Atomic orbitals and energy levels
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Introduction to Atomic Orbitals and Energy Levels
Atomic orbitals and energy levels are fundamental concepts in understanding the structure of atoms and electron behavior. The introduction video provides a crucial foundation for grasping these complex ideas. Quantum theory revolutionized our understanding of atomic structure, revealing that electrons don't orbit the nucleus like planets around the sun, but exist in probability clouds called orbitals. These orbitals represent regions where electrons are likely to be found and are characterized by specific energy levels. The quantum mechanical model explains how electrons occupy these energy levels and transition between them, giving rise to atomic spectra and chemical properties. This understanding is essential for explaining chemical bonding, spectroscopy, and various phenomena in materials science. By delving into atomic orbitals and energy levels, we gain insight into the microscopic world that shapes the macroscopic properties of matter, highlighting the profound impact of quantum theory on our comprehension of nature at its most fundamental level.
The Evolution of Atomic Models
The historical development of atomic models represents a fascinating journey through scientific discovery, challenging our understanding of the fundamental building blocks of matter. From early concepts to our current understanding of atomic orbitals, this evolution has been marked by groundbreaking experiments and revolutionary theories that have reshaped our view of the physical world.
The story begins with ancient Greek philosophers who first proposed the concept of atoms as indivisible units of matter. However, it wasn't until the 19th century that scientific evidence began to support atomic theory. John Dalton's atomic model in the early 1800s laid the foundation for modern atomic theory, proposing that all matter is composed of indivisible particles called atoms.
A significant leap forward came with J.J. Thomson's discovery of electrons in 1897. His "plum pudding" model suggested that atoms were composed of negatively charged electrons embedded in a positively charged "pudding." This model, while revolutionary for its time, was soon challenged by new experimental evidence.
Ernest Rutherford's gold foil experiment in 1909 led to a dramatic shift in our understanding of atomic structure. His findings resulted in the nuclear model of the atom, with a small, dense, positively charged nucleus surrounded by electrons. This model, however, couldn't explain the stability of atoms or the discrete emission spectra of elements.
The introduction of quantum concepts began to address these shortcomings. Max Planck's quantum theory, proposed in 1900, suggested that energy is emitted or absorbed in discrete packets called quanta. This revolutionary idea laid the groundwork for the quantum revolution in physics.
Building on Planck's work, Albert Einstein explained the photoelectric effect in 1905. This phenomenon, where light causes electrons to be emitted from a material, could only be explained by treating light as discrete particles (photons) rather than continuous waves. Einstein's work on the photoelectric effect, which later earned him the Nobel Prize, was crucial in establishing the particle nature of light and the wave-particle duality concept.
In 1913, Niels Bohr introduced his atomic model, which was a significant step towards our modern understanding of atomic structure. Bohr's model proposed that electrons orbit the nucleus in fixed energy levels, and that electrons can jump between these levels, emitting or absorbing specific amounts of energy. This model successfully explained the hydrogen spectrum and introduced the idea of quantized energy levels in atoms.
While Bohr's model was a major advancement, it still had limitations. The development of quantum mechanics in the 1920s by scientists like Werner Heisenberg, Erwin Schrödinger, and others led to our current understanding of atomic orbitals. This quantum mechanical model describes electrons not as particles in fixed orbits, but as probability distributions around the nucleus, forming orbitals of various shapes.
These developments in atomic theory fundamentally challenged classical physics. The deterministic view of the universe gave way to probabilistic interpretations. Concepts like wave-particle duality, the uncertainty principle, and quantum superposition emerged, revolutionizing our understanding of the microscopic world.
The journey from early atomic concepts to our current understanding of atomic orbitals illustrates the dynamic nature of scientific progress. Each model and theory built upon its predecessors, addressing limitations and incorporating new experimental evidence. This evolution not only transformed our view of atoms but also laid the foundation for numerous technological advancements, from semiconductors to medical imaging techniques.
Today, our understanding of atomic structure continues to evolve. Advanced experimental techniques and computational models allow scientists to probe deeper into the nature of matter and energy. The legacy of Planck, Einstein, Bohr, and other pioneers in quantum theory continues to influence modern physics, chemistry, and materials science, driving innovation and expanding our knowledge of the universe at its most fundamental level.
