Unit 3 Test Periodic Table

Carbon (C) has the smallest atomic radius among the given elements due to its higher nuclear charge and greater effective nuclear attraction on its electrons, leading to a more compact atomic structure compared to Lithium, Beryllium, and Boron.

As you move down a group in the periodic table, additional electron shells are added, which increases the distance between the nucleus and the outermost electrons. This results in an increase in atomic radius.

Oxygen (O) is the most electronegative element among the choices, with an electronegativity value of 3.44 on the Pauling scale, compared to hydrogen (2.20), carbon (2.55), and nitrogen (3.04).

Ionization energy increases across a period because the nuclear charge increases, attracting electrons more strongly. This makes it harder to remove an electron, thus requiring more energy.

Electronegativity decreases down a group because electrons are added to higher energy levels, which increases their distance from the nucleus. This greater distance reduces the nucleus's ability to attract bonding electrons.

As you move down a group in the periodic table, the ionization energy generally decreases. This is due to the increased distance of the outer electrons from the nucleus and increased electron shielding, making them easier to remove.

Noble gases have completely filled valence shells, which makes them stable and unreactive. This is in contrast to other elements, which often have partially filled shells and are more likely to react to achieve stability.

Carbon (C) has the highest first ionization energy among the given elements due to its higher nuclear charge and smaller atomic radius, which leads to a stronger attraction between the nucleus and the outermost electron.

Xenon (Xe) has the largest atomic radius among the given elements. Atomic radius increases down a group in the periodic table due to the addition of electron shells, making Xe larger than He, Ne, and Ar.

Alkali metals have a single electron in their outer shell, making it easy to remove this electron. This results in low ionization energies, as less energy is required to overcome the attraction between the electron and the nucleus.