Problem 88 - Chemistry Matter and Change

Question

Which electron configuration notation describes an atom in an excited state? $$\begin{array}{l}{\text { a. }[\mathrm{Ar}] 4 \mathrm{s}^{2} 3 \mathrm{d}^{10} \mathrm{4p}^{2}} \\ {\text { b. }[\mathrm{Ne}] 3 \mathrm{s}^{2} 3 \mathrm{p}^{5}} \\ {\text { c. }[\mathrm{Kr}] 5 \mathrm{s}^{2} 4 \mathrm{d}^{1}} \\ {\text { d. }[\mathrm{Ar}] 4 \mathrm{s}^{2} 3 \mathrm{d}^{8} 4 \mathrm{p}^{1}}\end{array}$$

Step-by-Step Solution

Answer

The configuration in option c represents an atom in an excited state.

Step 1: Understand Electron Configurations

Electron configurations describe the distribution of electrons in an atom's orbitals. It follows the order of filling from lower to higher energy levels and respects Hund’s rule and the Pauli exclusion principle. Atoms can be in a ground state (lowest energy configuration) or an excited state (one or more electrons are in higher energy orbitals than necessary).

Step 2: Identify Ground State Configurations

For each options a, b, c, and d, determine if the configuration corresponds to a ground state or if it violates the normal order of filling. Generally, electrons fill orbitals in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on.

Step 3: Check Option a

Option a is \[ \text{[Ar]} 4s^2 3d^{10} 4p^2 \]. This is a valid configuration for a ground state as electrons fill the 4s before starting to fill the 3d and then 4p orbitals. It is not in an excited state.

Step 4: Check Option b

Option b is \[ \text{[Ne]} 3s^2 3p^5 \]. This is also a ground-state configuration for Chlorine where all electrons are in the lowest energy orbitals available without skipping any order.

Step 5: Inspect Option c

Option c is \[ \text{[Kr]} 5s^2 4d^1 \]. For this atom, electrons normally fill 5s followed by 4d. Thus, jumping into 4d right after filling 5s with only one electron indicates the excited state. Normally, the 5s would be completed followed by filling of the 4d.

Step 6: Review Option d

Option d is \[ \text{[Ar]} 4s^2 3d^8 4p^1 \]. The electron configuration follows the normal filling order, where 4s is filled first, followed by partial filling of 3d and then starting to fill the 4p orbital.

Step 7: Conclusion

Option c includes an excited state because it skips the usual pattern of electron filling—having a 4d orbital filled before completely filling other lower-energy orbitals such as 4p. This aligns with the definition of an excited state where an electron is raised to a higher energy orbital.

Key Concepts

Electron Configurations

Electron configurations are a way to represent the arrangement of electrons in an atom. Think of an atom as a hotel, and each orbital as a room where electrons "stay." These configurations follow specific rules and patterns. Electrons fill available "rooms" in a specific sequence, based on increasing energy levels. We start with the 1s orbital, fill it, and move on to 2s, then 2p, and so on. Each orbital can hold a certain number of electrons: s orbitals can hold up to 2, p can hold up to 6, and d can hold up to 10.

• s orbitals: Hold 2 electrons maximum
• p orbitals: Hold 6 electrons maximum
• d orbitals: Hold 10 electrons maximum

Understanding this pattern is crucial as it tells us about the atom's reactivity and other properties. For instance, the electron configuration \( \text{[Ar]} 3d^{10} 4s^{2} 4p^{2} \) means we have two electrons in the 4p "room." This can indicate the atom's potential to react or bond with other atoms.

Ground State

The ground state of an atom is its most stable form. Here, all electrons occupy the lowest possible energy levels. Think of it like electrons are taking the simplest route to "stay" in their orbitals. This is often the configuration you find on the periodic table under each element's name. The ground state represents the starting point when describing an element's electron configuration.

• Lowest Energy Configuration: Electrons are in the lowest energy orbitals possible.
• Example: Chlorine, \( \text{[Ne]} 3s^{2} 3p^{5} \), follows the lowest energy pattern.

In theory, this is where every element "wants" to be, because it's the most stable configuration. Any change from this state gives rise to an excited state, which involves electrons jumping to higher energy levels than usual.

Energy Levels

Energy levels are like tiers in an apartment building where electrons "live." Each level represents a different amount of energy. The farther away an energy level (or electron) is from the nucleus, the more energy it has. You can think of the nucleus as the central hub and electrons, which have more energy, are often found farther from this hub.

• Lower Energy Levels: Closer to the nucleus; electrons here are in the ground state.
• Higher Energy Levels: Farther from the nucleus; associated with excited states.

When electrons absorb energy, they can move to higher energy levels, entering what's called the excited state. This movement can be temporary, and the electrons might return to their original levels, often releasing energy as light or heat.

Orbital Filling Order

The order in which electron orbitals are filled follows a strict pattern based on their energy level and type (s, p, d, f). This is known as the Aufbau principle, meaning "building up" in German. Electrons fill lower energy orbitals first before moving to higher ones. It's like filling seats in a movie theater from front to back: you fill the front rows first before moving to the back.

• 1s \( \rightarrow\) 2s \( \rightarrow\) 2p \( \rightarrow\) 3s \( \rightarrow\) 3p \( \rightarrow\) 4s \( \rightarrow\) 3d \( \rightarrow\) 4p \[...\]

Orbital filling order helps us predict the arrangement and behavior of electrons in different elements. When an electron skips from its expected orbital, such as moving from 4s to 4d instead of 4p, it creates an excited state. Understanding this concept is essential for identifying excited states versus ground states in atoms.