Understanding Electron Orbitals
Electron orbitals are a fundamental concept in modern chemistry and quantum mechanics, representing the three-dimensional regions around an atom's nucleus where electrons are most likely to be found. Unlike the earlier model of electron shells, which depicted electrons as particles orbiting the nucleus in fixed circular paths, the orbital model provides a more accurate and complex understanding of electron behavior.
The concept of electron orbitals is closely tied to the probabilistic nature of electron locations. Instead of precise, deterministic paths, electrons exist in probability distributions known as electron clouds. These clouds represent areas where an electron has a high likelihood of being present at any given moment. The shape and size of these clouds vary depending on the energy level and type of orbital.
One of the key reasons we can't pinpoint the exact position of electrons is explained by the Heisenberg Uncertainty Principle. This principle, formulated by Werner Heisenberg in 1927, states that it's impossible to simultaneously determine both the precise position and momentum of a particle, such as an electron, with absolute certainty. The more accurately we measure one property, the less accurately we can measure the other. This inherent uncertainty at the quantum level is a fundamental aspect of our understanding of electron behavior.
Electron orbitals differ significantly from the earlier concept of electron shells. While shells were visualized as concentric spheres around the nucleus, orbitals have various shapes and orientations. The main types of orbitals are s, p, d, and f, each with distinct characteristics. S orbitals are spherical, p orbitals are dumbbell-shaped, and d and f orbitals have more complex geometries. These shapes are determined by mathematical wave functions that describe the probability of finding an electron in a particular region of space.
The orbital model is crucial in explaining chemical bonding and molecular shapes. When atoms form chemical bonds, their orbitals interact and overlap. This interaction determines the type and strength of the bond formed. For example, in covalent bonding, the overlap of orbitals from different atoms allows electrons to be shared, creating stable molecules. The specific orbitals involved in bonding influence the geometry of the resulting molecule.
Understanding orbitals helps predict and explain molecular shapes through theories like Valence Shell Electron Pair Repulsion (VSEPR). This theory considers how the arrangement of electron pairs around a central atom, both bonding and non-bonding, determines the three-dimensional structure of molecules. The shapes of molecules, in turn, affect their properties and reactivity.
The concept of hybridization, where atomic orbitals combine to form new hybrid orbitals, further explains bonding in molecules like methane (CH4) or ethene (C2H4). Hybridization allows atoms to form bonds that wouldn't be possible with pure atomic orbitals, expanding our understanding of molecular structures and chemical reactions.
Electron orbitals also play a crucial role in spectroscopy and the study of atomic and molecular energy levels. When electrons transition between different energy levels or orbitals, they absorb or emit specific wavelengths of light. This phenomenon forms the basis for various spectroscopic techniques used in chemical analysis and astronomical observations.
In conclusion, the concept of electron orbitals represents a significant advancement in our understanding of atomic structure and chemical behavior. By embracing the probabilistic nature of electron locations and the principles of quantum mechanics, this model provides a powerful framework for explaining and predicting a wide range of chemical phenomena. From the formation of chemical bonds to the shapes of molecules and the interaction of matter with light, electron orbitals continue to be a cornerstone of modern chemistry and physics.
Types of Orbitals: S and P Orbitals
In the realm of atomic structure, s and p orbitals play a crucial role in determining the behavior and properties of elements. These orbitals, which represent regions of high electron probability, exhibit distinct characteristics that influence chemical bonding and reactivity. Understanding the nature of s and p orbitals is essential for grasping the fundamentals of atomic theory and quantum mechanics.
S orbitals are characterized by their spherical shape, which is symmetrical in all directions around the nucleus. This unique geometry allows s orbitals to have an equal probability of finding an electron at any given distance from the nucleus within the orbital. The spherical nature of s orbitals makes them the simplest type of atomic orbital, with no angular nodes. As we move up in energy levels, the size of the s orbital increases, but it maintains its spherical shape.
In contrast, p orbitals have a distinctive dumbbell shape, consisting of two lobes extending in opposite directions from the nucleus. This shape results from the presence of one angular node, which is a region where the probability of finding an electron is zero. The dumbbell shape of p orbitals allows for directional bonding, which is crucial in forming molecular structures and determining the geometry of compounds.
