Question

Write the symbol of each element described below. (a) largest atomic radius in Group 17 (b) smallest atomic radius in period 3 (c) largest first ionization energy in Group 2 (d) most electronegative in Group 16 (e) element(s) in period 2 with no unpaired p electron (f) abbreviated electron configuration is \([\mathrm{Ar}] 4 \mathrm{~s}^{2} 3 \mathrm{~d}^{3}\) (g) A \(+2\) ion with abbreviated electron configuration [Ar] \(3 \mathrm{~d}^{5}\) (h) A transition metal in period 4 forming \(\mathrm{a}+2\) ion with no unpaired electrons

Short answer

(a) The element with the largest atomic radius in Group 17 is Francium (Fr). (b) The element with the smallest atomic radius in period 3 is Argon (Ar). (c) The element with the largest first ionization energy in Group 2 is Beryllium (Be). (d) The most electronegative element in Group 16 is Oxygen (O). (e) The element in period 2 with no unpaired p electron is Neon (Ne). (f) The element with an abbreviated electron configuration of [Ar]4s²3d³ is Vanadium (V). (g) The ion with an abbreviated electron configuration of [Ar]3d⁵ and a +2 charge is Chromium ion (Cr²⁺). (h) A transition metal in period 4 forming a +2 ion with no unpaired electrons is Zinc ion (Zn²⁺).

Step-by-step solution

(a) Largest atomic radius in Group 17

To find the element with the largest atomic radius in Group 17, we must remember that atomic radius increases as we go down the group. Therefore, the element with the largest atomic radius in Group 17 will be the one at the bottom of the group. This element is Francium (symbol: Fr).

(b) Smallest atomic radius in period 3

To find the element with the smallest atomic radius in period 3, we must remember that atomic radius decreases as we move to the right across a period due to increasing effective nuclear charge. The element with the smallest atomic radius in period 3 is the one on the rightmost side, the element Argon (symbol: Ar).

(c) Largest first ionization energy in Group 2

To find the element with the largest first ionization energy in Group 2, we must remember that ionization energy increases as we go up a group because the outer electrons are closer to the nucleus and are more strongly attracted. Therefore, the element with the largest first ionization energy in Group 2 is Beryllium (symbol: Be).

(d) Most electronegative in Group 16

To find the most electronegative element in Group 16, we must remember that electronegativity increases as we go up a group and to the right across a period. In Group 16, the most electronegative element is the top one, Oxygen (symbol: O).

(e) Element(s) in period 2 with no unpaired p electron

To find the element(s) in period 2 with no unpaired p electrons, we look for elements in period 2 with a full p subshell. In period 2, the element Neon (symbol: Ne) has a full p subshell (2p6) and thus has no unpaired p electrons.

(f) Abbreviated electron configuration [Ar]4s²3d³

The element with an abbreviated electron configuration of [Ar]4s²3d³ has two 4s electrons and three 3d electrons after the Argon configuration. This configuration corresponds to the element Vanadium (symbol: V).

(g) A +2 ion with abbreviated electron configuration [Ar]3d⁵

If a +2 ion has an abbreviated electron configuration of [Ar]3d⁵, the neutral atom must have two more electrons than the ion. The neutral atom's electron configuration would then be [Ar]4s²3d⁵. This corresponds to the element Chromium (symbol: Cr). Therefore, the ion is Cr²⁺.

(h) A transition metal in period 4 forming a +2 ion with no unpaired electrons

For a transition metal in period 4 to form a +2 ion with no unpaired electrons, the +2 ion must have electron configuration with all of its d and s orbitals fully paired. One such element is Zinc (symbol: Zn) with electron configuration of [Ar]4s²3d¹⁰ as a neutral atom and [Ar]3d^10 after losing two electrons to form the Zn²⁺ ion.

Key concepts

Atomic Radius

Understanding atomic radius is crucial when analyzing chemical elements. It describes the size of an atom, typically measured from the nucleus to the boundary where the electron cloud ends. As we move down a group in the periodic table, the atomic radius increases because each subsequent element has an additional electron shell. Conversely, as we move across a period from left to right, the atomic radius decreases. This is due to the increasing nuclear charge, which pulls electrons closer to the nucleus, making the atom smaller.

For instance, in our problem's scenario (a) and (b), Francium (Fr) has the largest atomic radius in Group 17 because it's at the bottom of the group, whereas Argon (Ar) has the smallest atomic radius in period 3 due to its position on the far right of the period, affected by the greatest effective nuclear charge among its peers.

First Ionization Energy

First ionization energy refers to the amount of energy required to remove the most loosely bound electron—known as the valence electron—from a neutral atom in its gaseous state. The trend for the first ionization energy is opposite to that of atomic radius: it generally increases as you move up a group and across a period from left to right. This is because atoms with smaller sizes have their electrons closer to the nucleus, making it more difficult to remove them.

In step 3 of our solution, Beryllium (Be) in Group 2 has the largest first ionization energy, which correlates with its position at the top of the group where electrons are closer to the nucleus and more strongly attracted.

Electronegativity

Electronegativity measures an atom's tendency to attract and bond with electrons. It's a dimensionless quantity that reflects how strongly atoms pull on electrons in a chemical bond. In general, electronegativity increases as one moves from left to right across a period and from bottom to top within a group on the periodic table. This is because atoms with higher nuclear charges effectively attract bonding electrons and smaller atoms can bring these electrons closer to the nucleus.

For example, as indicated in step 4, Oxygen (O) is the most electronegative element in Group 16, conforming to its position at the top of the Group with higher electronegativity compared to its counterparts lower down.

Electron Configurations

An electron configuration is a distribution of electrons of an atom or molecule in atomic or molecular orbitals. It provides insight into the chemical behavior and bonding of an element. The electron configuration is built following the Aufbau principle, which states that electrons occupy orbitals of the lowest energy first. The notations such as 1s² or 3d⁵ represent the energy level, type of orbital, and the number of electrons in those orbitals, respectively.

As observed in steps (e) through (h), Neon has a full 2p6 configuration, implying filled p orbitals and no unpaired electrons. Vanadium features a 4s²3d³ configuration. For Chromium and Zinc, the electron configurations ([Ar]4s²3d⁵ for Cr and [Ar]4s²3d¹⁰ for Zn) are crucial in understanding their chemical properties and the nature of their ions.

Transition Metals

Transition metals include elements found in Groups 3 through 12 of the periodic table. These elements are characterized by their d orbitals and typically have valence electrons in both their s and d shells. Transition metals are known for their ability to form a variety of different ions with different charges, their colorful compounds, and their use as catalysts in numerous chemical reactions. Importantly, many of these metals can exhibit different oxidation states due to the loss of d-electron(s) as well as s-electron(s).

In the final part of our exercise, we encounter Zinc (Zn), a transition metal from period 4, which, in its +2 ion form, has a full d shell with no unpaired electrons. This is a characteristic example of how transition metals can form stable ions with distinct electronic structures.