Pseudo-Noble Gas: Unlock Electron Configuration Secrets!
Electron configuration, a concept central to understanding atomic structure, dictates the arrangement of electrons within an atom. Understanding the nuances of electron configuration is essential for predicting the chemical behavior of elements. Specifically, the concept of pseudo-noble-gas configurations, often addressed in materials science research, arises when certain transition metals attain a stable electron arrangement that mimics the electron configurations of noble gases, even though their total number of valence electrons differs. The understanding of stability in electron configurations is often approached using Hund's rules, which assist in determining how orbitals are filled, and further informed by the work of scientists in institutions such as the National Institute of Standards and Technology (NIST). Therefore, when we ask which of the following is a pseudo-noble-gas electron configuration?, we are exploring a subtle but significant aspect of chemical bonding and reactivity, often modeled using computational chemistry software to predict material properties.
Image taken from the YouTube channel Professor Dave Explains , from the video titled Pseudo Noble Gas Electron Configurations .
Electron configurations dictate the chemical behavior of elements. While a true noble gas configuration is the epitome of stability, certain ions achieve a semblance of this stability through what we call pseudo-noble gas configurations. These configurations, particularly prevalent among transition metals and in coordination complexes, grant unique properties to these species.
This section will delve into the intricacies of pseudo-noble gas configurations. We'll explore their definition, their distinction from true noble gas configurations, and their significance in the realm of transition metal chemistry. Finally, we will introduce the common question format: "which of the following is a pseudo-noble-gas electron configuration?"
Defining the Pseudo-Noble Gas Configuration
In layman's terms, a pseudo-noble gas configuration is an electron arrangement in an ion where the d subshell is completely filled, specifically the (n-1)d¹⁰ configuration. This is a stable arrangement, mimicking some aspects of the filled valence shell seen in true noble gases.
However, it's crucial to understand that the valence shell itself isn't entirely filled like in a noble gas. It is the filled (n-1)d subshell which imparts the stability.
Distinguishing Pseudo from True: A Key Difference
The key difference lies in the completeness of the valence shell. Noble gases, such as Neon ([He]2s²2p⁶) or Argon ([Ne]3s²3p⁶), have their outermost s and p orbitals completely filled, resulting in a full octet (or duet for Helium). This complete filling of the valence shell leads to their exceptional inertness.
In contrast, ions with pseudo-noble gas configurations, like Zn²⁺ ([Ar]3d¹⁰), possess a filled (n-1)d subshell, but their outermost s and p orbitals are typically empty.
This distinction is subtle but important, as it explains why pseudo-noble gas configurations offer relative stability, not the absolute inertness of a true noble gas.
Occurrence in Transition Metals and Their Ions
Pseudo-noble gas configurations are most frequently observed in transition metals and their ions. Transition metals, characterized by their partially filled d orbitals, can achieve a pseudo-noble gas configuration upon losing electrons to form cations.
For instance, Zinc (Zn), Cadmium (Cd), and Mercury (Hg) readily form +2 ions (Zn²⁺, Cd²⁺, Hg²⁺) with a d¹⁰ configuration. Similarly, Copper (Cu), Silver (Ag), and Gold (Au) can form +1 ions (Cu⁺, Ag⁺, Au⁺) that also exhibit this pseudo-noble gas stability. The ability to attain this configuration influences their chemical behavior and the types of compounds they form.
"Which of the Following..." - Decoding the Question
A common question in chemistry exams is, "Which of the following is a pseudo-noble-gas electron configuration?". This question tests your understanding of the definition and your ability to recognize the characteristic d¹⁰ configuration in an ion.
Answering this question requires you to:
- Identify the element or ion associated with each electron configuration.
- Determine if the ion has a filled (n-1)d¹⁰ subshell.
- Verify that the outer s and p orbitals are either empty or filled.
By mastering the concepts presented in this article, you'll be well-equipped to tackle this type of question and gain a deeper appreciation for the fascinating world of electron configurations.
Electron Configuration: A Quick Review
Understanding electron configurations is foundational to grasping chemical behavior. Before delving into the nuances of pseudo-noble gas configurations, a review of the fundamentals is essential. We will explore the notation, the rules governing electron filling, and how these concepts apply to both neutral atoms and ions.