The concept of nodes is fundamental in understanding the structure of atomic orbitals. Nodes are regions within an orbital where the probability of finding an electron is zero. S orbitals have no angular nodes but may have radial nodes in higher energy levels. P orbitals, on the other hand, always have one angular node, which gives rise to their characteristic dumbbell shape. The number and type of nodes in an orbital are directly related to its energy and quantum numbers.
Energy levels and sublevels are intimately connected to the structure of s and p orbitals. The principal quantum number (n) determines the main energy level, while the angular momentum quantum number (l) defines the sublevel. S orbitals belong to the l=0 sublevel, while p orbitals correspond to the l=1 sublevel. As we move up in energy levels, the number of available sublevels increases, allowing for more complex electron configurations.
To illustrate how electrons fill these orbitals, let's consider some examples. Hydrogen, the simplest atom, has one electron occupying the 1s orbital in its ground state. This single s orbital provides a spherical distribution of electron density around the nucleus. Helium, with two electrons, fills the 1s orbital completely, resulting in a stable noble gas configuration.
Lithium introduces the concept of electron shells and subshells. Its electron configuration is 1s² 2s¹, where two electrons occupy the inner 1s orbital, and the third electron resides in the 2s orbital. This arrangement demonstrates how electrons fill lower energy orbitals before occupying higher ones, following the Aufbau principle.
As we move to elements with more electrons, p orbitals come into play. For instance, carbon has the electron configuration 1s² 2s² 2p². Here, we see the first appearance of p orbitals, with two electrons distributed among the three available 2p orbitals (2px, 2py, and 2pz). This configuration illustrates how p orbitals contribute to the valence shell and influence chemical bonding.
The shapes and characteristics of s and p orbitals have profound implications for chemical behavior. The spherical nature of s orbitals allows for uniform electron distribution, making them ideal for forming single bonds. P orbitals, with their directional lobes, are instrumental in forming multiple bonds and determining molecular geometry.
Understanding the properties of s and p orbitals is crucial for predicting and explaining various chemical phenomena. From the hybridization of orbitals in molecular bonding to the magnetic properties of elements, these fundamental concepts form the basis of modern chemistry and materials science. As we delve deeper into atomic structure, the interplay between s and p orbitals, along with d and f orbitals in larger atoms, reveals the intricate beauty of quantum mechanics at the atomic level.
Energy Levels and Electron Configuration
Energy levels in atoms are fundamental to understanding the behavior of electrons and the chemical properties of elements. These discrete energy states, also known as electron shells, represent the regions where electrons are most likely to be found around an atom's nucleus. Each energy level corresponds to a specific set of electron orbitals, which are three-dimensional spaces that describe the probable locations of electrons.
The concept of energy levels is closely tied to electron orbitals. Orbitals are classified by their energy and shape, denoted by quantum numbers. The main energy levels are represented by the principal quantum number (n), which can have values of 1, 2, 3, and so on. Within each main energy level, there are sublevels (s, p, d, f) that further define the electron's energy and spatial distribution.
The Aufbau principle is a key concept in understanding how electrons fill these orbitals. According to this principle, electrons occupy orbitals in order of increasing energy. This means that the lowest energy orbital is filled first, followed by the next lowest, and so on. The order of filling follows a specific pattern: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p.
Electron configuration is a shorthand notation used to describe the arrangement of electrons in an atom. It shows how electrons are distributed among the various orbitals. For example, the electron configuration of carbon (atomic number 6) is written as 1s² 2s² 2p². This indicates that carbon has two electrons in the 1s orbital, two in the 2s orbital, and two in the 2p orbital.
As we move across the periodic table, electron configurations change systematically. For instance, oxygen (atomic number 8) has the configuration 1s² 2s² 2p, while sodium (atomic number 11) is 1s² 2s² 2p 3s¹. These configurations reflect the filling of orbitals according to the Aufbau principle and help explain the periodic trends in element properties.
