Questions: Why are the first ionization energies of some of the group 13 elements smaller than the first ionization energies of the nearby alkali earth metals in the same period? Choose one: It is easier to remove an electron in a p orbital than an s orbital because it experiences less effective nuclear charge. There is an increase in shielding by inner-shell electrons and an increase in atomic size. It is easier to remove an electron in an s orbital than a p orbital because it experiences less effective nuclear charge. There is a decrease in shielding by inner-shell electrons and a decrease in atomic size. Why are the first ionization energies of some of the group 16 elements smaller than the first ionization energies of the nearby group 15 elements in the same period?

Why are the first ionization energies of some of the group 13 elements smaller than the first ionization energies of the nearby alkali earth metals in the same period?

Choose one:
It is easier to remove an electron in a p orbital than an s orbital because it experiences less effective nuclear charge.
There is an increase in shielding by inner-shell electrons and an increase in atomic size.
It is easier to remove an electron in an s orbital than a p orbital because it experiences less effective nuclear charge.
There is a decrease in shielding by inner-shell electrons and a decrease in atomic size.

Why are the first ionization energies of some of the group 16 elements smaller than the first ionization energies of the nearby group 15 elements in the same period?

Transcript text: Why are the first ionization energies of some of the group 13 elements smaller than the first ionization energies of the nearby alkali earth metals in the same period? Choose one: It is easier to remove an electron in a $p$ orbital than an sorbital because it experiences less effective nuclear charge. There is an increase in shielding by inner-shell electrons and an increase in atomic size. It is easier to remove an electron in an s orbital than a porbital because it experiences less effective nuclear charge. There is a decrease in shielding by inner-shell electrons and a decrease in atomic size. Part 3 (1 point) Why are the first ionization energies of some of the group 16 elements smaller than the first ionization energies of the nearby group 15 elements in the same period?

Solution

Solution Steps

Step 1: Understanding Ionization Energy Trends

Ionization energy is the energy required to remove an electron from an atom. Generally, ionization energy increases across a period due to increasing nuclear charge, which pulls electrons closer to the nucleus.

Step 2: Comparing Group 13 and Alkali Earth Metals

Group 13 elements have their outermost electron in a $p$ orbital, while alkali earth metals have their outermost electron in an $s$ orbital. The $p$ orbitals are higher in energy and less tightly bound than $s$ orbitals, making it easier to remove an electron from a $p$ orbital.

Step 3: Analyzing the Options

Option 1: It is easier to remove an electron in a $p$ orbital than an $s$ orbital because it experiences less effective nuclear charge. This is correct because $p$ orbitals are less tightly bound.
Option 2: Increase in shielding and atomic size is not the primary reason for this trend.
Option 3: Incorrect, as $s$ orbitals are more tightly bound.
Option 4: Decrease in shielding and atomic size is not relevant here.

Final Answer

The answer is: $\boxed{\text{It is easier to remove an electron in a } p \text{ orbital than an } s \text{ orbital because it experiences less effective nuclear charge.}}$

Step 4: Understanding Group 15 and Group 16 Ionization Energies

Group 15 elements have half-filled $p$ orbitals, which are more stable due to electron pairing energy. Group 16 elements have one more electron, causing electron-electron repulsion, making it easier to remove an electron.

Final Answer

The answer is: $\boxed{\text{Electron-electron repulsion in Group 16 makes it easier to remove an electron compared to Group 15.}}$

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