Atomic orbital hybridization

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.

Understanding Hybridization Theory

Origins and Development of Hybridization Theory

Hybridization theory, a cornerstone concept in organic chemistry, was developed in the early 20th century to explain the observed properties of carbon compounds. This theory was primarily introduced by Linus Pauling in 1931 to bridge the gap between classical valence bond theory and the complex behavior of molecules, especially those containing carbon. The need for such a theory arose from the inability of existing models to explain the tetrahedral structure of methane and other carbon-based molecules.

The Concept of Hybridization

Hybridization refers to the mixing of atomic orbitals to form new hybrid orbitals. These hybrid orbitals can better explain the observed molecular structures and bond properties. In the case of carbon, its ground state electronic configuration (1s² 2s² 2p²) doesn't directly account for its ability to form four equivalent bonds, as seen in methane (CH). Hybridization theory proposes that the 2s and 2p orbitals of carbon mix to form four equivalent sp³ hybrid orbitals, allowing carbon to form four equivalent bonds.

Types of Hybridization in Carbon Compounds

Carbon exhibits three main types of hybridization:

• sp³ Hybridization: Involves the mixing of one s and three p orbitals, resulting in four equivalent sp³ hybrid orbitals. This explains the tetrahedral geometry of molecules like methane.
• sp² Hybridization: Involves the mixing of one s and two p orbitals, forming three sp² hybrid orbitals and leaving one unhybridized p orbital. This explains the planar structure of molecules like ethene (CH).
• sp Hybridization: Involves the mixing of one s and one p orbital, creating two sp hybrid orbitals and leaving two unhybridized p orbitals. This explains the linear structure of molecules like acetylene (CH).

Explaining Bond Strengths through Hybridization

Hybridization theory helps explain variations in bond strengths observed in different carbon compounds. For instance:

• In sp³ hybridized carbon (like in alkanes), the C-C single bond is formed by the overlap of two sp³ orbitals.
• In sp² hybridized carbon (like in alkenes), the C=C double bond consists of one σ bond (from sp² orbital overlap) and one π bond (from unhybridized p orbital overlap). This double bond is stronger than a single bond but not twice as strong.
• In sp hybridized carbon (like in alkynes), the CC triple bond consists of one σ bond (from sp orbital overlap) and two π bonds (from unhybridized p orbitals), resulting in the strongest carbon-carbon bond.

Molecular Geometries Explained by Hybridization

Hybridization theory is crucial in explaining the molecular geometries of various carbon compounds:

• Tetrahedral Geometry: sp³ hybridization results in a tetrahedral arrangement, as seen in methane (CH) with bond angles of 109.5°.
• Trigonal Planar Geometry: sp² hybridization leads to a trigonal planar structure, exemplified by ethene (CH) with bond angles of 120°.
• Linear Geometry: sp hybridization produces a linear structure, as observed in acetylene (CH) with bond angles of 180°.

Applications and Importance in Chemistry

Understanding hybridization is crucial for predicting and explaining:

sp3 Hybridization: The Tetrahedral Carbon

sp3 hybridization is a fundamental concept in organic chemistry that explains the bonding behavior of carbon atoms in many organic compounds. This type of hybridization is particularly important in understanding the structure of methane (CH4), the simplest hydrocarbon molecule. In sp3 hybridization, one s orbital and three p orbitals of the carbon atom combine to form four equivalent hybrid orbitals.

Let's delve into the process of sp3 hybridization using methane as our primary example. In its ground state, a carbon atom has the electron configuration 1s² 2s² 2p². To form four equivalent bonds, as observed in methane, the carbon atom needs to promote one of its 2s electrons to the empty 2p orbital, resulting in four unpaired electrons. These four orbitals (one 2s and three 2p) then combine to form four equivalent sp3 hybrid orbitals.

The sp3 hybrid orbitals have a unique tetrahedral arrangement in three-dimensional space. This arrangement minimizes electron repulsion and maximizes stability. In methane, these four sp3 orbitals of carbon overlap with the 1s orbitals of four hydrogen atoms to form four equivalent sigma (σ) bonds. The tetrahedral geometry results in bond angles of 109.5° between any two bonds, creating the characteristic shape of methane.

