As you move from left to right across a period, electrons are added to the same energy level while the nucleus gains more protons. The increased effective nuclear charge pulls the electron cloud closer, resulting in a decrease in atomic radius.

Ionization energy is defined as the energy required to remove an electron from a neutral atom in its gaseous state. This concept is important for understanding an element's reactivity and placement in the periodic table.

Electronegativity generally increases as you move from left to right across a period because atoms have a higher effective nuclear charge and a smaller radius. This enhances their ability to attract electrons in a chemical bond.

Down a group, each successive element has an additional electron shell compared to the previous one. The increase in the number of electron shells outweighs the increase in nuclear charge, leading to a larger atomic radius.

Electronegativity quantifies how strongly an atom attracts electrons within a chemical bond. This property varies across the periodic table and influences molecular structure and reactivity.

Neon is a noble gas with a completely filled valence shell, making it stable and less willing to lose an electron. This stability accounts for its high ionization energy compared to the metals listed.

Potassium is located in a lower period with more electron shells compared to the other elements listed. The additional shells significantly increase the atomic radius, making it the largest among the options.

As you move across a period, the number of protons in the nucleus increases, resulting in a stronger effective nuclear charge. This enhanced pull draws electrons closer, increasing the electronegativity of the element.

Electron shielding reduces the net attractive force exerted by the nucleus on the outer electrons. With a diminished effective nuclear charge, less energy is required to remove an electron, leading to a decrease in ionization energy.

Across a period, atoms tend to have a greater tendency to attract electrons, which results in more negative electron affinity values. This trend is due to the increased effective nuclear charge and decreased atomic size.

A filled valence shell indicates a stable electronic configuration that resists the removal of electrons. Therefore, more energy is required to ionize such an atom, resulting in a higher ionization energy.

The effective nuclear charge is largely determined by the total positive charge of the nucleus, which is a direct result of the number of protons. Although electron shielding moderates this effect, the proton count remains the primary factor.

As you progress from left to right across a period, the increased effective nuclear charge contracts the electron cloud. This results in a consistent decrease in atomic radius.

The addition of electron shells as you move down a group causes the outer electrons to be farther from the nucleus. This leads directly to an increase in the atomic radius.

Although the nuclear charge increases down a group, the effect of additional electron shells creates significant electron shielding. This shielding reduces the nucleus's pull on bonding electrons, leading to a decrease in electronegativity.

Beryllium exhibits a stable, filled 2s orbital configuration that tightly holds its electrons. In contrast, boron adds an electron to the 2p orbital, which is higher in energy and less effectively shielded, thereby reducing its ionization energy relative to beryllium.

Electronegativity is a relative scale that describes how strongly an atom attracts bonding electrons, whereas electron affinity is an absolute energy change when an electron is added. Differences in subshell configurations and electron-electron repulsions mean these two trends do not mirror each other exactly.

In transition metals, the electrons in d-orbitals shield the outer electrons less effectively than s- or p-electrons. This additional shielding introduces irregularities in trends such as atomic radius and ionization energy compared to the predictable patterns of main-group elements.

As elements descend a group, the addition of electron shells increases the electron shielding effect. This shielding reduces the hold of the nucleus on the outer electrons, thereby lowering the ionization energy.

A higher effective nuclear charge means that electrons are drawn more tightly toward the nucleus. This results in a smaller atomic radius and requires more energy to remove an electron, thereby increasing the ionization energy.