus. The attraction felt by the valence electrons is obtained by subtracting the inner shell electrons from the number of protons in the nucleus of the atom. We sometimes refer to this as the core charge. Take sodium for example.
The core charge is obtained by subtracting the inner shell electrons (10) from the 11 protons in the nucleus. The core charge is +1. So the valence electron is attracted towards the nucleus, effectively, by a single proton. The core charge of beryllium is +2. The valence electrons are attracted towards the nucleus by a core charge of +2
Calculate the effective charge for the following elements.
What is the trend for core charge as we go across a period?
What is the trend for core charge as we go down a group?
Consider the animation on the right. It shows the ionisation energies of the electrons in the sodium atom. Note that the closer an electron is to the nucleus the more energy required to remove it.
Removing the valence electron requires, relatively, little energy. After removing the valence electron we then move closer to the nucleus and start removing the second energy level electrons. You will notice a sharp rise in the energy required to remove the first electron from the second energy level. This is expected because we move closer to the nucleus, while experiencing a greater core and this results in a greater attraction to the nucleus.
Compare the first ionisation energies of He, Ne and Ar.
Do you notice a trend? decreases increases stays the same
How can this be best explained? The valence electrons are further from the nucleus as we go from He to Ar. The atom is getting smaller as we go from He through to Ar. The valence electrons experience a smaller core charge as we go from He through to Ar.