I need help understanding by boron does not follow the ionization trend across the period is lower than the first two elements in the same period, where as it should be higher.

has it got to do with the p block having an electron, which makes it easier to remove the electron hence the 1st ionization energy is lower than the general trend it should exhibit?

yes.

Yes, you are correct. The reason why boron does not follow the ionization trend across the period and has a lower first ionization energy compared to the first two elements in the same period (lithium and beryllium) is because it has an electron in a p orbital.

In general, the ionization energy of an atom increases as you move across a period from left to right because the atomic radius decreases and the effective nuclear charge (the positive charge felt by valence electrons) increases. As a result, it becomes more difficult to remove an electron.

However, when we specifically look at boron in the periodic table, we see that it has an electron in a p orbital. Unlike the s orbital, the p orbital has a higher energy and is further from the nucleus. As a result, the electron in the p orbital of boron experiences less effective nuclear charge and is less tightly held.

Due to this lower effective nuclear charge, it is easier to remove the 2p electron from boron compared to the 2s electrons from lithium and beryllium. As a result, boron has a lower first ionization energy than what the general trend would suggest.

So, the presence of the electron in a p orbital in boron allows it to deviate from the ionization energy trend across the period and have a lower first ionization energy.