Energy levels play a crucial role in determining an atom's chemical properties. The outermost energy level, known as the valence shell, is particularly important. Electrons in this shell are involved in chemical bonding and reactions. For example, elements in the same group of the periodic table have similar valence electron configurations, leading to similar chemical behaviors. Noble gases, with their full outer shells, are chemically inert due to their stable electron configurations.
Understanding energy levels and electron configurations is essential for predicting an element's reactivity, ionization energy, and electron affinity. For instance, alkali metals easily lose their single valence electron, making them highly reactive, while halogens readily gain an electron to achieve a stable configuration. This knowledge forms the basis for explaining chemical bonding, spectroscopic properties, and many other phenomena in chemistry and physics.
In conclusion, energy levels and electron configurations are fundamental concepts in atomic structure. They provide a framework for understanding how electrons are arranged in atoms, which in turn determines the chemical and physical properties of elements. By applying the Aufbau principle and analyzing electron configurations, scientists can predict and explain a wide range of atomic behaviors and chemical interactions, making these concepts indispensable in the study of matter and its transformations.
Applications of Orbital Theory
Orbital theory, a cornerstone of modern chemistry, has numerous practical applications that extend far beyond theoretical discussions. This fundamental concept provides invaluable insights into chemical reactivity, bond formation, and molecular structure, making it an essential tool for chemists and researchers across various scientific disciplines.
One of the primary applications of orbital theory is in predicting chemical reactivity. By understanding the arrangement and energy levels of electrons in atomic and molecular orbitals, scientists can anticipate how different elements and compounds will interact. This knowledge is crucial in designing new materials, developing more efficient catalysts, and optimizing chemical reactions for industrial processes.
In the realm of bond formation, orbital theory offers a detailed explanation of how atoms combine to form molecules. The concept of orbital overlap and hybridization helps chemists understand the geometry of molecules and the strength of chemical bonds. This understanding is vital in fields such as drug design, where the precise arrangement of atoms in a molecule can significantly impact its effectiveness and interactions within biological systems.
Molecular structure prediction is another area where orbital theory shines. By analyzing the distribution of electrons in molecular orbitals, scientists can determine the shape and properties of complex molecules. This capability is particularly valuable in materials science, where the structure of a compound often dictates its physical and chemical properties, such as conductivity, reactivity, and optical characteristics.
Spectroscopy, a powerful analytical technique, relies heavily on the principles of orbital theory. The interaction between electromagnetic radiation and the electrons in atomic or molecular orbitals forms the basis for various spectroscopic methods. For instance, UV-visible spectroscopy utilizes the transitions of electrons between different energy levels to identify and quantify chemical species. Similarly, infrared spectroscopy exploits the vibrational modes of molecules, which are intimately linked to their molecular orbitals.
The interpretation of atomic spectra is another critical application of orbital theory. The distinct spectral lines observed when atoms are excited correspond to specific electronic transitions between orbitals. This phenomenon not only allows for the identification of elements in unknown samples but also provides insights into the electronic structure of atoms and ions. Astronomers use this principle to analyze the composition of distant stars and galaxies, demonstrating the far-reaching implications of orbital theory beyond laboratory settings.
Recent advancements in orbital theory applications have been driven by computational chemistry. Sophisticated quantum mechanical calculations now allow for more accurate predictions of molecular properties and reactivity. These computational methods have become indispensable in fields such as nanotechnology, where understanding the behavior of materials at the atomic level is crucial for developing new technologies.
Ongoing research in orbital theory focuses on expanding its applicability to more complex systems. For example, scientists are working on improving models for describing electron correlation in multi-electron systems, which is essential for accurately predicting the properties of transition metal complexes and other challenging molecular systems. Additionally, researchers are exploring the role of orbital theory in understanding and manipulating quantum states, which has potential applications in quantum computing and advanced materials design.
In conclusion, orbital theory's practical applications span a wide range of scientific and technological domains. From predicting chemical reactivity and molecular structure to interpreting spectroscopic data and designing new materials, this fundamental concept continues to be a powerful tool in the hands of chemists, physicists, and materials scientists. As research progresses, the applications of orbital theory are likely to expand further, driving innovations in fields as diverse as energy storage, environmental remediation, and quantum technologies.