Decoding Electron Configuration Notation
Electron configuration notation provides a concise way to represent the arrangement of electrons within an atom. Each entry consists of three components: the principal quantum number (n), the orbital type (s, p, d, or f), and the number of electrons in that orbital.
For example, 1s² indicates that two electrons occupy the 1s orbital. The principal quantum number, 1, specifies the energy level, and 's' denotes the shape of the orbital. Superscript '2' represents the number of electrons populating that specific orbital. Understanding this notation is crucial for interpreting electron configurations and predicting chemical properties. Familiarity with this is essential.
The Aufbau Principle: Building Electron Configurations
The Aufbau principle dictates the order in which electrons fill atomic orbitals. Electrons first occupy the lowest energy orbitals available before moving to higher energy levels. This principle provides a systematic approach to constructing electron configurations.
The general filling order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. This order can be visualized using the Madelung rule, also known as the n+l rule, where n is the principal quantum number and l is the azimuthal quantum number (0 for s, 1 for p, 2 for d, and 3 for f). Orbitals with lower n+l values are filled first.
Exceptions to the Aufbau Principle
While the Aufbau principle offers a reliable guideline, there are exceptions. These deviations arise from the energetic stability associated with half-filled and fully filled d and f subshells.
For instance, Chromium (Cr) has an expected configuration of [Ar] 4s² 3d⁴, but its actual configuration is [Ar] 4s¹ 3d⁵. Similarly, Copper (Cu) deviates from [Ar] 4s² 3d⁹ to [Ar] 4s¹ 3d¹⁰. These exceptions highlight the subtle interplay of electron-electron interactions and energy minimization.
Hund's Rule and the Pauli Exclusion Principle
Beyond the Aufbau principle, Hund's rule and the Pauli Exclusion Principle further refine electron configuration. Hund's rule states that electrons individually occupy each orbital within a subshell before any orbital is doubly occupied. This minimizes electron-electron repulsion and maximizes stability. All electrons in singly occupied orbitals have the same spin (either +1/2 or -1/2).
The Pauli Exclusion Principle stipulates that no two electrons in an atom can have the same set of four quantum numbers. This principle implies that each orbital can hold a maximum of two electrons, each with opposite spin. These principles ensure a unique and stable electron arrangement for each atom.
Electron Configurations of Ions: Adding or Removing Electrons
Ions are formed when atoms gain or lose electrons. Cations (positive ions) result from the removal of electrons, while anions (negative ions) result from the addition of electrons. When determining the electron configuration of an ion, it is crucial to account for the change in the number of electrons.
For cations, electrons are removed from the outermost shell first. For example, Iron (Fe) has an electron configuration of [Ar] 4s² 3d⁶. To form Fe²⁺, two electrons are removed from the 4s orbital, resulting in [Ar] 3d⁶. For anions, electrons are added to the lowest energy unoccupied orbital. The overall charge of the ion must always be explicitly stated.
Having revisited the fundamentals of electron configuration, we now turn our attention to the exemplary stability exhibited by the noble gases. These elements serve as the benchmark against which other electron configurations are often compared, especially when discussing pseudo-noble gas configurations. Understanding their unique electronic structures is paramount to appreciating the driving forces behind chemical behavior.
Noble Gases: The Gold Standard
The noble gases—helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn)—occupy the extreme right of the periodic table (Group 18) and are renowned for their exceptional chemical inertness. Their reluctance to form chemical bonds stems directly from their uniquely stable electron configurations, making them the "gold standard" in terms of electronic stability.
Electron Configurations of Noble Gases
Each noble gas, with the exception of helium, possesses a full complement of eight electrons in its outermost, or valence, shell. Helium achieves stability with just two valence electrons. This configuration, denoted as ns²np⁶ (or 1s² for helium), is the key to their unreactive nature.
- Helium (He): 1s²
- Neon (Ne): 1s² 2s² 2p⁶
- Argon (Ar): 1s² 2s² 2p⁶ 3s² 3p⁶
- Krypton (Kr): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶
- Xenon (Xe): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶
- Radon (Rn): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p⁶
The Octet and Duet Rules
The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full valence shell containing eight electrons (an ns²np⁶ configuration). This rule governs the behavior of many elements in the periodic table, driving them toward achieving the stability of noble gases.
