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Atomic Orbital Hybridization: Foundation of Molecular Structure
Dive into the world of atomic orbital hybridization and revolutionize your understanding of molecular shapes and chemical bonding. Perfect for students seeking to excel in organic chemistry!

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Now Playing:Atomic orbital hybridization – Example 0a
Intros
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  1. Mixing atomic orbitals - hybridization.
  2. Same AOs, different bonding?
  3. sp3 hybrid orbitals and sigma bonds.
Atomic orbitals and energy levels
Jump to:NotesConceptExampleFAQsPrerequisites
Notes

In this lesson, we will learn:

  • To explain the bonding in simple molecules using atomic orbital hybridization.
  • To correctly draw sp3, sp2 and sp hybridized atomic orbitals.
  • To explain how atomic orbital hybridization is consistent with the observed properties of carbon-carbon bonds.
Notes:

  • In the lesson Atomic orbitals and energy levels, we learned that we use atomic orbitals (AOs) to explain bonding. We write electron configurations, like 1s2 2s2 2p2, to show electrons in different atomic orbitals that create molecules out of separate atoms by taking part in bonding.
    But atomic orbitals on their own don’t accurately reflect the bond energies and bond lengths in compounds. For example, carbon’s electron configuration is [1s2] 2s2 2p2. The two 2s electrons are paired up and two p electrons are unpaired in the higher energy 2p subshell. So it has four valence electrons and two are in a higher energy state than the other two.
    Does carbon actually make two bonds with a lone pair, like an analogue of NH3? Let’s look at two carbon compounds:
    • Methylene (:CH2) is a type of compound called a carbene. Carbenes have two electrons on carbon that are not shared; the other two are making covalent bonds. However, methylene is very reactive and unstable – most simple carbenes are.
    • In methane (CH4),carbon is bonded to four hydrogens and the bond angles and strengths are all equal. Methane is a stable compound that is easily stored at standard conditions.
    A simple comparison shows that compounds where carbon makes four bonds instead of two are generally far more common and stable.
    So how does carbon in CH4 make four bonds? How are the C-H bonds all equal energy and length? One current theory of bonding is that atomic orbitals can mix together or hybridize to become hybrid orbitals that are equal in energy. We call this idea atomic orbital hybridization. Like a hybrid animal or car, hybrid atomic orbitals are mixes of different atomic orbitals that show the characteristics of both.

  • Atomic orbital hybridization explains why in methane (CH4), carbon’s four valence electrons do not bond in this 2s2 2p2 ground state. If they did, there would be marked differences in bond energy. The four n=2 atomic orbitals (AOs) mix together to create four equivalent hybrid orbitals that are equal in energy.
    • This is done by a 2s electron being excited to a slightly higher 2p orbital, the one that is currently empty.
      This results in four singly occupied orbitals; three 2p orbitals equal in energy and one 2s orbital slightly lower in energy. Now carbon can form four single bonds.
    • These four orbitals mix together to create four hybrid orbitals of equal energy.
      In methane the one 2s orbital and three 2p orbitals create four sp3 hybrid orbitals. The character of the hybrids is related to the AOs that made them, so they are 25% ‘s character’ and 75% ‘p character’.
    • These sp3 hybrid orbitals can combine by overlapping with orbitals in other atoms (like the 1s AO of hydrogen) and form a sigma bond. Sigma bonds are formed by head-on orbital overlap.

    Like in VSEPR, the electrons in different orbitals repel one another so they will find maximum spacing around the atom, which for tetrahedral carbon is 109.5°. This explains the observation of equal C-H bond angles and bond strength in methane. These hybrids can be drawn as a mix of s and p orbitals, where they do still look similar to the p orbitals they have 75% of the character of, with the s character shown as one lobe of the p orbitals enlarged, the other shrunk (due to destructive/constructive overlap). See below for a diagram:

    Atomic orbitals
    Atomic orbitals

  • Hybridization can explain double bonds too, such as in ethene (C2H4). As with any scientific theory, hybridization must be able to explain the experimental evidence. Bond enthalpy data clearly shows C=C double bonds are less than twice as strong as two C-C single bonds. Something about the double bond orbital overlap is not as attractive in nature as the single bond overlap.
    • Ethene (C2H4) has its two carbon atoms bonded to three other atoms; two hydrogens each and the other carbon atom in the molecule. After a 2s electron is excited up to the empty 2p orbital, only the 2s and TWO 2p AOs need to be hybridized, not all three like in CH4. They are therefore called sp2 hybrid orbitals. They are one-third s character and two-thirds p character.
    • The last unhybridized 2p orbital containing one electron, which both carbon atoms have, will form a pi bond. Pi bonds are formed by planar orbital overlap, i.e. orbitals perpendicular to the positive nuclei, not in between them like the sigma bond. This out-of-plane pi interaction is further from the attractive nuclei and is therefore less strong than the sigma bond. This is why C=C bonds are not twice as strong as C-C single bonds.