The sigma bonds formed in sp3 hybridization are strong covalent bonds. They are characterized by head-on overlap between the sp3 orbital of carbon and the s orbital of hydrogen. This overlap occurs along the internuclear axis, resulting in a high electron density between the nuclei. The strength of these bonds contributes to the stability of methane and other sp3 hybridized molecules.

In terms of orbital shapes, sp3 hybrid orbitals have a distinctive appearance. They are asymmetrical, with a large lobe on one side and a smaller lobe on the other. This shape is a result of the combination of the spherical s orbital and the dumbbell-shaped p orbitals. The larger lobe is responsible for the majority of the bonding, while the smaller lobe has minimal contribution.

Energy levels play a crucial role in understanding sp3 hybridization. The energy of the sp3 hybrid orbitals lies between that of the original 2s and 2p orbitals. This intermediate energy level allows for more effective overlap with the orbitals of other atoms, leading to stronger bonds. The four sp3 hybrid orbitals are degenerate, meaning they have the same energy level, which explains the equivalence of the four C-H bonds in methane.

The concept of sp3 hybridization extends beyond methane to many other organic compounds. For instance, ethane (C2H6) features two sp3 hybridized carbon atoms, each bonded to three hydrogen atoms and one carbon atom. This results in a staggered conformation that minimizes steric hindrance. Similarly, larger alkanes and many other organic molecules utilize sp3 hybridization for their carbon-carbon and carbon-hydrogen bonds.

Understanding sp3 hybridization is crucial for predicting molecular geometries and reactivity in organic chemistry. The tetrahedral arrangement of bonds in sp3 hybridized carbons influences the three-dimensional structure of molecules, which in turn affects their physical and chemical properties. This concept is fundamental in explaining the behavior of saturated hydrocarbons, alcohols, and many other organic compounds.

In conclusion, sp3 hybridization, as exemplified by methane, demonstrates the remarkable ability of carbon to form stable, tetrahedral structures. The formation of four equivalent hybrid orbitals, their tetrahedral arrangement, and the resulting sigma bonds are key features of this hybridization. This understanding forms the basis for more complex organic structures and reactions, making sp3 hybridization an essential concept in the study of organic chemistry and molecular structure.

sp2 Hybridization: Double Bonds and Planar Structures

sp2 hybridization is a fundamental concept in organic reactions that plays a crucial role in understanding the structure and bonding of molecules with double bonds. This type of hybridization is particularly important in compounds like ethene (C2H4), which serves as an excellent example to explore the intricacies of sp2 hybridization and its consequences on molecular geometry and bonding.

In sp2 hybridization, one s orbital and two p orbitals of an atom combine to form three equivalent hybrid orbitals. These hybrid orbitals lie in a single plane and are arranged at 120° angles to each other. The remaining unhybridized p orbital is perpendicular to this plane. This arrangement is key to understanding the unique properties of molecules with double bonds.

Let's examine ethene as our primary example. In an ethene molecule, each carbon atom undergoes sp2 hybridization. The process begins with the promotion of one electron from the 2s orbital to an empty 2p orbital. Then, the 2s orbital mixes with two of the 2p orbitals to create three sp2 hybrid orbitals. The third 2p orbital remains unhybridized and is oriented perpendicular to the plane of the sp2 orbitals.

The formation of bonds in ethene involves both sigma (σ) and pi (π) bonds. The sp2 hybrid orbitals are responsible for forming sigma bonds. Two of these sp2 orbitals on each carbon atom form sigma bonds with hydrogen atoms. The third sp2 orbital from each carbon overlaps with the sp2 orbital of the other carbon, creating a strong sigma bond between the two carbon atoms. This sigma bonding framework forms the backbone of the ethene molecule.

The unique feature of sp2 hybridization is the formation of a pi bond, which gives ethene its characteristic double bond. This pi bond results from the side-by-side overlap of the unhybridized p orbitals of the two carbon atoms. The electron density in this pi bond is concentrated above and below the plane of the molecule, creating regions of high electron density that are crucial for many of ethene's chemical properties.

The combination of sigma and pi bonds in ethene results in a total of five bonds: four C-H sigma bonds and one C=C double bond (consisting of one sigma and one pi bond). This bonding arrangement is responsible for the high stability and reactivity of ethene and similar alkenes.