Helium, with its electron configuration of 1s², follows the duet rule, achieving stability with just two electrons in its valence shell.
Stability and Inertness: The Result of Filled Valence Shells
The chemical inertness of noble gases is a direct consequence of their filled valence shells. These elements already possess the most stable electron configuration possible, requiring a significant input of energy to either add or remove electrons. This high ionization energy and low electron affinity contribute to their unwillingness to participate in chemical reactions.
Because their valence shells are already complete, noble gases exhibit little tendency to form chemical bonds with other atoms. This absence of chemical reactivity is why they were historically referred to as "inert gases," though it's now known that heavier noble gases can form compounds under specific conditions.
Their filled valence shell represents a state of minimal potential energy, making the noble gases exceptionally stable and unreactive. It is this stability that other elements strive to emulate through chemical bonding and ionization. The concept of a filled valence shell is central to understanding chemical bonding and reactivity.
Having revisited the fundamentals of electron configuration, we now turn our attention to the exemplary stability exhibited by the noble gases. These elements serve as the benchmark against which other electron configurations are often compared, especially when discussing pseudo-noble gas configurations. Understanding their unique electronic structures is paramount to appreciating the driving forces behind chemical behavior.
Pseudo-Noble Gas Configurations Explained: The d¹⁰ Case
While the noble gases flaunt their perfectly filled valence shells, a select group of elements achieve a noteworthy semblance of stability through what we term pseudo-noble gas configurations. Among these, the d¹⁰ configuration stands out as the most prevalent and readily understood example. This configuration, while not mirroring the ns²np⁶ arrangement of noble gases, confers a relative stability that significantly influences the chemical behavior of the elements that exhibit it.
Defining the d¹⁰ Configuration
A d¹⁰ configuration signifies that an element or ion possesses a completely filled d subshell in the (n-1) shell, where n represents the outermost principal quantum number. In simpler terms, all ten possible electron slots within the d orbitals of the penultimate (second to last) shell are occupied. It's crucial to note that this doesn't necessarily equate to a fully occupied valence shell.
The Stability of a Filled (n-1)d Subshell
The stability associated with the d¹⁰ configuration arises from the symmetrical distribution of electron density within the d orbitals. This symmetrical arrangement minimizes electron-electron repulsions and leads to a lower overall energy state.
Furthermore, completely filled subshells, including the d¹⁰ configuration, exhibit a spherical symmetry, contributing to increased stability and reduced reactivity. The electronic arrangement is balanced, making the element or ion less prone to participate in chemical reactions.
Examples: Zinc, Cadmium, and Mercury
Several elements readily form ions with stable d¹⁰ configurations. Prominent examples include zinc (Zn²⁺), cadmium (Cd²⁺), and mercury (Hg²⁺).
These elements, upon losing two electrons, attain a d¹⁰ configuration in their (n-1) shell, resulting in stable and relatively unreactive ions.
Oxidation States and Pseudo-Noble Gas Configurations
It is important to understand the oxidation states in which these elements exhibit these configurations. Zinc, cadmium, and mercury most commonly achieve the d¹⁰ configuration in their +2 oxidation state (Zn²⁺, Cd²⁺, Hg²⁺). This is because losing two electrons from their s orbitals results in the stable d¹⁰ arrangement.
Other elements, such as copper, silver, and gold, achieve the d¹⁰ configuration in their +1 oxidation state, forming Cu⁺, Ag⁺, and Au⁺ ions, which also display relative stability due to the filled d subshell.
Filled d-orbital vs. Filled Valence Shell
It's essential to distinguish between a filled d-orbital and a filled valence shell. A filled valence shell, like those found in noble gases (except Helium), involves the complete occupancy of both the s and p orbitals in the outermost shell (ns²np⁶).
A filled d-orbital (the d¹⁰ configuration), on the other hand, refers only to the complete occupancy of the d orbitals in the (n-1) shell.
