Question

Refer to the periodic table and predict which element in each of the following pairs has the lower ionization energy: (a) \(\mathrm{Rb}\) or \(\mathrm{Cs}\) (b) He or Ar (c) \(\mathrm{B}\) or \(\mathrm{Al}\) (d) For I

Short answer

(a) Cs, (b) Ar, (c) Al, (d) Evaluated in context of its group, I has a relatively low ionization energy.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. Generally, ionization energy increases as you move across a period (left to right) and decreases as you move down a group (top to bottom) on the periodic table.

Compare Rb and Cs

Rubidium (Rb) and Cesium (Cs) are both in Group 1 of the periodic table. Since Cs is below Rb in the same group, it has a lower ionization energy due to its larger atomic size and greater shielding effect, making it easier to remove an outer electron.

Compare He and Ar

Helium (He) and Argon (Ar) are both noble gases. However, He is at the top of Group 18, while Ar is below it. Generally, ionization energy decreases as you move down a group, but because He is so small and tightly holds its electrons, it actually has a higher ionization energy than Ar.

Compare B and Al

Boron (B) and Aluminum (Al) are in the same group (Group 13), with B above Al. Since ionization energy decreases down a group, Al has a lower ionization energy than B.

Evaluate Iodine

Iodine (I) is part of Group 17 (the halogens). To estimate its relative ionization energy, note its position in the periodic table and compare it generally to elements around it. Across a period, ionization energy typically increases, and down a group, it decreases. Iodine should have a relatively low ionization energy compared to other elements in its period but higher than those below it in its group.

Key concepts

Periodic Table

The periodic table is a comprehensive map of all known elements, organized in such a way that it reflects recurring trends in their properties. It's essentially a grid where elements are positioned by their atomic number (the number of protons in an atom’s nucleus), electronic configuration, and recurring chemical properties. Each row of the table is called a 'period', and each column is referred to as a 'group'. Understanding the structure of the periodic table is essential for predicting the characteristics of elements, such as ionization energy. Elements within the same group (i.e., vertically aligned) share similar traits, as they have the same number of electrons in their outermost shell. This is a key reason why the periodic table is so useful, offering a visual reference for many fundamental chemical properties and trends.

Group Trends

Group trends in the periodic table refer to the recurring trends in properties that can be observed as you move down a group or column. One noteworthy trend is the decrease in ionization energy as you move from top to bottom within a group. This happens because:

• Atomic size increases as new electron shells are added.
• The outermost electrons are further away from the nucleus.
• There's an increase in the shielding effect by inner shells, reducing the effective nuclear charge experienced by the outermost electrons.

These factors make it easier to remove an electron as you move down a group, which is why elements such as Cesium (Cs) have lower ionization energies than those like Rubidium (Rb) above them in the same group.

Period Trends

Period trends are patterns that can be observed across a period or row in the periodic table. A significant trend is the increase in ionization energy as you move from left to right across a period. The increase occurs due to:

• A greater nuclear charge with each successive element, as more protons are added.
• A decrease in atomic size due to greater attraction between electrons and the nucleus.
• No increase in shielding effect, as electrons are being added to the same shell.

This means that, within a period, elements like Helium (He) have much higher ionization energies than those like Argon (Ar), which fit the general rule despite being further into the noble gas group.

Atomic Size and Shielding Effect

Atomic size refers to the distance from the center of the nucleus to the outermost shell of an electron cloud. As you move down a group in the periodic table, atomic size increases; conversely, as you move across a period from left to right, atomic size tends to decrease. The "shielding effect" describes how inner electrons can block the pull of the nucleus on outer electrons. This effect plays a crucial role in determining ionization energy:

• Increased shielding generally leads to lower ionization energy, as outer electrons are less tightly bound.
• Conversely, less shielding means outer electrons feel a stronger pull from the nucleus, increasing ionization energy.

Thus, in larger atoms like those found further down the periodic table, increased shielding reduces the nucleus' grip on the outermost electrons, thereby lowering the ionization energy.