1. Home
  2. >Chemistry
  3. >The Atom
Help

Electronic structure: Subshells

Get the most by viewing this topic in your current grade. Pick your course now.

?
Intros
Lessons
  1. Building on electron structure
  2. Electron Shells: Recap and introduction to sub-shells.
  3. The evidence for electron shells or 'quantum' energy levels.
  4. Electron subshells.
  5. Evidence for electron subshells
  6. The energy level diagram.
  7. Subshell notation: Walkthrough.
  8. Noble gas notation: Walkthrough.
?
Examples
Lessons
  1. Apply knowledge of electron subshells to write electron configuration.
    Write the electron configuration of the following species using subshell notation:
    1. B
    2. Ca
    3. F−^-
    4. Mg2+^{2+}
  2. Apply knowledge of energy level diagrams to write subshell electron configuration.
    Write the electron configuration of the following species using an energy level diagram:
    1. Al
    2. Ne
    3. O2−^{2-}
    4. As
  3. Apply knowledge of core notation to write subshell electron configuration.
    Write the electron configuration of the following species using noble gas or 'core' notation:
    1. Be
    2. N3−^{3-}
    3. Br−^-
    4. Se
Topic Notes
?

Introduction to Electronic Structure: Subshells

Welcome to our exploration of electron subshells, a crucial concept in understanding atomic structure! Subshells are the building blocks of electronic configuration, helping us visualize how electrons are arranged within an atom. In the introduction video, we'll dive into this fascinating topic, breaking it down in a way that's easy to grasp. You'll learn about the different types of subshells (s, p, d, and f) and how they relate to an atom's energy levels. Understanding subshells is key to predicting an element's chemical properties and behavior. As we progress, you'll see how this knowledge forms the foundation for more advanced concepts in chemistry and physics. The video will use clear examples and visual aids to make these abstract ideas more concrete. By the end, you'll have a solid grasp of subshells and their significance in atomic structure. Let's embark on this exciting journey together!

Understanding Electron Subshells

Electron subshells are a fundamental concept in chemistry that helps us understand the complex structure of atoms. To grasp what a subshell is, we first need to revisit the basic shell model of an atom. In this model, electrons orbit the nucleus in distinct energy levels or shells. However, this simplified view doesn't fully explain the intricate behavior of electrons. This is where subshells come into play.

So, what is a subshell? A subshell is a subdivision of an electron shell that provides a more detailed description of electron arrangement and energy states within an atom. Think of subshells as specialized compartments within each main shell, each with its own unique properties and capacity for electrons.

There are four types of subshells, designated by the letters s, p, d, and f. Each type has a distinct shape and capacity for electrons:

  • The s subshell is spherical and can hold up to 2 electrons.
  • The p subshell has a dumbbell shape and can accommodate 6 electrons.
  • The d subshell is more complex, resembling a four-leaf clover, and can hold 10 electrons.
  • The f subshell has an even more intricate shape and can contain 14 electrons.

To better understand subshells, let's use an analogy. Imagine an apartment building where each floor represents a main shell. On each floor, there are different types of apartments (subshells) - studio apartments (s), one-bedroom apartments (p), two-bedroom apartments (d), and luxury suites (f). Just as these apartments differ in size and capacity, subshells vary in their electron-holding capabilities and energy levels.

The relationship between subshells and energy levels is crucial in subshell chemistry. Within each main shell, subshells have slightly different energy levels. Generally, the energy increases in the order s < p < d < f. However, as we move to higher main shells, the energy levels of subshells can overlap between shells. This overlapping explains why the periodic table is arranged the way it is and why elements in the same group have similar properties.

Let's explore each subshell in more detail:

The s subshell is the simplest and lowest energy subshell. It's present in every main shell and is always filled first. Its spherical shape means electrons in an s subshell have an equal probability of being found anywhere at a given distance from the nucleus.

The p subshell appears from the second main shell onwards. Its dumbbell shape results in electrons being distributed in three possible orientations along the x, y, and z axes. This arrangement gives rise to the directional bonding properties of many elements.

The d subshell doesn't appear until the third main shell and is responsible for the unique properties of transition metals. Its complex shape allows for a variety of electron configurations, leading to the diverse colors and magnetic properties of these elements.

