Question

Based on their positions in the periodic table, predict which atom of the following pairs will have the smaller first ionization energy: (a) \(\mathrm{Cl}, \mathrm{Ar} ;\) (b) \(\mathrm{Be}, \mathrm{Ca}\) (c) \(\mathrm{K}\), Co; (d) \(\mathrm{S}, \mathrm{Ge} ;\) (e) Sn. Te.

Short answer

Considering the periodic trends in first ionization energy, we can predict the following: (a) Cl will have a smaller first ionization energy than Ar. (b) Ca will have a smaller first ionization energy than Be. (c) K will have a smaller first ionization energy than Co. (d) S will have a smaller first ionization energy than Ge. (e) Sn will have a smaller first ionization energy than Te.

Step-by-step solution

a) Cl and Ar

First, locate Cl and Ar in the periodic table. Cl is in Group 17 and Period 3, while Ar is in Group 18 and Period 3. Since both elements share the same Period, they are in the same horizontal row. Recall that ionization energy typically increases from left to right across a period. Given that Cl is to the left of Ar, it is expected to have the smaller first ionization energy.

b) Be and Ca

Locate Be and Ca in the periodic table. Be is in Group 2 and Period 2, while Ca is in Group 2 and Period 4. Both elements are in the same Group and therefore in the same vertical column. Ionization energy typically decreases from top to bottom within a group. As Be is higher up in the periodic table than Ca, Ca should have the smaller first ionization energy.

c) K and Co

Locate K and Co in the periodic table. K is in Group 1 and Period 4, while Co is in Group 9 and Period 4. Both elements are in the same Period and so are in the same horizontal row. As mentioned earlier, ionization energy increases from left to right across a period. Due to K being located more to the left compared to Co, K is predicted to have the smaller first ionization energy.

d) S and Ge

Locate S and Ge in the periodic table. S is in Group 16 and Period 3, while Ge is in Group 14 and Period 4. Neither elements share a Period or Group, so we need to consider both periodic trends. Since ionization energy increases across a period and decreases down a group, the smaller ionization energy should be in an element that is more to the left and/or lower down on the periodic table. Ge is one place lower, but three places to the right of S. Therefore, S is predicted to have the smaller first ionization energy.

e) Sn and Te

Locate Sn and Te in the periodic table. Sn is in Group 14 and Period 5, while Te is in Group 16 and Period 5. Both elements are in the same Period, which means they are in the same horizontal row. Since ionization energy increases from left to right across a period, Sn, which is located to the left of Te, is predicted to have the smaller first ionization energy.

Key concepts

Periodic Table Trends

Understanding the periodic table is key to mastering concepts in chemistry. It's not just a chart of elements, but a map of atomic behavior and reactivity. The periodic table trends are critical in predicting how elements will interact, such as their ionization energies.

One of the essential trends is the variation of ionization energy across periods (horizontal rows) and down groups (vertical columns). Ionization energy generally increases as one moves from left to right across a period. This is because elements gain more protons and electrons, increasing the attraction between the positively charged nucleus and the negatively charged electrons. Consequently, more energy is required to remove an electron.

In contrast, ionization energy typically decreases moving down a group as the outer electrons are located in increasingly distant orbitals from the nucleus, which weakens their electrostatic attraction to it, and therefore, less energy is needed to remove them. These principles allow us to predict relative ionization energies simply upon an element's position in the periodic table.

Atomic Structure

The atomic structure is a foundational concept in chemistry, explaining the configuration of protons, neutrons, and electrons in an atom. Central to this structure is the atomic nucleus, consisting of protons and neutrons, around which electrons orbit in defined regions of space called electron shells.

The number of protons (atomic number) determines the identity of the element, while the distribution of electrons across the shells influences how an atom will interact with others. These electron arrangements are known as electron configurations and play a vital role in properties such as ionization energy. The closer an electron is to the nucleus, the stronger the electrostatic forces holding it, making atoms with tightly packed inner electrons generally have higher ionization energies.

First Ionization Energy

The first ionization energy is a significant concept explaining the energy required to remove the most loosely held electron from an atom in its gaseous state to form a cation. This 'first' pertains to the energy required to remove the outermost electron, which is significant as it indicates the strength of the electron's binding to the atom.

Various factors affect this value, such as electron shielding, atomic radius, and nuclear charge. For example, the more shielded an outer electron is by inner electrons, the lower the ionization energy. Likewise, a larger atomic radius typically means lower ionization energy because the outer electrons are further away from the influence of the nucleus. Understanding this energy is useful in predicting chemical reactivity and periodic trends.

Electron Configuration

Electron configuration is the arrangement of electrons in an atom's shells and subshells. It follows a set of rules, such as the Pauli Exclusion Principle and Hund's Rule, to minimize repulsion and maximize stability.

The configuration deeply influences an atom's chemical properties, including ionization energy. Atoms strive for the most stable electron configuration, often resembling that of the nearest noble gas with a complete valence shell. As you progress through the periodic table, the electron configuration patterns help to explain why elements exhibit distinct periodic trends in their ionization energies. Elements with a complete or half-filled valence shell often possess higher ionization energies due to their relatively stable electron arrangements. This stability must be overcome by energy input to remove an electron, thus explaining the trend in ionization energy seen across different elements.