The key difference lies in the location of the filled orbitals and the completeness of the outermost shell. While a filled d-orbital contributes to stability, it doesn't necessarily result in the same level of inertness as a fully occupied valence shell. Elements with d¹⁰ configurations still exhibit some reactivity, particularly in coordination chemistry, which we will explore in a later section.
Having established the underlying principles of the d¹⁰ configuration, it is crucial to illustrate these concepts with concrete examples. Examining specific elements and their corresponding ions allows for a deeper understanding of how pseudo-noble gas configurations arise and influence their chemical behavior.
Examples of Elements Exhibiting Pseudo-Noble Gas Configurations
Several elements, primarily transition metals, readily form ions that exhibit a d¹⁰ electron configuration. This section will delve into specific examples, highlighting the electron configurations of both the neutral atom and the resulting ion. This comparison will underscore the electronic rearrangement that leads to the pseudo-noble gas configuration.
Group 12 Elements: Zinc, Cadmium, and Mercury
The Group 12 elements – zinc (Zn), cadmium (Cd), and mercury (Hg) – are textbook examples of elements that readily form stable divalent cations (2+ ions) with a d¹⁰ configuration.
Zinc (Zn) and Zinc(II) ion (Zn²⁺)
Neutral zinc (Zn), with an atomic number of 30, has the electron configuration [Ar] 3d¹⁰ 4s². Upon losing two electrons to form the Zn²⁺ ion, the 4s electrons are removed. This results in the ion having the electron configuration [Ar] 3d¹⁰.
This configuration showcases a completely filled 3d subshell, mimicking the electronic stability of noble gases. It's important to note that the 4s orbital is empty in the Zn²⁺ ion.
Cadmium (Cd) and Cadmium(II) ion (Cd²⁺)
Cadmium (Cd), element number 48, exhibits a similar behavior. Its neutral electron configuration is [Kr] 4d¹⁰ 5s². Upon ionization to Cd²⁺, the two 5s electrons are lost, resulting in the ion having the electron configuration [Kr] 4d¹⁰.
Again, we observe a completely filled d subshell in the (n-1) shell, lending stability to the Cd²⁺ ion. The 5s orbital in this instance is, again, empty.
Mercury (Hg) and Mercury(II) ion (Hg²⁺)
Mercury (Hg), with atomic number 80, presents a slightly more complex electron configuration due to the presence of f-orbitals. The neutral atom has the configuration [Xe] 4f¹⁴ 5d¹⁰ 6s². Upon forming the Hg²⁺ ion, the two 6s electrons are removed, leaving the ion with the configuration [Xe] 4f¹⁴ 5d¹⁰.
The filled 4f subshell doesn't directly contribute to the pseudo-noble gas stability. The critical factor is the filled 5d subshell. In this case, the 6s orbital is also empty.
Group 11 Elements: Copper, Silver, and Gold
The Group 11 elements – copper (Cu), silver (Ag), and gold (Au) – also exhibit pseudo-noble gas configurations, but in their monovalent (1+) ionic state. This is due to the removal of a single electron.
Copper (Cu) and Copper(I) ion (Cu⁺)
Copper (Cu), with an atomic number of 29, has an exception to the typical filling order due to electron-electron repulsion. Its electron configuration is [Ar] 3d¹⁰ 4s¹. Upon losing one electron to form the Cu⁺ ion, the 4s electron is removed.
The Cu⁺ ion then has the electron configuration [Ar] 3d¹⁰, giving it a pseudo-noble gas configuration.
Silver (Ag) and Silver(I) ion (Ag⁺)
Silver (Ag), element number 47, also benefits from the same exceptional behavior seen in Copper. Its neutral electron configuration is [Kr] 4d¹⁰ 5s¹. Upon ionization to Ag⁺, the single 5s electron is lost.
This results in the Ag⁺ ion having the electron configuration [Kr] 4d¹⁰. The 5s orbital is empty in the Ag⁺ ion.
Gold (Au) and Gold(I) ion (Au⁺)
Gold (Au), with atomic number 79, again presents a more complex electron configuration due to the presence of f-orbitals. The neutral atom has the configuration [Xe] 4f¹⁴ 5d¹⁰ 6s¹. Upon forming the Au⁺ ion, the 6s electron is removed, leaving the ion with the configuration [Xe] 4f¹⁴ 5d¹⁰.