Conclusion
Atomic orbitals and energy levels are fundamental concepts in understanding atomic structure. The introduction video provides a crucial visual representation of these complex ideas, helping to demystify quantum mechanics in atomic structure. Key points include the quantized nature of electron energy levels, the shapes and orientations of orbitals, and how electrons fill these orbitals according to specific rules. The Pauli Exclusion Principle and Hund's Rule play vital roles in determining electron configurations. To deepen your understanding, explore additional resources on quantum mechanics in atomic structure. These concepts form the foundation for many aspects of chemistry and physics, from chemical bonding to spectroscopy. Apply this knowledge in your studies of chemistry, materials science, and related fields to gain a more comprehensive understanding of matter at the atomic level. By mastering these concepts, you'll be better equipped to tackle advanced topics and contribute to scientific advancements in the future.
Example:
What are electron orbitals? What affects molecular/atomic properties?
Step 1: Introduction to Atomic Orbitals and Energy Levels
In this lesson, we will delve into the topic of atomic orbitals and energy levels. Depending on your prior knowledge or the sequence of lessons you have followed, you might already be familiar with atomic orbitals. This section serves as an introduction to the nature of atomic orbitals and the historical development of our understanding of them. The concept of what electrons are doing in atoms has evolved significantly since the first atomic models were proposed.
Step 2: Historical Development and Experimental Evidence
Our current understanding of atomic orbitals is based on a series of experimental evidence and theoretical developments. Initially, the models of the atom were quite different from what we understand today. Over time, through various experiments and observations, scientists have refined the model to include the concept of atomic orbitals. These orbitals help us understand the behavior of electrons in atoms and their role in chemical bonding.
Step 3: Shape and Nature of Electron Orbitals
Electron orbitals have specific shapes and energy levels. The lower energy electron orbitals have distinct shapes, which are crucial for understanding how electrons occupy these orbitals. The energy level diagrams help visualize the distribution of electrons in different orbitals. Understanding the shape and nature of these orbitals is essential for comprehending how atoms bond to form molecules.
Step 4: Filling Electron Orbitals
Filling electron orbitals follows specific rules and notations, such as subshell notation. This process determines the arrangement of electrons in an atom, which in turn affects the atom's chemical properties. The shape and nature of the orbitals play a significant role in this arrangement, influencing how atoms interact with each other.
Step 5: Role of Orbitals in Chemical Bonding
Orbitals are fundamental to understanding chemical bonding. They explain how atoms join together to form molecules. Various methods and theories have been developed to describe chemical bonding, with orbitals being a central concept. For example, the shape of molecules, such as the trigonal planar shape of borane (BH3) or the tetrahedral shape of methane (CH4), is determined by the arrangement of orbitals and the electrons within them.
Step 6: Influence of Electrons on Molecular Shape
The shape of a molecule is influenced by the electrons and their states. For instance, changing the central atom in a molecule can alter its shape due to the different electron configurations. The same number of electrons can result in different molecular shapes depending on the states of the electrons. This is evident in molecules like ammonia (NH3) and phosphine (PH3), where the arrangement of atoms is dictated by the electrons.
Step 7: Stability of Molecules and Atoms
The stability of molecules and atoms is also influenced by the electron configuration. For example, helium, with its two electrons, does not form stable molecules but is stable as an individual atom. The comparison of different atoms and molecules with the same number of electrons but different shapes highlights the role of electron states in determining molecular properties.
Step 8: Quantum Theory and Atomic Orbitals
Our understanding of atomic orbitals is rooted in quantum theory. This theory explains the behavior of electrons in atoms and how they occupy specific orbitals. The development of quantum theory and the experiments that led to the concept of orbitals are crucial for understanding the current model of the atom.
Step 9: Conclusion
In conclusion, atomic orbitals and energy levels are essential concepts for understanding the behavior of electrons in atoms and their role in chemical bonding. The shape and nature of orbitals, the arrangement of electrons, and the influence of electron states on molecular shape and stability are all critical factors that affect molecular and atomic properties. Our current understanding is based on a combination of experimental evidence and theoretical developments, particularly quantum theory.
FAQs
Here are some frequently asked questions about atomic orbitals and energy levels:
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What are the orbitals in order of energy level?
