Question

Why do we believe that the three electrons in the \(2 p\) sublevel of nitrogen occupy different orbitals?

Short answer

We believe that the three electrons in the 2p sublevel of nitrogen occupy different orbitals because of Hund's rule, which states that electrons will occupy different orbitals with parallel spins when multiple orbitals have the same energy level. Nitrogen has an electron configuration of \(1s^2 2s^2 2p^3\), and the three electrons in the 2p sublevel will occupy the three different orbitals with parallel spins to create the most stable and lowest energy configuration.

Step-by-step solution

Writing the electron configuration for nitrogen

Nitrogen has an atomic number of 7, which means it has 7 electrons. We will write the electron configuration for nitrogen to know the arrangement of these electrons in subshells. The electron configuration for nitrogen is: \[1s^2 2s^2 2p^3\]

Understanding the quantum numbers

Quantum numbers are used to describe the properties of an electron in an atom. There are four quantum numbers: principal quantum number (n), azimuthal quantum number (l), magnetic quantum number (m_l), and spin quantum number (m_s). For the 2p subshell in nitrogen: - n = 2 (indicates the energy level) - l = 1 (indicates the p orbital) - m_l can be -1,0,1 (indicates the three different orbitals within the p sublevel) - m_s can be +1/2 or -1/2 (indicates the two possible spin values of an electron)

Understanding Hund's Rule

Hund's rule states that when there are multiple orbitals with the same energy level, electrons will occupy different orbitals with parallel spins. This means that they will prefer to fill one electron in each orbital before pairing up.

Applying Hund's Rule to nitrogen 2p sublevel

Nitrogen has three electrons in the 2p sublevel. Based on Hund's rule, these electrons will occupy the three different orbitals with parallel spins because this configuration has the lowest energy. This is the most stable configuration, and it is why we believe the electrons in the 2p sublevel of nitrogen occupy different orbitals.

Key concepts

Quantum Numbers

Imagine electrons as tiny particles with very detailed 'addresses' that describe their exact location within an atom. This is where quantum numbers come into play. They are essentially the ZIP code for electrons, pinpointing their positions in the vast 'city' of an atom. There are four quantum numbers—principal (n), azimuthal (l), magnetic (m_l), and spin (m_s)—each representing a specific characteristic of an electron's orbit.

For instance, the principal quantum number (n) is like the neighborhood, telling us the energy level where you'd find the electron. The larger the number, the farther from the nucleus and the higher the energy. Then there's the azimuthal (l) number, which is akin to the type of building the electron lives in, with shapes ranging from spherical to dumbbell. The magnetic quantum number (m_l) narrows this down further to the exact apartment within the building, revealing which direction in space the electron's orbital points. Finally, the spin quantum number (m_s) indicates whether the electron is 'spinning' clockwise or counterclockwise.

This intricate addressing system is essential not just for identifying electrons but also for predicting how they will bond with other atoms. It's a fundamental concept that helps explain the intricate dance of electrons in their atomic orbits.

Hund's Rule

Imagine you enter a movie theater with your friends and find it almost empty. You're likely to sit in your own seat rather than sharing a seat with a friend, right? That's somewhat similar to Hund's rule in the world of electrons. Hund's rule acts like social guidance for electrons, suggesting they'd rather have their own 'seat' (orbital) when they 'watch the atomic movie.'

Hund's rule tells us that electrons will fill empty orbitals of the same energy (degenerate orbitals) with parallel spins, avoiding each other until they have to pair up. Think of electrons as tiny, negative magnets. They spin in a way that creates a small magnetic field. If two electrons with the same spin are in different orbitals, they are like magnets repelling each other, which is low-energy and stable. But when orbitals are forced to share electrons, or when electrons 'sit together,' their opposite spins create an awkward tension, like holding hands with someone while giving them the silent treatment—uncomfortable and high-energy. This is why, in the case of nitrogen with three electrons in the 2p level, each electron claims its own orbital throne before sharing space.

Atomic Structure

The beauty and complexity of the atomic structure are akin to an enthralling cosmic ballet. At the center of this dance is the nucleus—the sun of the atom—made up of protons and neutrons. Swirling around it are the electron clouds, which are regions where electrons are likely to be found, not unlike planets orbiting their star.

Breaking down the atom even further, the electron clouds contain various energy levels, often visualized as concentric circles. Each level is further subdivided into subshells, named with letters like s, p, d, and f. These subshells resemble different orbits with unique shapes and capacities for housing electrons.

Layering Complexity

Within these subshells, there are orbitals—think of them as individual seats on a Ferris wheel, each able to hold a maximum of two electrons. Following our rules of quantum numbers and Hund's rule, these electrons fill the atomic structure in a way that minimizes energy and maximizes stability. Through their complicated yet methodical movements and interactions, electrons in their respective orbitals define not just the chemical properties of an element but also the very nature of the material universe.