The question asks about the trend of atomic radius across a period (from left to right) in the periodic table. To answer this, we need to understand the factors that influence atomic radius, specifically the number of electron shells and the effective nuclear charge.
Therefore, as we move from left to right across a period, the atomic radius decreases due to an increase in effective nuclear charge.
B) Decreases due to increase in effective nuclear charge — This option correctly identifies both the trend (decrease) and the primary reason for it (increase in effective nuclear charge). As the number of protons increases across a period, the nuclear pull on the valence electrons becomes stronger, drawing them closer to the nucleus and reducing the atomic size, while the number of electron shells remains constant.
The question asks why atomic radius generally increases down a group in the periodic table. This trend is a fundamental concept in atomic structure and periodic properties. We need to analyze how atomic properties change as we move down a column (group) in the periodic table.
Correct Option: B) Principal quantum number increases and shielding by inner electrons increases. This option correctly identifies the two primary reasons for the increase in atomic radius down a group: the addition of new electron shells (increasing principal quantum number) and the increased shielding effect from inner electrons, both of which cause the valence electrons to be held less tightly and further from the nucleus.
The atomic radius is a periodic property that generally increases down a group and decreases across a period in the periodic table. All the given options (Li, Na, K, Cs) belong to Group 1 (alkali metals). Therefore, to determine the largest atomic radius, we need to identify the element that is lowest in this group.
Correct Option: D) Cs
The atomic radius is a periodic property that varies across a period and down a group in the periodic table. To determine the smallest atomic radius among the given elements, we need to understand these trends.
Correct Option: A) F
Fluorine (F) is located furthest to the right in the second period among the given options. As atomic radius decreases across a period due to increasing effective nuclear charge, Fluorine will have the smallest atomic radius.
The bond length in a diatomic molecule formed by two identical atoms is approximately twice the covalent radius of that atom. This is because the bond length represents the distance between the nuclei of the two bonded atoms.
Therefore, the bond length in a Cl\(_2\) molecule is 198 pm.
C) 198 pm — This is calculated by doubling the covalent radius, which is the correct method for finding the bond length between two identical atoms.
The question asks why the atomic radii of noble gases cannot be directly compared with the covalent radii of other elements. This requires understanding the different types of atomic radii and how they are measured.
A) Because noble gases are monoatomic and their radii are van der Waals radii — This is correct. Noble gases exist as individual atoms (monoatomic) and do not form covalent bonds under normal conditions. Therefore, their atomic size is measured using van der Waals radii, which are based on non-bonded interactions. Covalent radii, on the other hand, are based on bonded interactions. Since these are fundamentally different types of measurements, they are not directly comparable.