The f subshell, found in the inner transition elements (lanthanides and actinides), has the most complex shape. It's responsible for the special chemical and physical properties of these elements, including their use in powerful magnets and nuclear reactions.

Understanding subshells is key to grasping many aspects of chemical behavior. For instance, the way electrons fill these subshells (following the Aufbau principle, Hund's rule, and the Pauli exclusion principle) determines an element's electron configuration. This configuration, in turn, influences the element's chemical properties, bonding behavior, and position in the periodic table.

In conclusion, electron subshells provide a more nuanced view of atomic structure than the basic shell model. By dividing shells into s, p, d, and f subshells, we can better explain the complex behavior of electrons and the diverse properties of elements. Whether you're studying basic chemistry or delving into quantum mechanics, a solid understanding of subshells is essential for comprehending the fascinating world of atoms and molecules.

Subshell Electronic Configuration

Understanding how electrons are arranged in subshells is crucial for grasping the fundamentals of atomic structure and chemical behavior. The arrangement of electrons follows specific principles that govern their distribution within an atom's atomic orbitals. In this section, we'll explore the key concepts of subshell electronic configuration, including the Aufbau principle, Hund's rule, and the Pauli exclusion principle.

Electrons occupy atomic orbitals, which are grouped into subshells. These subshells are designated by the letters s, p, d, and f. Each subshell has a specific number of orbitals: s has 1, p has 3, d has 5, and f has 7. The arrangement of electrons in these subshells follows a set of rules that ensure the most stable configuration for the atom.

The Aufbau principle is the primary guide for electron arrangement. This principle states that electrons fill orbitals in order of increasing energy. The order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p. This sequence ensures that lower-energy orbitals are filled before higher-energy ones.

Hund's rule complements the Aufbau principle by dictating how electrons occupy orbitals within a subshell. According to Hund's rule, electrons will occupy orbitals of equal energy individually before pairing up. This maximizes the number of unpaired electrons, resulting in a lower-energy, more stable configuration.

The Pauli exclusion principle adds another layer to electron arrangement. It states that no two electrons in an atom can have the same set of quantum numbers. In practical terms, this means that each orbital can hold a maximum of two electrons, and these electrons must have opposite spins.

Let's look at some examples of electron configurations for various elements:

1. Hydrogen (H): 1s¹
2. Helium (He): 1s²
3. Lithium (Li): 1s² 2s¹
4. Carbon (C): 1s² 2s² 2p²
5. Oxygen (O): 1s² 2s² 2p
6. Sodium (Na): 1s² 2s² 2p 3s¹
7. Chlorine (Cl): 1s² 2s² 2p 3s² 3p

To write electronic configurations using subshell notation, follow these steps:

1. Determine the total number of electrons for the atom (equal to the atomic number for neutral atoms).
2. Start with the 1s orbital and fill it with up to two electrons.
3. Move to the next subshell in the Aufbau order (2s, then 2p, etc.).
4. Fill each subshell with electrons, following Hund's rule for p, d, and f orbitals.
5. Continue until all electrons are placed.
6. Write the configuration using the notation nl^x, where n is the principal quantum number, l is the subshell letter, and x is the number of electrons in that subshell.

For example, let's write the electronic configuration for phosphorus (P, atomic number 15):

1. Phosphorus has 15 electrons.
2. Fill 1s: 1s²
3. Fill 2s: 1s² 2s²
4. Fill 2p: 1s² 2s² 2p
5. Fill 3s: 1s² 2s² 2p 3s²

Subshell Energy Levels and Order

Understanding the energy levels of subshells and their order of filling is crucial in grasping the structure of atoms and the organization of the periodic table. The arrangement of electrons in an atom follows specific patterns, but these patterns are not always as straightforward as one might expect.

