Hypothesis 2, including explanation, for reactivity trend down a column:
Metal reactivity will increase as you go down a group, because the nuclear shielding increases and causes the nucleus’ hold on the valence electron to weaken, therefore it is easier to remove the valence electrons.
Is this right to you? It is what I came up with, I just wanted it checked. Anything added, etc.
That sounds OK to me but I think of it as the shielding effect coupled with the distance. The outer electrons is farther and farther out as one goes down the table for group I metals. That increasing distance means the attractive forces between the nucleus and the lone electron in the outer shell is weaker and weaker so it is easier to pull that last electron away. The shielding effect of the inner electrons means less attraction between the nucleus and the outer electron, too, so both contribute to removing the outermost electron.
Your hypothesis is almost correct, but there is an important addition that you can make to strengthen your explanation.
Metal reactivity does increase as you go down a group because of the increase in atomic size due to additional electron shells. This increase in atomic size leads to greater distance between the positively charged nucleus and the valence electrons. As a result, the attractive forces between the nucleus and valence electrons weaken, making it easier for the valence electrons to be removed.
Additionally, the increase in atomic size also results in an increase in the number of inner electron shells, which provides increased shielding for the valence electrons from the positively charged nucleus. This increased shielding further weakens the attractive forces between the nucleus and valence electrons.
So, in summary, as you move down a group, the combination of increased atomic size and increased shielding of valence electrons makes it easier for metals to lose their valence electrons, leading to an increase in their reactivity.
Your hypothesis is partly correct. The reactivity trend down a column in the periodic table generally follows the pattern of increasing reactivity. This is due to the increase in atomic size and the decreasing effective nuclear charge as you move down the group.
Here's an explanation of why the reactivity increases:
1. Atomic Size: As you go down a group, the atomic size increases because of the addition of new electron shells. The increase in size means the valence electrons are further away from the nucleus. Consequently, the attractive force between the nucleus and valence electrons decreases, making it easier for the valence electrons to be involved in chemical reactions.
2. Effective Nuclear Charge: The effective nuclear charge experienced by the valence electrons also decreases as you move down a group. Effective nuclear charge refers to the positive charge experienced by the valence electrons, after taking into account the shielding effect of inner electrons. As the number of electron shells increases down the group, there is greater electron-electron repulsion and increased shielding by the inner shells. This reduces the effective nuclear charge on the valence electrons, weakening the attraction to the nucleus and making it easier for them to be involved in chemical reactions.
Considering these two factors, the hypothesis that the metal reactivity increases as you go down a group is supported. However, it's important to note that this trend may not be true for all elements and there can be exceptions based on specific elements or groups. Nonetheless, the general trend of increasing reactivity down a column holds in most cases.