Question

Refer to the periodic table and predict which element in each of the following pairs has the lower ionization energy: (a) \(\mathrm{Mg}\) or \(\mathrm{Si}\) (b) \(\mathrm{Pb}\) or \(\mathrm{Bi}\) (c) Ca or Ga (d) \(P\) or \(C l\)

Short answer

(a) Mg, (b) Pb, (c) Ca, (d) P

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 decreases as you move down a group in the periodic table (due to increased distance of the outer electrons from the nucleus) and increases across a period from left to right (due to increased nuclear charge).

Comparing Mg and Si

Both magnesium (Mg) and silicon (Si) are in the same period, with Mg to the left of Si. According to periodic trends, as you move across a period from left to right, ionization energy generally increases. Therefore, Mg has a lower ionization energy than Si.

Comparing Pb and Bi

Lead (Pb) is in Group 14, while bismuth (Bi) is in Group 15, but both are in the same period. As you move across a period from left to right across these groups, ionization energy increases slightly due to increased nuclear charge. Thus, Pb has a lower ionization energy than Bi.

Comparing Ca and Ga

Calcium (Ca) and gallium (Ga) are in the same period with Ca to the left of Ga. As ionization energy increases across a period, Ca will have a lower ionization energy compared to Ga.

Comparing P and Cl

Phosphorus (P) and chlorine (Cl) are found in the same period, with P located to the left of Cl. Since ionization energy increases across a period, P has lower ionization energy than Cl.

Key concepts

Periodic Table Trends

The periodic table is a powerful tool that helps chemists predict how different elements will behave. One of the core trends within the periodic table is the change in ionization energy — the energy needed to remove an electron from an atom. As you move across a period from left to right, ionization energy generally increases. This is because the nuclear charge (the positive charge of the nucleus) increases, pulling electrons closer and making them harder to remove.
On the other hand, as you move down a group in the periodic table, ionization energy usually decreases. This occurs because the electrons are further from the nucleus due to the addition of electron shells, which reduces the nuclear pull on the outermost electrons. So, the further down the group an element is, the easier it is to remove an electron.
To summarize, keep these points in mind:

• Ionization energy increases across a period (left to right)
• Ionization energy decreases down a group (from top to bottom)

Ionization Energy Comparison

Understanding ionization energy is critical when comparing the reactivity of different elements. In the exercise examples, we are comparing pairs of elements to determine which has the lower ionization energy.
First, consider magnesium (Mg) and silicon (Si). They are in the same period, with Mg to the left of Si. Since ionization energy increases across a period, Mg will have a lower ionization energy than Si.
Next, look at lead (Pb) and bismuth (Bi). Similarly, these elements are in the same period, with Pb to the left of Bi. Therefore, Pb has a lower ionization energy than Bi due to the increasing nuclear charge as you move to the right.
In the case of calcium (Ca) and gallium (Ga), both are in the same period, but once again, Ca is positioned to the left of Ga. As expected, Ca has a lower ionization energy.
Lastly, phosphorus (P) and chlorine (Cl) show the same trend. P, being to the left of Cl, has a lower ionization energy.
Each example underscores the principle that moving from left to right in a period increases ionization energy, effectively illustrating how this vital periodic trend functions in practical comparisons.

Electron Removal

Removing an electron from an atom can tell us much about an element's chemical behavior. This process, known as ionization, requires energy because electrons are attracted to the positively charged nucleus. Understanding which elements have lower ionization energies can help predict their reactivities.
Electrons that are easiest to remove are typically those farthest from the nucleus. In elements lower on the periodic table or farther to the left, outer electrons are less tightly held, making them easier to remove, hence exhibiting lower ionization energy.
The exercise shows how the concept of electron removal helps us determine which element in a pair has a lower ionization energy by exploiting these periodic table understandings. The trends imply that elements positioned further down or to the left in a period are more apt to lose electrons, making them generally more reactive or more readily ionizable compared to their counterparts.
Thus, knowing about electron removal processes not only clears up concepts of ionization energy but also broadens the comprehension of periodic trends and atomic reactivity.
In summary, a grasp of electron removal aids in understanding and predicting the chemical characteristics of elements.