Subshells are subdivisions of electron shells, denoted by the letters s, p, d, and f. Each subshell has a unique energy level and can hold a specific number of electrons: s (2), p (6), d (10), and f (14). The general order of filling these subshells, known as the Aufbau principle, typically follows the sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

Shell Subshells Max Electrons
1 1s 2
2 2s, 2p 8
3 3s, 3p, 3d 18
4 4s, 4p, 4d, 4f 32

However, this order is not always strictly followed, leading to some apparent anomalies. One notable example is the filling of the 4s subshell before the 3d subshell. This phenomenon occurs due to the complex interplay of various factors affecting electron energy levels, including nuclear charge, electron-electron repulsion, and quantum mechanical effects.

The reason for the 4s subshell filling before 3d lies in the concept of effective nuclear charge. As we move across a period in the periodic table, the nuclear charge increases, pulling electrons closer to the nucleus. This effect is more pronounced for s electrons, which can penetrate closer to the nucleus than d electrons. Consequently, the 4s orbital experiences a lower energy state than the 3d orbital, despite being in a higher principal quantum number shell.

This seemingly counterintuitive order of filling is reflected in the structure of the periodic table. The transition metals, which involve the filling of d orbitals, appear as a distinct block between the s-block and p-block elements. This arrangement showcases how the periodic table is not just a simple list of elements but a sophisticated representation of electron configurations and chemical properties.

The periodic table's organization closely mirrors the order of subshell filling. The s-block elements (Groups 1 and 2) correspond to the filling of s orbitals. The p-block (Groups 13-18) represents the filling of p orbitals. The d-block (transition metals) and f-block (lanthanides and actinides) correspond to the filling of d and f orbitals, respectively. This structure allows chemists and physicists to predict the electronic behavior and chemical properties of elements based on their position in the table.

It's important to note that while the general trend of subshell filling follows the Aufbau principle, there are exceptions. Elements like chromium (Cr) and copper (Cu) exhibit electron configurations that deviate from the expected pattern due to the stability gained from having half-filled or fully filled d subshells. These exceptions highlight the complexity of atomic structure and the limitations of simplified models.

Understanding subshell energy levels and their filling order is essential for predicting chemical behavior, explaining spectroscopic observations, and comprehending the periodic trends in element

Evidence for Subshells

The concept of subshells in atomic structure is fundamental to our understanding of modern chemistry. Experimental evidence supporting the existence of subshells has been crucial in developing the current model of atomic structure. This evidence comes primarily from spectroscopic data, ionization energy patterns, and other sophisticated experiments.

Spectroscopic data has been one of the most significant sources of evidence for subshells in atomic structure. When atoms are excited, they emit light at specific wavelengths, creating unique spectral lines. These lines, when analyzed closely, reveal fine structure that cannot be explained by a simple shell model. The splitting of spectral lines, known as fine structure splitting, provides direct evidence for the existence of subshells. For instance, the hydrogen spectrum shows a splitting of lines that can only be accounted for by the presence of subshells with different angular momentum quantum numbers.

Ionization energy patterns have also played a crucial role in supporting the subshell model. Experiments measuring the energy required to remove electrons from atoms have revealed patterns that align perfectly with the subshell structure. As electrons are removed from an atom, there are distinct jumps in ionization energy that correspond to the completion of subshells. These jumps are particularly noticeable when moving from one principal energy level to another, but smaller variations within a principal level provide evidence for subshells.

X-ray spectroscopy has been another valuable tool in confirming the existence of subshells. When high-energy electrons bombard atoms, they can knock out inner-shell electrons, causing outer electrons to fill these vacancies. This process results in the emission of X-rays with characteristic energies. The fine structure observed in these X-ray spectra provides additional evidence for the subshell model, as it reflects the energy differences between various subshells.

Electron spin resonance (ESR) spectroscopy has also contributed to our understanding of subshells. This technique detects unpaired electrons in atoms and molecules, providing information about the electronic structure. The patterns observed in ESR spectra are consistent with the subshell model and have helped refine our understanding of electron configurations.

The photoelectric effect, while primarily known for its role in quantum theory, has also provided indirect evidence for subshells. The energy of ejected electrons from metals exposed to light of varying frequencies shows discrete jumps that correspond to electrons being removed from different subshells.

Atomic beam experiments, such as the Stern-Gerlach experiment, have demonstrated the quantization of angular momentum, which is a key feature of the subshell model. These experiments showed that electrons possess intrinsic angular momentum (spin) in addition to orbital angular momentum, both of which are fundamental to the concept of subshells.

