The observed general trend in ionization energies shows two slight discontinuities in each period. Examine the electron configuration at these points and explain why such a discontinuity may occur.
Look at the electron configurations for example One of the discontinuities happen from Be to B and the other from N to O
In both these situations
Be and N were in a stable position ( full orbital or half full orbitals)
So they need more energy to go to next level
therefore, they lost some energy and their energy level went lower.
Well, let me put on my electron configuration joke hat for this one!
You see, when it comes to ionization energies, it's all about stealing electrons and playing hide and seek with the protons. Picture it like this:
In each period, as we move across the periodic table, we add one proton but also one electron to our atoms. So it's like playing a sneaky game of catch-up with protons and electrons. As the number of protons increases, they have a stronger pull on the electrons, making it more difficult for them to break free.
Now, imagine two electrons having a race. The first electron gets stolen by an atom pretty easily, but the second, oh boy, it puts up a fight! It needs a bit more energy to leave its comfort zone and join the big bad ion party. Hence, you observe a slight discontinuity in ionization energy. It's like the second electron is saying, "Nah, I'm cozy here, why would I leave?"
So, these slight discontinuities occur because those electrons are stubborn little creatures, making the ionization energy jump up a bit. It's like the electron world has its own version of the "two steps forward, one step back" dance move. They can't make things too easy for us, can they?
Hope that explanation didn't shock you too much with its electrifying humor!
The observed general trend in ionization energies shows two slight discontinuities in each period. These discontinuities occur when moving from the end of one subshell to the beginning of the next subshell within a given period.
To understand why these discontinuities occur, let's examine the electron configuration at these points. In the periodic table, the electron configuration represents the arrangement of electrons in an atom's energy levels or orbitals.
Each subshell has a different energy level, denoted by the principal quantum number (n) and is represented by a series of letters (s, p, d, f). The energy of these subshells generally increases with increasing values of n. For example, the 3p subshell has a higher energy than the 2p subshell.
Now, as we move across a period from left to right, the number of protons in the nucleus increases, resulting in a stronger attractive force between the nucleus and the electrons. This increased nuclear charge tends to pull the outermost electrons closer to the nucleus, making it harder to remove them.
However, there are two exceptions to this general trend. In the transition from the end of the s-block to the beginning of the p-block, and from the end of the d-block to the beginning of the p-block, we observe a slight discontinuity in ionization energies.
Let's take an example of the second period. The electron configuration at the end of the 2s subshell is 1s^2 2s^2. At the beginning of the 2p subshell, the electron configuration starts with 1s^2, and then it fills the 2p subshell. During this transition, there is a slight decrease in ionization energy.
This decrease occurs because the 2p subshell is higher in energy than the 2s subshell. The outermost electrons in the 2p subshell experience less attractive force from the nucleus compared to the 2s electrons. Thus, it is relatively easier to remove an electron from the 2p subshell, resulting in a lower ionization energy.
Similarly, in the transition from the end of the d-block to the beginning of the p-block, there is another slight discontinuity. For example, in the fourth period, at the end of the 3d subshell, the electron configuration is 3d^10 4s^2. At the beginning of the 4p subshell, the electron configuration starts with 3d^10, and then the 4p subshell is filled. During this transition, there is another slight decrease in ionization energy.
This decrease occurs because the 4p subshell is higher in energy than the 3d subshell and experiences less attractive force from the nucleus. Therefore, it is relatively easier to remove an electron from the 4p subshell compared to the 3d subshell, resulting in a lower ionization energy.
In summary, these slight discontinuities in ionization energies occur due to the energy differences between the subshells. The higher energy of the next subshell allows for a relatively easier removal of electrons, resulting in a temporary decrease in ionization energy.
To examine the electron configuration at the points where slight discontinuities occur in the observed general trend of ionization energies, we need to understand a few concepts about the periodic table and electron configurations.
First, let's recall what ionization energy is. Ionization energy is the amount of energy required to remove one electron from a neutral atom in the gaseous state. The general trend of ionization energies across a period on the periodic table typically increases from left to right.
Now, let's consider the electron configuration at these points. There are two common discontinuities in ionization energies in each period: a decrease from the group 2 to group 13 elements and a decrease from the group 15 to group 16 elements.
The electron configurations of these elements can help explain these discontinuities. For example, let's focus on the first discontinuity between group 2 (alkaline earth metals) and group 13 (boron group). Group 2 elements have a fully-filled s subshell (ns²) in their valence shell, while group 13 elements have an extra electron in the p subshell (ns²np¹). Removing an electron from the p subshell requires less energy than removing an electron from a fully-filled s subshell. Hence, the ionization energy decreases from group 2 to group 13.
Moving on to the second discontinuity between group 15 (nitrogen group) and group 16 (oxygen group), we can observe a similar pattern. Group 15 elements have a half-filled p subshell (ns²np³) in their valence shell, whereas group 16 elements have a fully-filled p subshell (ns²np⁴). Removing an electron from the half-filled p subshell is easier compared to removing an electron from the fully-filled p subshell, resulting in a decrease in ionization energy from group 15 to group 16.
In summary, the observed discontinuities in ionization energies occur due to the differences in electron configurations at those points. These changes in electron configurations affect the ease of removing electrons, thus leading to deviations from the general trend of increasing ionization energies across a period on the periodic table.