Question

Use electron configurations to explain why the ionization energy for fluorine is greater than that for oxygen.

Short answer

The ionization energy for fluorine is greater than for oxygen due to their electron configurations. In oxygen, there is one 'p' orbital that contains a single electron, making it easier to remove this electron. However, in fluorine, each electron in the 'p' orbital is paired up, thus it takes more energy to remove an electron from these stable pairs.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The ionization energy of an atom depends on the attraction between the positive nucleus and the negative electrons, and the repulsion between electrons. Electrons closer to the nucleus are more tightly held and require more energy to remove. Additionally, it's important to keep in mind that lonization energy generally increases from left to right across a period, due to increasing nuclear charge.

Electron Configuration of Fluorine and Oxygen

We need to recall the electron configurations for both oxygen and fluorine. For oxygen, it's 1s² 2s² 2p⁴, and for fluorine, it's 1s² 2s² 2p⁵.

Analyzing the Ionization Energy based on Electron Configuration

For oxygen, the second electron in one 'p' orbital doesn't have an electron pair, making it easier to remove. Therefore, less energy is required. However, in the case of fluorine, each electron in the 'p' orbital has its own electron pair and the energy required to remove an electron from these stable pairs is higher. That's why the ionization energy for fluorine is greater than for oxygen.

Key concepts

Understanding Electron Configuration

Electron configuration refers to the arrangement of electrons in an atom's orbitals. Each element has a unique electron configuration that determines its chemical properties and reactions. For instance, **oxygen** has the electron configuration of 1s² 2s² 2p⁴, which means it has a total of six electrons placed within its orbitals. **Fluorine**, on the other hand, has an electron configuration of 1s² 2s² 2p⁵, indicating it has nine electrons. The placement of electrons within these orbitals follows a specific order, dictated by principles like the Aufbau principle, Hund’s Rule, and Pauli’s Exclusion Principle.
- **Aufbau Principle:** Electrons fill orbitals starting from the lowest energy level to the highest. - **Hund’s Rule:** Electrons will occupy all degenerate orbitals singly before pairing up. - **Pauli Exclusion Principle:** No two electrons can have identical quantum numbers, limiting each orbital to two electrons of opposite spins.
Understanding these rules helps explain how atoms like oxygen and fluorine differ, and why their ionization energies are not the same.

Significance of Periodic Trends

Periodic trends describe patterns in the periodic table that help predict the behaviors of elements. **Ionization energy**, atomic radius, and electronegativity are among these trends. As you move across a period from left to right, ionization energy generally increases. This happens because the electrons are pulled closer to the nucleus due to increased nuclear charge, making them harder to remove. In this case, fluorine, which is to the right of oxygen in the same period, shows higher ionization energy. This is due to its higher nuclear charge and more compact electron cloud, compared to oxygen. Knowing these trends can help predict elements’ chemical properties and reactions.

Role of Nuclear Charge

Nuclear charge is essentially the total charge of the nucleus, determined by the number of protons it contains. A greater number of protons means a higher nuclear charge, which exerts a stronger attractive force on the surrounding electrons.
- **Impact on Electrons:** The higher the nuclear charge, the more strongly electrons are attracted towards the nucleus. This decreases the size of the electron cloud and increases ionization energy. - **Comparison between Oxygen and Fluorine:** Fluorine has one more proton than oxygen, thus a higher nuclear charge. The electron cloud is pulled in more tightly, and the electrons are held more firmly.
Therefore, even though both oxygen and fluorine have electrons in similar energy levels, fluorine’s higher nuclear charge means its electrons are more tightly bound, and thus, it requires more energy to ionize compared to oxygen.