Electronic Configurations Of Al, Ne, And S: A Chemistry Problem

Hey guys! Let's dive into a fascinating chemistry problem that involves figuring out the electronic configurations of Aluminum (Al), Neon (Ne), and Sulfur (S). This is a fundamental concept in chemistry that helps us understand how atoms interact with each other to form molecules. We'll tackle each part step-by-step, making it super easy to grasp. So, buckle up and let's get started!

a) Electronic Configurations: Unlocking the Atom's Secrets

Let's start with the first part: determining the electronic configurations of Aluminum (13Al), Neon (10Ne), and Sulfur (16S). Electronic configuration essentially maps out how electrons are arranged within an atom's energy levels and orbitals. Think of it like the atom's address system for its electrons! This arrangement dictates the chemical behavior of an element, so it's pretty crucial stuff.

Aluminum (13Al)

Aluminum has an atomic number of 13, which means it has 13 protons and, in its neutral state, 13 electrons. To write its electronic configuration, we need to fill the electron shells and subshells according to the Aufbau principle (electrons first fill the lowest energy levels) and Hund's rule (electrons individually occupy orbitals within a subshell before pairing up). The order of filling is generally: 1s, 2s, 2p, 3s, 3p, and so on. So, for Aluminum, the electronic configuration is:

• 1s²: The first two electrons fill the 1s orbital.
• 2s²: The next two electrons fill the 2s orbital.
• 2p⁶: Six electrons fill the three 2p orbitals.
• 3s²: Two electrons fill the 3s orbital.
• 3p¹: The last electron goes into one of the 3p orbitals.

Putting it all together, the electronic configuration of Aluminum is 1s²2s²2p⁶3s²3p¹. This tells us a lot about Aluminum's reactivity and how it will likely interact with other elements. For example, the single electron in the 3p orbital makes Aluminum prone to forming bonds.

Neon (10Ne)

Next up is Neon, a noble gas with an atomic number of 10. This means it has 10 electrons to arrange. Neon is famously unreactive, and its electronic configuration explains why. Let's break it down:

• 1s²: The first two electrons fill the 1s orbital.
• 2s²: The next two electrons fill the 2s orbital.
• 2p⁶: Six electrons fill the three 2p orbitals.

So, Neon's electronic configuration is 1s²2s²2p⁶. Notice that the 2p orbitals are completely filled. This full outer shell (or octet) is what makes Neon so stable and unreactive. It doesn't need to gain, lose, or share electrons to achieve a stable configuration, unlike many other elements.

Sulfur (16S)

Finally, let's tackle Sulfur, which has an atomic number of 16 and thus 16 electrons. Sulfur is a nonmetal known for its ability to form various compounds. Its electronic configuration will give us clues about its bonding behavior:

• 1s²: The first two electrons fill the 1s orbital.
• 2s²: The next two electrons fill the 2s orbital.
• 2p⁶: Six electrons fill the three 2p orbitals.
• 3s²: Two electrons fill the 3s orbital.
• 3p⁴: Four electrons fill the 3p orbitals.

Therefore, the electronic configuration of Sulfur is 1s²2s²2p⁶3s²3p⁴. Sulfur has four electrons in its 3p orbitals, meaning it needs two more electrons to achieve a stable octet. This explains why Sulfur often forms compounds where it gains or shares two electrons, leading to its versatile chemistry.

b) Electrons in s Orbitals: A Matter of Shape

Now, let's move on to the second part of the problem: identifying the element whose electrons are placed only in s orbitals. This is a fun question that tests our understanding of orbital shapes. Remember, s orbitals are spherical, while p orbitals are dumbbell-shaped, and d orbitals have more complex shapes. The number of orbitals available at each energy level dictates how many electrons can be accommodated.

The key here is Neon (10Ne) whose electronic configuration is 1s²2s²2p⁶. No element exclusively fills only s orbitals when they have enough electrons to also fill p orbitals. This is because after the 1s and 2s orbitals are filled, electrons start filling the 2p orbitals. For an atom to have electrons only in s orbitals, it would need to have either 2 electrons (filling 1s and 2s) or just 2 electrons (filling the 1s orbital). Helium (He) has the electronic configuration 1s², hence has electrons only in s orbitals.

However, let's consider a hypothetical scenario. If the question is interpreted very literally, the element in our list that comes closest to having electrons primarily in s orbitals before filling p orbitals is Neon (10Ne). Although Neon has electrons in the 2p orbitals, it fills all the s orbitals (1s² and 2s²) before filling the 2p orbitals (2p⁶). It is important to note that this is a very strict interpretation of the question.

In a real-world scenario, Helium (He) is the only element that truly has electrons placed only in s orbitals. Its electronic configuration is 1s², meaning all its electrons reside in the spherical 1s orbital. It's a tiny atom with a simple electron arrangement, which contributes to its stability and inertness.

c) Outermost Shell Electrons: The Valence Crew

Finally, let's tackle the last part: determining the number of electrons in the outermost shell (also known as the valence shell) for each element. The electrons in the outermost shell are called valence electrons, and they're the ones primarily involved in chemical bonding. Knowing the number of valence electrons helps us predict how an element will react with others.

Aluminum (13Al)

Aluminum's electronic configuration is 1s²2s²2p⁶3s²3p¹. The outermost shell is the third shell (n=3), which contains the 3s and 3p subshells. Aluminum has 2 electrons in the 3s subshell and 1 electron in the 3p subshell. Therefore, Aluminum has a total of 3 valence electrons. This makes Aluminum prone to losing these three electrons to form a +3 ion, which is why it's commonly found in compounds like Al₂O₃.

Neon (10Ne)

Neon's electronic configuration is 1s²2s²2p⁶. The outermost shell is the second shell (n=2), which contains the 2s and 2p subshells. Neon has 2 electrons in the 2s subshell and 6 electrons in the 2p subshell. This gives Neon a total of 8 valence electrons. Remember our earlier discussion about Neon's stability? This full octet of valence electrons is the key! It makes Neon exceptionally stable and chemically inert.

Sulfur (16S)

Sulfur's electronic configuration is 1s²2s²2p⁶3s²3p⁴. The outermost shell is the third shell (n=3), containing the 3s and 3p subshells. Sulfur has 2 electrons in the 3s subshell and 4 electrons in the 3p subshell. So, Sulfur has a total of 6 valence electrons. This means Sulfur needs two more electrons to complete its octet, explaining its tendency to form bonds by gaining or sharing electrons.

Wrapping Up: Decoding Atomic Behavior

So, there you have it! We've successfully navigated through the electronic configurations of Aluminum, Neon, and Sulfur. We've figured out how their electrons are arranged, which elements have electrons solely in s orbitals (or very close to it!), and how many valence electrons each element possesses. Understanding these concepts is like having a secret decoder ring for atomic behavior. You can now predict how these elements will interact with others and form compounds. Chemistry is awesome, isn't it?

I hope this breakdown helped you guys understand electronic configurations a little better. Keep exploring the fascinating world of chemistry!