Question

For each pair listed here, state which one has a higher first ionization energy and explain your choice: (a) H or \(\mathrm{H}_{2},(\mathrm{~b}) \mathrm{N}\) or \(\mathrm{N}_{2},\) (c) \(\mathrm{O}\) or \(\mathrm{O}_{2},\) (d) \(\mathrm{F}\) or \(\mathrm{F}_{2}\).

Short answer

(a) H has higher first ionization energy than H2, (b) N has higher first ionization energy than N2, (c) O has higher ionization energy than O2, and (d) F has higher ionization energy than F2. This is because atoms usually have higher first ionization energy than their molecular forms.

Step-by-step solution

Comparing Hydrogen Atom with Hydrogen Molecule

Let's first compare the hydrogen atom (H) and the hydrogen molecule (H2). The hydrogen molecule is more stable and requires less energy to remove an electron because of the electron pairing in molecules. Therefore, a hydrogen atom (H) has a higher first ionization energy than a hydrogen molecule (H2).

Comparing Nitrogen Atom with Nitrogen Molecule

The nitrogen molecule (N2) has a triple bond which is very strong and stable. This makes it hard to remove an electron from N2, lowering its ionization energy. Therefore, nitrogen atom (N) has a higher first ionization energy than nitrogen molecule (N2).

Comparing Oxygen Atom with Oxygen Molecule

For the oxygen atom and molecule (O and O2), the oxygen molecule (O2) also has a double bond which makes it more stable requiring less energy to remove an electron compared to the oxygen atom. Therefore, an oxygen atom (O) has a higher first ionization energy than an oxygen molecule (O2).

Comparing Fluorine Atom with Fluorine Molecule

Lastly, comparing the fluorine atom vs. molecule (F vs. F2), the F2 molecule is more stable because of its electron pairing, so it is easier to remove an electron from F2. Consequently, a fluorine atom (F) has a higher first ionization energy than a fluorine molecule (F2).

Key concepts

Ionization Energy Comparison

Ionization energy can be thought of as the amount of energy needed to strip an electron from an atom or molecule. One of the fundamental questions in chemistry is how this energy requirement differs across various substances. When comparing atomic versus molecular ionization energies, the main consideration is the overall stability provided by chemical bonds. Atoms, such as Hydrogen (H), Nitrogen (N), Oxygen (O), and Fluorine (F), tend to have higher ionization energies than their molecular counterparts - diatomic molecules H₂, N₂, O₂, and F₂.

This enhanced ionization energy arises from the state of the electrons within the atoms and molecules. In a stand-alone atom, electrons occupy orbital that require more energy to remove them than when they are in a molecule. The chemical bonds in molecules often lead to a state where electrons are more shared and less tightly bound to any particular atom, making them easier to remove. Therefore, it is consistent across each pair that the atomic form has a higher first ionization energy than the molecular form.

Atomic vs Molecular Ionization

Diving deeper into the comparison between atomic and molecular ionization, it's essential to understand how the structure of molecules influences their ionization energies. In a molecule, electrons are involved in bond formation between atoms, which has a consequential effect on their ease of removal. For instance, in diatomic hydrogen (H₂), the two hydrogen atoms create a stable electron pair that can be ionized with less energy than for isolated hydrogen atoms.

Bond Strength and Ionization

Furthermore, the type and strength of bonds play a significant role. The triple bonds in nitrogen molecules (N₂) make them incredibly stable and resistant to having electrons removed, although not quite to the extent that surpasses the ionization energy of isolated nitrogen atoms. Understanding this concept is crucial for explaining the trends observed in ionization energies across different chemical species.

Chemical Bonding and Stability

The stability of chemical substances is directly tied to the nature of the chemical bonds they possess. Electrons are the glue that holds atoms together, creating molecules, and the chemical bond type – whether covalent, ionic, or metallic – influences the energy required to alter these molecules.

In our textbook examples, we're dealing with covalent bonds where electron sharing provides stability. The stability of molecules like N₂ and O₂, with triple and double bonds respectively, is so great that even though you might expect a molecule to usually have a lower first ionization energy than an atom, molecules with exceptionally strong bonds can approach or exceed the ionization energies of their constituent atoms. The first ionization energy reflects how tightly an electron is held within the atomic or molecular environment and is a valuable indicator of stability. To sum up, stronger chemical bonds generally lead to increased stability and lower ionization energies for molecules.