The orbitals in order of increasing energy level are: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. This order follows the Aufbau principle, which states that electrons fill orbitals from lowest to highest energy.
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What are the 4 types of orbitals?
The four main types of orbitals are s, p, d, and f. Each type has a distinct shape: s orbitals are spherical, p orbitals are dumbbell-shaped, d orbitals have a clover-leaf shape, and f orbitals have more complex geometries.
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How are orbitals related to energy levels?
Orbitals are directly related to energy levels. Each energy level (principal quantum number n) contains one or more sublevels (s, p, d, f), which in turn contain orbitals. As the energy level increases, so does the number of available sublevels and orbitals.
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Which orbital has the lowest energy?
The 1s orbital has the lowest energy. This is because it is closest to the nucleus, where the electron experiences the strongest attractive force from the positively charged protons.
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How do you find the energy level of an orbital?
The energy level of an orbital is determined by its principal quantum number (n). For example, in the notation 2p, the number 2 indicates the energy level. The letter (s, p, d, or f) represents the sublevel within that energy level.
Prerequisite Topics
Understanding atomic orbitals and energy levels is a fundamental concept in chemistry and quantum mechanics. While there are no specific prerequisite topics provided for this subject, it's important to recognize that a strong foundation in basic chemistry and physics principles can greatly enhance your comprehension of this complex topic.
Atomic orbitals and energy levels are at the heart of our understanding of atomic structure and electron behavior. To fully grasp these concepts, students should have a solid understanding of basic atomic structure, including the components of atoms (protons, neutrons, and electrons) and their arrangement. Familiarity with the periodic table and electron configuration is also beneficial, as these provide context for how electrons are distributed in different elements.
Additionally, a basic understanding of quantum mechanics can be incredibly helpful when delving into atomic orbitals and energy levels. The wave-particle duality of electrons and the principles of quantum mechanics form the basis for our modern understanding of electron behavior in atoms. While not strictly prerequisites, these concepts are closely intertwined with the study of atomic orbitals and energy levels.
Mathematical skills, particularly in algebra and basic calculus, can also be valuable when exploring this topic. Many of the principles governing atomic orbitals and energy levels are expressed through mathematical equations and models. A comfort with mathematical concepts can make it easier to understand and apply these principles.
Furthermore, knowledge of spectroscopy and how atoms interact with light can provide practical applications and real-world context for the study of atomic orbitals and energy levels. Understanding how electrons transition between energy levels and emit or absorb specific wavelengths of light is crucial in many areas of chemistry and physics.
While these topics are not listed as formal prerequisites, having a strong background in general chemistry, basic physics, and introductory quantum mechanics can significantly enhance your ability to understand and appreciate the complexities of atomic orbitals and energy levels. As you explore this fascinating subject, you'll find that it builds upon and integrates many fundamental concepts from various areas of science.
Remember, learning is a journey, and it's okay if you don't have a perfect understanding of all these related topics. The study of atomic orbitals and energy levels itself will help reinforce and expand your knowledge in these areas. Approach the subject with curiosity and patience, and don't hesitate to review foundational concepts as needed. This approach will not only help you grasp atomic orbitals and energy levels more effectively but also deepen your overall understanding of chemistry and physics.
In this lesson, we will learn:
- To understand the experimental evidence that led to our current understanding of electrons and orbitals.
- To recall the shape and nature of early electron orbitals and energy level diagrams.
- To fill electron orbitals for the first two rows of elements correctly according to the Pauli and Aufbau principles.
Notes:
- There are many methods we use to learn about chemical bonding, find out the shape and structure of compounds, but why do atoms even form molecules and compounds in the first place? What causes different molecules to have different shapes?
VSEPR helps us predict molecular shapes. Recall the shapes of some molecules from Molecular Geometry:- BH3 is trigonal planar, but NH3 and PH3, with the same number of atoms, don't form the same shape as BH3.
- CH4 is tetrahedral and isoelectronic molecules (same number of electrons) like NH3 and water H2O still form the same tetrahedral structure without four H atoms, using their lone electron pairs instead.
- Some atoms do not form molecules; for some reason they are stable enough to exist as individual atoms, like helium (He).
