Question

In each of the following sets of elements, which element would be expected to have the highest ion ization energy? a. \(\mathrm{Cs}, \mathrm{K}, \mathrm{Li}\) b. Ba, Sr, Ca c. I, Br, Cl d. \(\mathrm{Mg}, \mathrm{Si}, \mathrm{S}\)

Short answer

The elements with the highest ionization energy in each set are: a. Li b. Ca c. Cl d. S

Step-by-step solution

Set a: Cs, K, Li

In this set, all elements belong to group 1 (Alkali metals). Since ionization energy decreases down a group, the element with the highest ionization energy will be the element at the top of this group. Between Cs, K, and Li, lithium (Li) is the highest, so it has the highest ionization energy.

Set b: Ba, Sr, Ca

In this set, all elements belong to group 2 (Alkaline earth metals). Similar to set a, we compare their positions in the periodic table within the same group. Between Ba, Sr, and Ca, calcium (Ca) is the highest in group 2, so it has the highest ionization energy.

Set c: I, Br, Cl

In this set, all elements belong to group 17 (Halogens). These elements are in the same period, so ionization energy will increase as we move across the period from left to right. Between I, Br, and Cl, chlorine (Cl) is at the rightmost position in this period, so it has the highest ionization energy.

Set d: Mg, Si, S

In this set, all elements are in the same period (period 3). As mentioned earlier, ionization energy increases across a period from left to right. Between Mg, Si, and S, sulfur (S) is at the rightmost position in this period, so it has the highest ionization energy.

Key concepts

Periodic Table Trends

Understanding the periodic table is crucial for grasping why certain elements exhibit specific properties. One of the vital trends to understand is the ionization energy, which is the energy required to remove an electron from an atom in its gaseous state. Ionization energy generally increases across a period (from left to right) as the nuclear charge of the atoms increases. Simultaneously, it decreases down a group because the electrons are further from the nucleus and shielding effects become stronger.

These trends are a result of the balance between the attractive force of the atomic nucleus and the repulsive forces between electrons. Thus, elements in the top right of the periodic table (apart from noble gases) have the highest ionization energies due to their large nuclear charge and smaller atomic radii. In contrast, those in the bottom left have the lowest ionization energies. Recognizing these trends empowers students to predict the behavior of elements during chemical reactions.

Alkali Metals

Alkali metals, residing in group 1 of the periodic table, are known for having low ionization energies, which make them highly reactive. This group includes lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). As we descend the group, the ionization energy decreases because of the increasing distance of the valence electron from the nucleus and the increased effect of electron shielding.

Key Properties of Alkali Metals:

• They have a single electron in their outermost shell, which they lose easily, making them highly reactive.
• As the atomic number increases down the group, the reactivity increases due to lower ionization energy.
• They are characterized by their softness and low melting and boiling points relative to the rest of the metals.

Understanding these properties helps in explaining why lithium, although an alkali metal, will have a higher ionization energy compared to cesium or potassium.

Alkaline Earth Metals

Alkaline earth metals are found in group 2 of the periodic table and include beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). These elements have two electrons in their outermost shell and exhibit slightly higher ionization energies compared to alkali metals, attributable to their fully filled s-orbitals.

Common Characteristics of Alkaline Earth Metals:

• Their two valence electrons make them less reactive than alkali metals, yet they too readily lose electrons to form cations with a +2 charge.
• Ionization energy decreases from Be to Ra within the group, following the general trend seen across all groups in the periodic table.
• They typically have higher melting points than alkali metals and are harder and denser.

By remembering these characteristics, one can understand why calcium has a higher ionization energy than strontium and barium, following the trend for this group of elements.

Halogens

Halogens occupy group 17 of the periodic table, consisting of fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). These nonmetals are notorious for their high ionization energies, which decrease down the group. Their high electronegativities and small atomic radii contribute to this trait. Halogens are one electron short of a full valence shell, which makes them highly reactive, especially towards alkali and alkaline earth metals.

Distinguishing Features of Halogens:

• All halogens have seven electrons in their outer shells, making them eager to gain an electron and form a -1 anion.
• Their physical states vary at room temperature, from gases (F and Cl) to liquids (Br) and solids (I and At).
• Due to their high ionization energies and electronegativities, halogens are excellent oxidizing agents.

We can apply this knowledge to deduce that among iodine, bromine, and chlorine, Cl has the highest ionization energy because it is the highest up in the group and closest to completing its valence electron shell.