Question

Choose the element with the higher ionization energy from each pair. (a) As or Bi (b) As or Br (c) \(\mathrm{S}\) or \(\mathrm{I}\) (d) \(S\) or \(S b\)

Short answer

The elements with higher ionization energy in each pair are: (a) As, (b) Br, (c) S, (d) S.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. Generally, ionization energy increases across a period from left to right and decreases down a group on the periodic table.

Comparing Ionization Energy of As and Bi

Arsenic (As) and Bismuth (Bi) are both in Group 15 of the periodic table but As is above Bi. Since ionization energy decreases as you move down a group, As has a higher ionization energy than Bi.

Choosing Between As and Br

Arsenic (As) is in Group 15 and Bromine (Br) is in Group 17 of the same period on the periodic table. Within a period, ionization energy increases moving to the right, thus Br has a higher ionization energy than As.

Comparing Ionization Energy of S and I

Sulfur (S) and Iodine (I) are both in Group 16, but S is above I on the periodic table. As ionization energy decreases down a group, S has a higher ionization energy than I.

Determining the Higher Ionization Energy between S and Sb

Sulfur (S) is in Group 16 and Antimony (Sb) is in Group 15. Moreover, S is one period above Sb in the table. Both these factors contribute to S having a higher ionization energy than Sb.

Key concepts

Periodic Table Trends

Understanding the trends in the periodic table is crucial for predicting the properties of elements, including their ionization energy. Ionization energy generally tends to increase as you move from left to right across a period. This is because the addition of protons in the nucleus across a period enhances the positive charge, pulling electrons closer and making them harder to remove.

Conversely, as you move down a group, ionization energy decreases. This is due to the increase in the number of electron shells, which causes a greater distance between the nucleus and the outermost electron and reduces the effective nuclear charge experienced by those electrons. Therefore, electrons are easier to remove from atoms lower down in a group.

These simple but essential trends enable a comparison of ionization energies among various elements and predict which might have a higher ionization energy in comparison exercises.

Atomic Structure

The atomic structure has a profound impact on an element's ionization energy. Every atom consists of a dense nucleus surrounded by a cloud of electrons. These electrons occupy regions called shells or energy levels, which can be thought of as the 'floors' in an 'atomic hotel'.

Electrons nearer to the nucleus feel a stronger attraction due to the positive charge of the protons and require more energy to be removed. This concept helps explain why ionization energy increases across a period - the electrons are being pulled closer as the positive charge in the nucleus increases.

Moreover, the distribution of electrons into various sublevels (s, p, d, f) also affects the ionization energy. Electrons in the same sublevel repel each other, and as more electrons are added (especially in 'p' and 'd' sublevels), they are slightly easier to remove. Therefore, understanding the principles of atomic structure aids in grasing the variations in ionization energy across different elements.

Group and Period Comparison

When comparing elements based on groups and periods in the periodic table, you're essentially looking at a vertical and horizontal perspective of ionization energy trends. Along a period (horizontal rows on the table), ionization energy increases due to more protons being added to the nucleus without adding a new electron shell, thereby increasing nuclear charge and pull on the electrons.

In a group comparison (vertical columns of the table), the elements further down a group have additional shells of electrons. These additional shells 'shield' the outermost electron from the nucleus' pull, reducing the ionization energy required to remove it.

The step-by-step solutions in the original exercise demonstrate this beautifully. For instance, arsenic (As) is higher in the group than bismuth (Bi), hence As has a higher ionization energy. Such comparisons provide learners with the conceptual groundwork to apply these trends and extract conclusions on the elements' reactivity and bonding characteristics.

Electron Removal Energy

Electron removal energy, more commonly referred to as ionization energy, is the quantifiable amount of energy required to strip away an electron from an atom in the gas phase, forming a cation. This energy is a critical indicator of how likely an atom is to participate in chemical reactions that lead to the formation of compounds.

Ionization energy is often measured in electron volts (eV) or kilojoules per mole (kJ/mol). An element with a high ionization energy is more reluctant to lose electrons and is less likely to form positive ions (cations). On the other hand, elements with low ionization energies tend to lose electrons easily and are more reactive.

This concept is essential when determining the chemical behavior of elements and can be especially helpful in understanding and predicting the formation of ionic bonds. For example, elements with significantly different ionization energies are likely to form ionic compounds, such as sodium (low ionization energy) and chlorine (high ionization energy) combining to form sodium chloride (NaCl).