Choose the element with the higher first ionization energy from each pair. a. P or I b. Si or Cl c. P or Sb d. Ga or Ge

Video transcript

Hey everyone in this example, we need to figure out which of the below statements are true based on periodic trends. Now, reading out all of these answer choices. They're all in regards to ionization energy. We want to recall that ionization energy. I'm abbreviating it here is represented as I With a subscript of the # one. And so our next step is to recall our periodic trend of ionization energy, which is going to be increasing as we go from the bottom towards the top and towards the right of our periodic tables. And so what we're going to want to do is since these choices are comparing magnesium to aluminum is we're going to list out our location of magnesium versus aluminum. So based on our periodic table, we would find magnesium in Group two A And across period three on the periodic table. Now, aluminum we would find Group three A. On our periodic table, so more towards the right And across period three as well. So we would say that therefore because aluminum is more towards the right and a higher group number than our magnesium atom, it follows the trend of increasing ionization energy. And so we can say that therefore aluminum's first ionization energy is going to be greater than magnesium seems first ionization energy. And so the only answer choice that corresponds with what we have concluded is stating that B magnesium has a lower first ionization energy than aluminum, which we would agree with because we determined that it's a smaller, it has a smaller value due to it being more towards the left of our periodic table, following the opposite direction of our ionization energy trend in the direction of decreasing. And so B is the correct choice to complete this example. If you have any questions, please leave them down below, and I will see everyone in the next practice video.