Clearly, it is not the number of atoms or number of electrons that determine molecule shape and other properties, it is the different states electrons are in. Today we use atomic orbitals to describe the behaviour of electrons in atoms.
- To understand why we use orbitals to describe electrons it is very helpful to know how science got here. Some extremely important experiments and developments were made in the early 20th century. All the developments below led to Nobel Prizes for the scientists contributions to quantum mechanics:
- In 1900, Max Planck derived Planck’s law, which solved the ultraviolet catastrophe. Classical theory predicted that the emission spectrum of a black body (think of the sun as a rough example) would amount to emitting infinite amounts of UV and other high frequency radiation.
This is obviously not true; humans wouldn’t exist if it was! By assuming energy is transferred and emitted in quantized amounts, Planck’s revision to the obviously flawed theory correctly fitted real observations. - In 1905, Einstein’s experiment, the photoelectric effect, was explained using Planck’s idea of quantization the Planck-Einstein relation, E = hv. Einstein explained that light travels as quantized packets of energy called photons.
Although light was long considered a wave, it also shows particulate properties. - In 1913, Niels Bohr used the quantization of energy to develop his model of the atom. Bohr could explain the atomic emission spectrum (AES) of hydrogen; the electron orbited the nucleus in quantized energy states. If it absorbed energy, an electron would be excited to a higher energy state, before returning to a ground state, emitting a quantized amount of energy (E = hv) in the process. These energy ‘states’ are the electron shells.
- In 1923, the de Broglie hypothesis proposed that like light, all matter including electrons has wave-particle nature. He developed the de Broglie wavelength formula = h/p, which applies to particles like electrons, relating momentum of a particle to wavelength. Later experiments of electrons diffracting through slits (a wave property) confirmed the wave properties of electrons.
- In 1927, Werner Heisenberg developed Heisenberg’s uncertainty principle. A simple consequence of wave-particle duality is that we cannot exactly know the location AND momentum of an electron at the same time.
- Waves have no definitive position; they are continuously spread through space. However, we can identify features of waves like wavelength and frequency, which is related to energy and momentum.
- Particles definitively exist in one position at a given time, but they do not have a wavelength and as such its momentum is unknown.
- In 1926, the Schrodinger equation was developed to describe the wave behaviour of electrons over time using wave functions.
- The most important point from the above is that drawing electrons as particles orbiting around a nucleus is WRONG. Showing electrons in shells ‘orbiting’ around the nucleus suggest they are single point charges (they aren’t) and we know exactly where they are (we don’t).
Electrons are better thought of as ‘clouds’ of density inside orbitals where they are most likely to be found. Orbitals are like electron houses; they aren’t always/definitely there, but they are most likely there (or more likely there than anywhere else). - All orbitals can hold two electrons, no more. This is part of the Pauli principle. This first 1s orbital, then, covers the electrons for hydrogen and helium, which you can show using an energy-level diagram or subshell notation.
- The first, lowest energy orbital is called a 1s orbital.
- 1 refers to an energy level; the lower it is, the more stable it is and the closer to the nucleus it is.
- S refers to shape – think s for sphere .
- After the first two electrons are filled in the 1s orbital, the third and fourth electrons are found in the 2s orbital. This is where the third electron of Li and third/fourth for Be are found.
- Again, s means spherical but 2 shows a higher energy level. This orbital is higher in energy than 1s because the electrons are normally further away from the nucleus.
- The positive region will, for example, overlap with positive parts of other orbitals (constructive interference).
- The positive part will cancel out negative parts of other orbitals (destructive interference).
- The 2nd energy level also has orbitals where the node is not spherical, but in a single plane. These are called p orbitals, and there are three in each energy level - one each for the x, y and z axes they occupy. They are considered “lobe shaped”. See below for images:
- With two electrons each, this gives the electron configurations for the elements boron through neon. See the diagram below for the energy level diagrams:
- The evidence for electron sublevels or subshells largely comes from photoelectron spectroscopy (PES), a development from the photoelectric effect. This is covered in our lesson Photoelectron spectroscopy.
Because this zero node separates the orbital into two halves, it is convention to draw these halves as opposing regions (red/blue, +/-, shaded/unshaded). This is how we try and show the part wave nature of the electrons in orbitals:
