Question

Compare the elements \(\mathrm{B}, \mathrm{Al}, \mathrm{C},\) and \(\mathrm{Si}\) (a) Which has the most metallic character? (b) Which has the largest atomic radius? (c) Which has the most negative electron attachment enthalpy? (d) Place the three elements \(\mathrm{B}, \mathrm{Al}\), and \(\mathrm{C}\) in order of increasing first ionization energy.

Short answer

(a) Al, (b) Al, (c) C; Order: Al < B < C.

Step-by-step solution

Determine Metallic Character

In general, metallic character increases as you move down a group in the periodic table and decreases from left to right across a period. Given elements: B (Boron), C (Carbon), Si (Silicon), Al (Aluminum). Among these, Aluminum is below Boron in the periodic table and therefore, has the most metallic character.

Compare Atomic Radii

The atomic radius increases as you move down a group and decreases across a period from left to right. Comparing: Boron (group 13, period 2), Aluminum (group 13, period 3), Carbon (group 14, period 2), Silicon (group 14, period 3). Since Aluminum is in the same group as Boron but in a higher period, it has the largest atomic radius.

Determine Electron Affinity

Electron affinity trends become more negative from left to right across a period and less negative as you move down a group. Carbon and Silicon are non-metals and generally have more negative electron attachment enthalpy compared to metals. Between Carbon and Silicon, Carbon's electron affinity is more negative. Therefore, Carbon has the most negative electron attachment enthalpy.

Order by First Ionization Energy

Ionization energy generally increases across a period and decreases down a group. For Boron, Aluminum, and Carbon: - Boron's position relative to Aluminum and Carbon means it has higher ionization energy than Aluminum but lower than Carbon. - Arrange them: Aluminum < Boron < Carbon.

Key concepts

Metallic Character

The metallic character of an element describes its ability to lose electrons, and typically elements with higher metallic character are better conductors of heat and electricity. On the periodic table, the metallic character increases as you move from right to left across a period and as you move down a group.

This is because, as you move down a group, the atomic size increases due to the addition of more electron shells. This added distance makes it easier for these outer electrons to be lost, enhancing metallic character.

In our case, Aluminum (\( \text{Al} \)) is below Boron (\( \text{B} \)) in Group 13 of the periodic table, making Aluminum more metallic since it can lose electrons more easily compared to Boron, which is higher up in the group.

Atomic Radius

Atomic radius refers to the size of an atom, measured from the center of the nucleus to the boundary of the surrounding cloud of electrons. On the periodic table, the atomic radius increases as you move down a group and decreases as you move across a period from left to right.

This increasing trend down a group occurs because each row adds a new electron shell, making the atom larger. However, atomic radii decrease across a period because additional protons in the nucleus exert a stronger attraction on the electrons, pulling them closer.

In our comparison, Aluminum (\( \text{Al} \)) has the largest atomic radius because it is located in the third period and the same group below Boron (\( \text{B} \)), indicating it has an extra shell of electrons.

Electron Affinity

Electron affinity describes an atom's tendency to accept an electron. It is often greater (more negative) in non-metals, and electron affinity generally becomes more negative moving from left to right across a period and less negative as you move down a group.

The more negative the electron affinity, the more energy is released when the atom gains an electron, indicating a stronger attraction. Generally, non-metals have higher (more negative) electron affinities compared to metals as they close to completing their valence shells.

Between Carbon (\( \text{C} \)) and Silicon (\( \text{Si} \)), Carbon has the more negative electron affinity. This is because Carbon is higher in the periodic table, implying a stronger effective nuclear charge which enables it to attract additional electrons more effectively than Silicon.

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in the gaseous state. It is an indicator of how strongly an atom holds onto its electrons. Ionization energy increases across a period and decreases down a group.

As you go across a period, increased nuclear charge leads to a stronger force on the electrons, requiring more energy to remove one. Conversely, as you move down a group, additional electron shells reduce the nucleus’s hold on the outer electrons, decreasing ionization energy.

In our exercise, Boron (\( \text{B} \)) has a higher ionization energy than Aluminum (\( \text{Al} \)) but lower than Carbon (\( \text{C} \)). Thus, ordered by increasing ionization energy, the sequence is Aluminum < Boron < Carbon.