      • Atomic orbitals
      Atomic orbitals


  • Hybridization can be used to show the bonding for a carbon triple bond, too. Carbon in ethyne (the simplest alkyne) is sp hybridized . How many orbitals do you think will be hybridized here?
    • Ethyne has only two atoms bonded to each carbon atom – it only needs to hybridize for a C-C and C-H sigma bond, taking the 2s orbital and one 2p orbital, hence it the label sp and 50:50 s and p character. The two perpendicular unhybridized p orbitals form the two pi bonds between the two carbon atoms to make the triple bond overall. See the diagram below.
    • Equally as we saw with ethene, ethyne has a carbon-carbon triple bond, which is not three times as strong as a C-C single bond because it is made of two pi bonds which are weaker than sigma bonds.

  • The type of hybridization a carbon atom has with its orbitals is easily found by the number of atoms the carbon center is bonded to:

    • Number of atoms bonded (to carbon)

    Hybridization

    Shape around carbon

    Example

    1

    Sp

    Linear

    C2H2

    2

    Sp2

    Trigonal

    C2H4

    3

    Sp3

    Tetrahedral

    CH4


  • Orbital hybridization is an extension of valence bond theory , which helps us explain the bonding and geometry in many compounds. It has a few basic principles:
    • Electrons exist in atomic orbitals or hybrid atomic orbitals which are localized to their principle atoms.
    • Electrons in atomic or hybrid orbitals repel each other so will maximise space between them around a central atom (similar to how VSEPR finds its bond angles).
    • Chemical bonds are formed by the overlap of separate atomic or hybrid orbitals, where sigma bonds are created by constructive head-on overlap of orbitals, and pi bonds are created by constructive planar overlap of orbitals.

    While this does accurately predict the properties of many compounds, it falls short in some areas.
    Another theory of chemical bonding is known as molecular orbital theory, which also accurately predicts the properties of many compounds, succeeding in places where valence bond theory does not. This is our focus for the next lesson.
Concept

Introduction to Atomic Orbital Hybridization

Atomic orbital hybridization is a fundamental concept in organic chemistry that explains the formation of chemical bonds in molecules. This process involves the mixing of atomic orbitals to create new hybrid orbitals with different shapes and energies. The introduction video on atomic orbital hybridization serves as an essential resource for students and enthusiasts alike, offering a visual and comprehensive explanation of this complex topic. By watching this video, viewers can gain a clearer understanding of how hybridization affects molecular geometry and bonding. Grasping the principles of hybridization is crucial for anyone studying organic chemistry, as it forms the basis for understanding molecular structures, reactivity, and properties. From simple molecules like methane to complex organic compounds, hybridization plays a pivotal role in determining their shapes and behaviors. By mastering this concept, students can better predict and explain various chemical phenomena, making it an indispensable tool in their organic chemistry toolkit.

Understanding the mixing of atomic orbitals is essential for predicting the shapes of molecules. The principles of molecular geometry and bonding are deeply intertwined with hybridization. This knowledge is not only applicable to simple molecules but also extends to complex organic compounds, making it a versatile and valuable concept in the study of chemistry.

Example

Mixing atomic orbitals - hybridization. Same AOs, different bonding?

Step 1: Introduction to Atomic Orbital Hybridization

In this lesson, we will explore a theory of bonding known as atomic orbital hybridization. This theory complements our existing knowledge and helps explain the properties observed in certain organic compounds. The main objective is to understand how atomic orbital hybridization explains bonding in simple molecules. The term "hybrid" is not new and is used for a reason. We will also learn to draw orbital diagrams using hybrid terms and see how this theory aligns with the observed properties of carbon-carbon bonds and the geometry of carbon compounds.

Step 2: Understanding the Concept of Hybridization

The term "hybrid" is familiar in various contexts, such as hybrid vehicles that use both electric and gasoline power. Similarly, in chemistry, hybridization refers to the mixing of atomic orbitals to form new hybrid orbitals. These hybrid orbitals have different shapes and energies compared to the original atomic orbitals. This concept helps explain the bonding and geometry of molecules, particularly those involving carbon atoms.

Step 3: Electron Configuration of Carbon

To understand hybridization, let's start with the electron configuration of carbon. Carbon has an electron configuration of 1s² 2s² 2p². The 1s electrons are in the core and not involved in bonding. The 2s and 2p electrons are in the outer shell and participate in bonding. In an orbital diagram, the 2s orbital contains two paired electrons, while the 2p orbitals contain two unpaired electrons. This configuration suggests that carbon can form two bonds using the unpaired 2p electrons.

Step 4: Example of Methylene (CH)

Methylene (CH) is an example of a compound called a carbene. Carbenes are highly reactive and unstable, often requiring special conditions to exist. In methylene, the two unpaired 2p electrons form two covalent bonds with hydrogen atoms. The paired 2s electrons do not participate in bonding, resulting in a highly reactive and unstable compound.