The development of the subshell model was a gradual process that evolved as new experimental evidence emerged. Early atomic models, such as Bohr's model, could not fully explain the observed spectral patterns. The introduction of quantum mechanics and the concept of electron orbitals by Schrödinger and others provided the theoretical framework for subshells. However, it was the accumulation of experimental evidence that solidified the subshell model as an accurate representation of atomic structure.

In conclusion, the existence of subshells in chemistry is supported by a wealth of experimental evidence. Spectroscopic data, ionization energy patterns, X-ray spectra, ESR, photoelectric effect studies, and atomic beam experiments have all contributed to our understanding of subshells. This evidence has been crucial in developing and refining the subshell model, which is now an integral part of our understanding of atomic structure and chemical behavior. The subshell model not only explains observed phenomena but also provides a framework for predicting and understanding complex atomic and molecular interactions in chemistry and related fields.

Applications and Importance of Subshell Theory

Subshell theory in chemistry plays a crucial role in understanding atomic structure and predicting chemical behavior. This fundamental concept has far-reaching applications in various fields, from materials science to spectroscopy. The practical implications of subshell theory are evident in how it helps scientists and researchers interpret chemical bonding, reactivity, and spectroscopic data.

One of the primary applications of subshell theory is in predicting chemical bonding. By understanding the arrangement of electrons in subshells, chemists can anticipate how atoms will interact and form bonds. For instance, the concept of valence electrons, which are found in the outermost subshells, is essential in determining an element's reactivity and bonding capabilities. This knowledge is invaluable in fields such as materials science, where researchers design new compounds with specific properties based on their electronic configurations.

In spectroscopy, subshell theory is indispensable for interpreting atomic and molecular spectra. The energy transitions between different subshells give rise to characteristic spectral lines, which serve as fingerprints for identifying elements and compounds. For example, flame tests used to identify metal ions rely on the unique colors emitted when electrons in specific subshells are excited and then return to their ground state. More advanced spectroscopic techniques, such as X-ray photoelectron spectroscopy (XPS), directly probe the energies of electrons in different subshells, providing valuable information about chemical environments and oxidation states.

The concept of subshells is also crucial in understanding the periodic table and trends in chemical properties. The periodic arrangement of elements is a direct result of the progressive filling of subshells, which explains patterns in atomic size, ionization energy, and electronegativity. This understanding is fundamental in predicting the behavior of elements and compounds in various chemical reactions and processes.

Real-world applications of subshell theory extend to many technological advancements. In the semiconductor industry, the design of electronic devices relies heavily on manipulating the electron configurations in materials. The behavior of electrons in different subshells determines the conductivity and other electronic properties of semiconductors, which are essential for creating transistors, solar cells, and other electronic components.

Moreover, subshell theory is vital in understanding complex phenomena such as magnetism and color in transition metal compounds. The partially filled d-subshells of transition metals give rise to unique magnetic and optical properties, which are exploited in various applications, from MRI machines to colorful pigments used in art and industry.

In conclusion, subshell theory is not just an abstract concept but a powerful tool with wide-ranging practical applications. Its importance in predicting chemical behavior, understanding bonding, and interpreting spectroscopic data makes it an essential foundation for numerous scientific and technological advancements. From developing new materials to analyzing chemical compounds, the insights provided by subshell theory continue to drive innovation and discovery in chemistry and related fields.

Conclusion

Understanding atomic structure is crucial for grasping atomic structure and electronic configuration. The introduction video provides a solid foundation for this complex topic, illustrating how electrons occupy different energy levels and subshells within an atom. Key points include the arrangement of subshells (s, p, d, f), their varying capacities, and the Aufbau principle governing electron filling. This knowledge is fundamental to predicting chemical behavior and bonding. We encourage you to delve deeper into atomic structure, exploring concepts like orbital shapes and electron spin. Remember, mastering electron subshells opens doors to understanding periodic trends, chemical reactivity, and spectroscopy. As you continue your chemistry studies, apply this knowledge to solve problems and predict atomic properties. Challenge yourself to draw electronic configurations for various elements and ions. By doing so, you'll reinforce your understanding and develop critical thinking skills essential for success in chemistry. Keep exploring and applying these concepts in your academic journey!