Similar to Hg²⁺, the filled 4f subshell doesn't directly contribute to the pseudo-noble gas stability. The critical factor is again the filled 5d subshell.
Summary of Electron Configurations
| Element/Ion | Neutral Atom Electron Configuration | Ion Electron Configuration |
|---|---|---|
| Zn | [Ar] 3d¹⁰ 4s² | Zn²⁺: [Ar] 3d¹⁰ |
| Cd | [Kr] 4d¹⁰ 5s² | Cd²⁺: [Kr] 4d¹⁰ |
| Hg | [Xe] 4f¹⁴ 5d¹⁰ 6s² | Hg²⁺: [Xe] 4f¹⁴ 5d¹⁰ |
| Cu | [Ar] 3d¹⁰ 4s¹ | Cu⁺: [Ar] 3d¹⁰ |
| Ag | [Kr] 4d¹⁰ 5s¹ | Ag⁺: [Kr] 4d¹⁰ |
| Au | [Xe] 4f¹⁴ 5d¹⁰ 6s¹ | Au⁺: [Xe] 4f¹⁴ 5d¹⁰ |
This table clearly demonstrates how these elements, upon ionization, achieve a d¹⁰ configuration in the (n-1) shell. This electronic arrangement contributes significantly to their chemical stability and reactivity. When answering the question "which of the following is a pseudo-noble-gas electron configuration?", keep these examples in mind.
Beyond d¹⁰: Exploring Less Common Pseudo-Noble Gas Configurations
While the d¹⁰ configuration reigns supreme as the most frequently encountered and conceptually straightforward example of a pseudo-noble gas configuration, it is important to acknowledge that other electronic arrangements can, under specific circumstances, mimic the stability associated with noble gases. These instances are generally less prevalent and involve more intricate electronic interactions, making them less emphasized in introductory chemistry curricula.
The Role of Filled f-Orbitals (f¹⁴)
Filled f-orbitals (f¹⁴) can, under certain conditions, contribute to a pseudo-noble gas-like stability. This is particularly relevant in the chemistry of lanthanides and actinides.
The completely filled f-subshell exhibits a spherical symmetry, similar to filled s and p orbitals in noble gases. This symmetry leads to a more stable electronic environment.
However, the impact of f-orbitals on chemical behavior is more complex than that of d-orbitals. This is due to their deeper penetration into the electron core and weaker interactions with ligands in coordination complexes. Consequently, f¹⁴ configurations are less straightforward in conferring "pseudo-noble gas" character than the d¹⁰ configuration.
d¹⁰ Dominance in Introductory Chemistry
It is crucial to emphasize that the d¹⁰ configuration remains the most significant and widely studied case of pseudo-noble gas configurations, especially within the context of introductory chemistry. Elements such as zinc, cadmium, mercury, copper(I), silver(I), and gold(I) readily adopt this configuration upon ionization. This makes them excellent examples for understanding the principles governing electronic stability and chemical reactivity.
The relative simplicity and clear-cut examples associated with d¹⁰ configurations make them invaluable for grasping the foundational concepts. Tackling the nuances of f¹⁴ configurations is generally reserved for more advanced courses in inorganic and coordination chemistry.
The Inert Pair Effect: A Related Phenomenon
The inert pair effect provides another perspective on pseudo-noble gas stability, particularly for heavier elements in groups 13-16.
This effect describes the tendency of the heavier elements to form ions with oxidation states two less than their group oxidation state. This is often attributed to the reluctance of the ns² electrons to participate in bonding.
The inert pair effect doesn't directly result in a d¹⁰ configuration. However, it does lead to a relatively stable electron configuration where the s-electrons remain tightly bound to the nucleus, effectively mimicking the behavior of a filled valence shell.
For example, thallium (Tl) exhibits a stable +1 oxidation state (Tl⁺) in addition to the +3 oxidation state. The Tl⁺ ion has an electron configuration of [Xe] 4f¹⁴ 5d¹⁰ 6s², where the 6s² electrons are relatively inert. While not strictly a pseudo-noble gas configuration in the d¹⁰ sense, it reflects a similar principle of achieving stability through a specific electronic arrangement.