Step 5: Example of Methane (CH)

Methane (CH) is a more stable compound compared to methylene. In methane, carbon forms four covalent bonds with hydrogen atoms. These bonds are of equal energy, and the bond angles between them are all equal, resulting in a tetrahedral geometry. This observation is inconsistent with the electron configuration of carbon, which suggests only two unpaired electrons available for bonding.

Step 6: Explanation of Hybridization in Methane

The inconsistency in methane's bonding can be explained by the concept of hybridization. In methane, the 2s and 2p orbitals of carbon mix to form four equivalent hybrid orbitals. These hybrid orbitals, known as sp³ hybrid orbitals, have equal energy and form four equivalent bonds with hydrogen atoms. This mixing of s and p orbitals results in the observed tetrahedral geometry and equal bond energies in methane.

Step 7: Conclusion

Hybridization is a crucial concept in understanding the bonding and geometry of molecules, particularly those involving carbon atoms. By mixing atomic orbitals, hybridization explains how molecules like methane can have equal bond energies and specific geometries. This theory aligns with experimental observations and helps us make sense of the bonding in various organic compounds.

FAQs

  1. How to determine the hybridization of an atom?

    To determine the hybridization of an atom, follow these steps:

    • Count the number of sigma (σ) bonds and lone pairs on the atom.
    • Use the formula: Hybridization = Number of σ bonds + Number of lone pairs
    • If the result is 4, it's sp³; if 3, it's sp²; if 2, it's sp hybridization.

  2. What is hybridization of atoms?

    Hybridization is the process of mixing atomic orbitals to form new hybrid orbitals with different shapes and energies. This concept explains how atoms form chemical bonds in molecules, particularly in organic compounds. Hybridization helps predict molecular geometry and bond properties.

  3. What are the 4 types of hybridization?

    The four main types of hybridization are:

    • sp³: Tetrahedral geometry (e.g., methane)
    • sp²: Trigonal planar geometry (e.g., ethene)
    • sp: Linear geometry (e.g., acetylene)
    • sp³d: Trigonal bipyramidal geometry (e.g., PCl)

  4. How to tell if a carbon is sp² or sp³?

    To distinguish between sp² and sp³ hybridized carbon:

    • sp² carbon: Forms three σ bonds and one π bond; planar geometry with ~120° bond angles.
    • sp³ carbon: Forms four σ bonds; tetrahedral geometry with ~109.5° bond angles.
    • Look for double bonds or aromatic rings (sp²) vs. single bonds only (sp³).

  5. What is the importance of understanding hybridization in organic chemistry?

    Understanding hybridization is crucial in organic chemistry because it:

    • Explains molecular geometries and bond angles
    • Predicts reactivity and chemical properties of molecules
    • Helps in understanding complex organic structures and reactions
    • Provides insights into bond strengths and molecular stability
Prerequisites

Understanding atomic orbital hybridization is crucial in organic chemistry and molecular structure analysis. However, to fully grasp this concept, it's essential to have a solid foundation in two key prerequisite topics: atomic orbitals and energy levels and molecular geometry and VSEPR.

Atomic orbital hybridization builds upon the fundamental principles of atomic structure and electron configuration. A thorough understanding of atomic orbitals and energy levels is essential because hybridization involves the mixing of these orbitals to form new hybrid orbitals. This prerequisite topic provides the necessary background on how electrons are distributed in atoms and the shapes of different orbitals, which directly influences the hybridization process.

For instance, knowing the differences between s, p, d, and f orbitals and their energy levels helps explain why certain hybridizations occur in specific atoms. The concept of electron promotion, which is covered in the study of atomic orbitals, is also crucial for understanding how atoms can form hybrid orbitals that differ from their ground state configurations.

Similarly, molecular geometry and VSEPR (Valence Shell Electron Pair Repulsion) theory are closely related to orbital hybridization. This prerequisite topic provides the foundation for predicting molecular shapes, which are directly influenced by the hybridization of atomic orbitals. Understanding VSEPR theory helps explain why certain hybridizations lead to specific molecular geometries.

For example, the tetrahedral shape of a methane molecule is a result of sp³ hybridization in the carbon atom. Without a solid grasp of molecular geometry principles, it would be challenging to connect the dots between hybridization and the three-dimensional structures of molecules.

Moreover, the concepts learned in molecular geometry and bonding help in understanding how hybrid orbitals overlap to form chemical bonds. This knowledge is essential when studying more complex molecules and their properties, which are often determined by the type of hybridization and resulting molecular structure.

By mastering these prerequisite topics, students can more easily comprehend the intricacies of atomic orbital hybridization. The atomic orbitals and energy levels provide the foundational knowledge of electron behavior, while molecular geometry and VSEPR offer insights into how these hybridized orbitals influence molecular shapes and bonding patterns.

In conclusion, a solid understanding of these prerequisite topics is not just beneficial but essential for grasping the concept of atomic orbital hybridization. They provide the necessary context and background knowledge, allowing students to build a comprehensive understanding of molecular structure and bonding in organic chemistry.