Question

Refer to the periodic table and predict which element in each of the following pairs has the higher ionization energy: (a) Ga or Ge (b) Si or P (c) Br or Cl (d) As or Sb

Short answer

Ge, P, Cl, and As have higher ionization energies.

Step-by-step solution

Understand Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. Generally, it increases across a period (left to right in the periodic table) and decreases down a group (top to bottom).

Analyze the Pair (Ga or Ge)

Gallium (Ga) and Germanium (Ge) are both located in the fourth period of the periodic table, with Ge to the right of Ga. Since ionization energy increases across a period, Ge has a higher ionization energy than Ga.

Analyze the Pair (Si or P)

Silicon (Si) and Phosphorus (P) are both in the third period, with Phosphorus being to the right of Silicon. Therefore, due to the trend of increasing ionization energy across a period, P has a higher ionization energy than Si.

Analyze the Pair (Br or Cl)

Bromine (Br) and Chlorine (Cl) are in the same group (Group 17), but Cl is above Br. Ionization energy decreases down a group, so Cl, being above Br, has a higher ionization energy.

Analyze the Pair (As or Sb)

Arsenic (As) and Antimony (Sb) are both in Group 15, with As above Sb. Since ionization energy decreases as you move down the group, As has a higher ionization energy than Sb.

Key concepts

Periodic Table

The periodic table is an organized display of all known chemical elements. Arranged by increasing atomic number (the number of protons in an atom's nucleus), it reveals patterns in element properties.
Columns in the periodic table are known as groups, while rows are called periods. Elements in the same group have similar chemical properties because they have the same number of valence electrons.
Each period represents a new level of electron shell being filled. Understanding the layout of the periodic table is essential in predicting various element properties, such as ionization energy.

Trends in Ionization Energy

Ionization energy reflects how much energy is needed to remove an electron from an atom in the gaseous state.
Across a period, ionization energy typically increases. Why? Atoms gain more protons and electrons, pulling the electron closer to the nucleus, requiring more energy to remove it.
Down a group, the trend is reversed. Ionization energy decreases due to additional electron shells. This increased distance reduces the effective nuclear pull, making it easier to remove an electron.

Element Comparison

When comparing elements, a clear understanding of their position on the periodic table aids in predicting properties.
For example, Germanium ( Ge ) has a higher ionization energy than Gallium ( Ga ) because Ge is further right in the same period.
Phosphorus ( P ), also for being further right in its period, has a higher ionization energy than Silicon ( Si ).
Considering elements in the same group like Chlorine ( Cl ) and Bromine ( Br ), Cl is higher up the group, giving it a higher ionization energy.
In the case of Arsenic ( As ) and Antimony ( Sb ), As, being above Sb, also possesses a higher ionization energy.

Educational Chemistry Concepts

Understanding the foundational concepts of chemistry, such as ionization energy, demands grasping how these concepts interlink with periodic trends.
Ionization energy directly connects with electron configurations and the structure of atoms. The need to remove electrons relates to atomic radii and overall element stability.
Also, in educational chemistry, applying these concepts through comparisons and predictions, such as distinguishing higher ionization energies, helps reinforce deeper understanding, moving beyond rote memorization to analytical thinking.

• Apply knowledge of periodic trends to predict outcomes.
• Relate electron configurations to real-world element behavior.
• Use comparisons for practical learning and retention.

Understanding these core principles offers a strong foundation for further studies in chemistry.