Building on Electron Structure

Electron Shells: Recap and Introduction to Sub-shells

Step 1: Introduction to Electron Configuration and Structure

In this lesson, we delve deeper into the electron configuration and structure beyond the basic understanding. Initially, we have learned about the 2AA rule, which provides a fundamental introduction to electron configuration. Now, we will explore the concept of electron subshells, which adds more detail and complexity to our understanding of electron structure.

Step 2: Understanding Electron Subshells

Electron subshells are subdivisions within electron shells. Each shell can contain one or more subshells, and these subshells are designated by the letters s, p, d, and f. The types of subshells and their arrangement within shells are crucial for understanding the distribution of electrons in an atom.

Step 3: Evidence for Electron Shells and Subshells

Scientific theories require experimental evidence to be validated. The existence of electron shells and subshells is supported by various experiments and observations. These experiments provide the necessary evidence to confirm the theoretical models of electron configuration.

Step 4: Energy Level Diagrams

Energy level diagrams are a visual representation of the electron structure of an atom. These diagrams help in understanding the arrangement of electrons in different shells and subshells. Constructing these diagrams involves placing electrons in the appropriate energy levels and subshells according to specific rules.

Step 5: Subshell Notation

Subshell notation is a method of writing electron configurations that provides more detail than the basic 2AA rule. This notation includes the designation of subshells and the number of electrons in each subshell. Additionally, core notation or noble gas notation is another way to represent electron configurations, which simplifies the notation by using the electron configuration of the nearest noble gas as a reference point.

Step 6: Recap of the 2AA Rule

The 2AA rule states that the first electron shell can hold up to two electrons, which corresponds to the first row of the periodic table (hydrogen and helium). The second shell can hold up to eight electrons, corresponding to the second row of the periodic table. This rule provides a basic framework for understanding electron configuration but does not account for the differences between electrons in the same shell.

Step 7: Limitations of the 2AA Rule

While the 2AA rule is useful for a basic understanding, it has limitations. It suggests that all electrons in a shell are the same, which is not accurate. In reality, electrons in the same shell can occupy different subshells, which have different energy levels and properties. This discrepancy is addressed by the concept of subshells and orbitals.

Step 8: Introduction to Orbitals

Orbitals are regions within subshells where electrons are likely to be found. Each type of subshell (s, p, d, f) contains a specific number of orbitals, and each orbital can hold up to two electrons. Understanding orbitals is essential for accurately describing the electron configuration of atoms.

Step 9: Filling Electron Subshells

The process of filling electron subshells follows specific rules, such as the Aufbau principle, Hund's rule, and the Pauli exclusion principle. These rules dictate the order in which electrons fill subshells and orbitals, ensuring that the electron configuration is correctly represented.

Step 10: Summary and Next Steps

This lesson provides a comprehensive overview of electron configuration, subshells, and orbitals. By understanding these concepts, we can accurately describe the electron structure of atoms and predict their chemical behavior. In the next sections, we will explore these topics in more detail and apply our knowledge to various elements and compounds.

FAQs

Here are some frequently asked questions about electron subshells:

1. What is a subshell in simple words?

A subshell is a subdivision of an electron shell in an atom. It represents a group of orbitals with the same energy level and similar shapes. Subshells are designated by the letters s, p, d, and f, each with a specific capacity for electrons.

2. What is the difference between an orbital and a subshell?

An orbital is a region in an atom where an electron is likely to be found, while a subshell is a group of orbitals with the same energy level. For example, a p subshell contains three p orbitals.

3. How many electrons are in each subshell?

The number of electrons in each subshell varies: s subshells can hold 2 electrons, p subshells can hold 6, d subshells can hold 10, and f subshells can hold 14 electrons.

4. What is the meaning of 1s, 2s, 2p, 3s, 3p?

These notations describe specific subshells. The number indicates the principal quantum number (energy level), and the letter represents the subshell type. For example, 2p refers to the p subshell in the second energy level.

5. What is the order of energy in the subshells?

The general order of increasing energy for subshells is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p. However, there can be some variations in this order for specific elements.