The exploration of these alternative scenarios serves to broaden the understanding of electronic stability. They provide a more comprehensive perspective on the factors influencing chemical behavior beyond the standard d¹⁰ paradigm.
Relevance and Applications: Coordination Chemistry and Beyond
Having explored the fundamental principles and examples of pseudo-noble gas configurations, it's crucial to understand their practical significance within the broader context of chemistry. The impact of these electronic arrangements extends far beyond simple ionic stability. They play a pivotal role, particularly in coordination chemistry, significantly shaping the properties and reactivity of metal complexes.
Pseudo-Noble Gas Configurations and Coordination Complex Stability
The stability of coordination complexes, intricate structures formed by a central metal ion surrounded by ligands, is profoundly influenced by the electronic configuration of the metal. When a metal ion adopts a pseudo-noble gas configuration, such as d¹⁰, it can contribute significantly to the overall stability of the resulting complex.
This stability arises from the filled d-subshell, which minimizes ligand field stabilization energy (LFSE) in certain geometries.
Ions with d¹⁰ configurations typically exhibit a preference for tetrahedral or linear geometries, as these geometries do not introduce any LFSE. This preference is in contrast to other d-electron counts that may favor octahedral or square planar geometries due to LFSE considerations.
The absence of LFSE in these cases effectively reduces the destabilizing effects from ligand repulsion, thereby enhancing the overall stability of the complex.
Influence on Electronic and Magnetic Properties
Beyond stability, pseudo-noble gas configurations also exert a considerable influence on the electronic and magnetic properties of coordination complexes. The d¹⁰ configuration, being completely filled, results in diamagnetic behavior.
This means that complexes containing metal ions with this electron configuration are not attracted to an external magnetic field.
Moreover, the electronic spectra of these complexes are often simplified due to the absence of d-d transitions, which are common in complexes with partially filled d-orbitals. This simplification can make these complexes useful in various applications, such as catalysis and sensing.
Example: Tetracyanozincate(II)
A classic example illustrating the impact of pseudo-noble gas configurations is the tetracyanozincate(II) ion, [Zn(CN)₄]²⁻. Zinc(II) possesses a d¹⁰ electron configuration.
The complex adopts a tetrahedral geometry. As previously noted, this is in part because that geometry avoids any ligand field stabilization energy.
The resulting complex is colorless (due to the absence of d-d transitions) and diamagnetic, directly reflecting the electronic characteristics conferred by the zinc(II) ion's d¹⁰ configuration. This example showcases how a pseudo-noble gas configuration contributes to predictable and well-defined properties in a coordination complex, making them valuable in various chemical applications and studies.
Video: Pseudo-Noble Gas: Unlock Electron Configuration Secrets!
FAQ: Pseudo-Noble Gas Electron Configuration
What exactly is a pseudo-noble gas electron configuration?
A pseudo-noble gas configuration is when a transition metal ion achieves an electron arrangement resembling a noble gas, specifically with a filled d subshell (d¹⁰). It's stable, but not exactly the same as a true noble gas configuration with a filled s and p subshell in the outermost shell. Example: Zn²⁺.
How does a metal ion attain a pseudo-noble gas configuration?
Metals lose electrons to form positive ions (cations). Some transition metals, after losing electrons, end up with a filled d subshell. This electron loss can lead to the stable d¹⁰ arrangement, which is which of the following is a pseudo-noble-gas electron configuration? - one with 10 d-electrons.
Why are pseudo-noble gas configurations considered relatively stable?
The filled d subshell represents a stable electron arrangement due to the complete pairing of electrons and the inherent stability associated with filled orbitals. This stability is not as robust as true noble gas configurations, but still significant.
Are elements with pseudo-noble gas configurations always unreactive?
No. While the d¹⁰ configuration is stable, these ions can still participate in chemical reactions. Their reactivity depends on factors like charge density and the polarizing ability of the ion. Remember, even with a filled d subshell, these elements can still form complexes.
So, after all that electron configuration talk, hopefully you have a better grasp on which of the following is a pseudo-noble-gas electron configuration? Keep experimenting, and see what other interesting electron arrangements you can find!