Prerequisite Topics

Understanding the electronic structure of atoms, particularly subshells, is a crucial concept in chemistry. However, to fully grasp this topic, it's essential to have a solid foundation in several prerequisite areas. These foundational concepts provide the necessary context and background knowledge to comprehend the intricacies of subshells within electronic structures.

One of the most fundamental prerequisites is the understanding of atomic orbitals and energy levels. This concept is vital because subshells are essentially groupings of these atomic orbitals. By grasping how electrons occupy different energy levels and orbitals, students can better visualize the arrangement of electrons in subshells. The shapes and energies of atomic orbitals directly influence the structure and properties of subshells, making this prerequisite topic indispensable.

Another crucial prerequisite is the history and development of the periodic table. The periodic table's arrangement is intrinsically linked to the electronic structure of atoms, including subshells. Understanding how scientists discovered and organized elements based on their properties provides valuable insights into why subshells are structured the way they are. The periodic table's layout reflects the filling of electron subshells, making this historical context essential for a deeper comprehension of electronic structure.

Additionally, knowledge of periodic trends, particularly ionization energy, is highly relevant to the study of subshells. Ionization energy patterns across the periodic table are directly related to the electronic configuration of atoms, including the arrangement of electrons in subshells. By understanding how ionization energy changes across periods and down groups, students can better appreciate the stability and reactivity of different electronic configurations, which is closely tied to subshell structure.

These prerequisite topics form a strong foundation for delving into the complexities of electronic structure and subshells. Atomic orbitals provide the basic building blocks, the periodic table offers a framework for organizing this information, and ionization energy trends demonstrate the practical implications of electronic structure. By mastering these concepts, students will be well-equipped to tackle the intricacies of subshells, understanding not just the 'what' but also the 'why' behind electronic structures.

In conclusion, the journey to understanding electronic structure and subshells is paved with these essential prerequisite topics. Each concept builds upon the others, creating a comprehensive understanding of atomic structure and behavior. By recognizing the importance of these foundational ideas, students can approach the study of subshells with confidence, ready to explore the fascinating world of electronic configurations and their far-reaching implications in chemistry and beyond.

In this lesson, we will learn:
  • To understand energy level diagrams of electron shells.
  • How to construct energy level diagrams using electron subshells.
  • How to write electron configurations using electron subshell notation.
  • How to write electron configurations using noble gas or 'core' notation.
  • How an element’s electron configuration relates to its chemical properties.

Notes:
  • In Electronic Configuration 1 a basic electron structure was introduced with the 2-8-8 rule. This is useful only for the first 3 rows of the periodic table before the transition metals and it doesn’t look at electron subshells which we need for later elements.

  • Evidence that electrons exist in quantum shells comes from the results of atomic emission spectroscopy (AES). When energized, samples of elements like atomic hydrogen give off photons of a fixed wavelength – we call these its spectral lines and they act like a fingerprint for the element. These specific photon wavelengths and energies are evidence that the electrons shifted directly from a state of one specific (quantized) amount of energy to another state of specific quantized energy, which is why they are called quantum shells or energy levels.
    Ionization energies (see Periodicity: Ionization energies) are also evidence of quantum levels because of the distinct gaps in energy required at regular intervals as you progress through the periodic table in order of atomic number. For now it is ok, but try to replace ‘shell’ with ‘energy level’ when talking about where electrons are. We will learn why later.

  • Each energy level (shell) is comprised of orbitals (subshells). We cover orbitals in detail in Introduction to atomic orbitals and energy levels, but you can think of them as shells within a shell or energy level. An orbital is a region of space which can hold up to two electrons each, and there are a few types of orbitals:

    • For the first energy level (n=1n=1), one s-orbital exists which contains up to 2 electrons. S-orbitals are spherical shaped; think “s for sphere”. Helium, which has two electrons and therefore a full first energy level, has the electron configuration “1s2”.

    • In the second energy level (n=2n=2), three p-orbitals also exist, which can contain up to 6 electrons. P-orbitals are lobe shaped; each p-orbital is two lobes running along an axis in the opposite direction.
      With three p-orbitals and one s-orbital, the second energy level holds a total of 8 electrons, so an atom with full first and second energy levels would have the electron configuration “1s2 2s2 2p6”.

    • In the third energy level (n=3n=3), five d-orbitals also exist which can contain up to 10 electrons. This gives the third energy level a total of 18 electrons.
      An atom with completed first, second and third energy levels would have the electron configuration “1s2 2s2 2p6 3s2 3d10 3p6”.

    • There is also a fourth energy level that contains f-orbitals which is beyond the scope of chemistry courses at this level.

    • Orbitals fill with electrons in a specific order according to the energy level diagram, shown below.

  • While the evidence for quantized shells comes largely from AES, the evidence for electron subshells largely comes from photoelectron spectroscopy (PES).
    • PES is an analytical tool where photons bombard electrons in an elemental sample and the kinetic energy of the electrons ejected is found. By knowing the energy of the photons used (E=hvE = hv) and measuring the ejected electrons’ kinetic energy, we can work out the binding energy of electrons to the atoms when they were in the sample. This is how much energy it took to ‘kick out’ the electrons from whatever state they were in.
    • In any given elemental sample, the electrons’ binding energy show distinct gaps. This supports the idea of quantized electron energy levels with fixed energy gaps between them.
      But where the binding energies appear, there are often groups of different binding energies. These are the s, p, d subshells or ‘shells within shells’ better known as orbitals.
    PES is one method of finding ionization energies that help us understand atomic structure, and is explored in more detail in our lesson Photoelectron spectroscopy.

  • Electron configurations in an atom can be shown in three ways.
    • You can use an energy level diagram, just by filling in the labelled orbitals (represented as lines) from lowest energy to highest. The energy level diagram is shown below. It reflects a few pieces of important information that is essential to writing electron configurations correctly:

energy level diagram

      • An s-orbital (say, 3s) is lower energy than a same level p-orbital (3p), which is lower energy than a d-orbital (3d).
      • Electrons always fill lower energy subshells before higher energy subshells. For example, 2s will always fill before 2p. The energy-level diagram shown above has the correct order of subshells.
      • Electrons must fill orbitals singly first and only pair up after this. They must pair up with ‘opposite spin’; one pointing up, one pointing down.
      • The gaps between increasing energy levels gets increasingly small, which is why 3d is higher energy and fills up after 4s, which is why it is the 2-8-8 rule and not 2-8-18!
      • Up to including the f-orbitals, the shells fill with the following number of electrons: 2 (1s); 8 (2s, 2p); 8 (3s, 3p); 18 (4s, 3d, 4p); 18 (5s, 4d, 5p); 32 (6s, 4f, 5d, 6p).

    • You can also show electron configuration by writing in subshell notation.
      • Subshell notation does not require drawing an energy level diagram. It is written in the format nxy where nn is the energy level, xx is the subshell and yy is the number of electrons in that subshell. For example: 2p4 would mean there are four electrons in the 2nd level p-subshell.
      • Like the energy-level diagram, the lowest energy subshells is written first. For example, the electron configuration for carbon is written: 1s2, 2s2 2p2.

    • There is another way to write electron configuration, known as noble gas notation or core notation.
      • This uses the chemical symbols of noble gases in square brackets, and only writes the subshells for the highest incomplete energy level. Noble gas symbols are used because their electron configuration are full shells; it is writing a shorthand.
      • For example: [Ar] has the configuration (2,8,8) so the configuration of K could be written: [Ar] 4s1. Other than this shortcut, core notation is very similar to regular subshell notation.

  • To write the configuration for ions, see whether the ion is negative (it has gained an electron) or positive (it has lost an electron) and add or remove electrons as needed. If a subshell was the last to fill up, it is the first to empty! The only exception to this is with the 3d subshell.

  • The highest (partially) filled subshell for an element is how we say what “block” an element is in.
    • For example, lithium’s electron configuration of 1s2 2s1 shows its highest energy electron in an s-orbital, so it is an s-block element.
    • Likewise, boron is called a p-block element because its configuration of 1s2 2s2 2p1 shows its highest energy electron in a p-orbital.
    Because the highest energy electrons are what make bonds and take part in chemical reactions, an element’s electron configuration is what determines its chemical properties.