One of the most significant consequences of sp2 hybridization is the planar geometry it imparts to molecules. In ethene, all atoms lie in the same plane. This planar structure is a direct result of the arrangement of the sp2 hybrid orbitals. The 120° bond angles between these orbitals create a trigonal planar geometry around each carbon atom. This planar configuration is essential for the optimal overlap of p orbitals to form the pi bond.

The planar geometry of sp2 hybridized molecules has important implications for their properties and reactivity. It allows for maximum overlap of the p orbitals, strengthening the pi bond. This planarity also influences the molecule's interaction with light, contributing to the unique spectroscopic properties of alkenes and other sp2 hybridized compounds.

sp2 hybridization is not limited to carbon atoms in alkenes. It's also observed in other organic compounds like aldehydes, ketones, and carboxylic acids, as well as in some inorganic molecules. In each case, the principles of orbital hybridization, sigma and pi bond formation, and planar geometry apply, though the specific atoms involved may differ.

Understanding sp2 hybridization is crucial for predicting and explaining the behavior of molecules in various chemical organic reactions. For instance, the reactivity of the pi bond in alkenes is the basis for many important organic reactions, including addition reactions and polymerization processes. The planar structure also plays a role in determining the stereochemistry of reaction products.

In conclusion, sp2 hybridization, as exemplified by ethene, is a cornerstone concept in organic chemistry. It explains the formation of double bonds, the planar geometry of certain molecules, and provides a foundation for understanding more complex organic structures and reactions. By mastering the

sp Hybridization: Triple Bonds and Linear Structures

sp hybridization is a fundamental concept in organic chemistry that plays a crucial role in understanding the structure and bonding of molecules with triple bonds, such as ethyne (also known as acetylene). This type of hybridization results in a unique linear geometry and is essential for explaining the formation of one sigma bond and two pi bonds in triple-bonded molecules.

To understand sp hybridization, let's focus on the example of ethyne (C2H2). In an ethyne molecule, each carbon atom undergoes sp hybridization. This process involves the mixing of one s orbital and one p orbital from the carbon atom's valence shell. The result is two equivalent sp hybrid orbitals, while two unhybridized p orbitals remain.

The formation of sp hybrid orbitals can be visualized as follows:

1. The carbon atom's 2s orbital combines with one of its 2p orbitals (typically the 2px orbital).
2. This combination results in two new, equivalent sp hybrid orbitals.
3. The remaining two p orbitals (2py and 2pz) remain unhybridized.

The sp hybrid orbitals have a unique shape and orientation. They are directed 180 degrees apart from each other, forming a linear arrangement. This linear orientation is key to understanding the overall geometry of molecules like ethyne.

In ethyne, the sp hybridization of both carbon atoms leads to the following bonding arrangement:

1. One sp hybrid orbital from each carbon atom overlaps head-on to form a strong sigma (σ) bond between the two carbon atoms.
2. The remaining sp hybrid orbital on each carbon atom forms a sigma bond with a hydrogen atom.
3. The two unhybridized p orbitals on each carbon atom (py and pz) are perpendicular to the molecular axis and to each other.

The unhybridized p orbitals play a crucial role in forming the triple bond characteristic of ethyne. Here's how the triple bond is created:

1. The sigma bond between the carbon atoms forms the first part of the triple bond.
2. The unhybridized py orbitals from both carbon atoms overlap side-by-side to form one pi (π) bond.
3. Similarly, the unhybridized pz orbitals overlap to form a second pi bond.

This combination of one sigma bond and two pi bonds constitutes the triple bond in ethyne. The resulting molecular structure is linear, with bond angles of 180 degrees. This linear geometry is a direct consequence of the sp hybridization and is characteristic of molecules with triple bonds.

The sp hybridization in ethyne contributes to several important properties of the molecule:

• High bond strength: The triple bond is very strong, making ethyne a relatively stable molecule.
• Reactivity: Despite its stability, the pi bonds can participate in addition reactions, making ethyne an important starting material in organic synthesis.
• Acidity: The sp-hybridized carbon atoms make the hydrogen atoms in ethyne slightly acidic, allowing for unique chemical reactions.

Understanding sp hybridization is crucial for chemists and students alike, as it provides insights into the structure and reactivity of many important organic compounds. Beyond ethyne, sp hybridization is also observed in other molecules with triple bonds, such as nitriles (R-CN) and some metal complexes.

In conclusion, sp hybridization in ethyne (acetylene) exemplifies the formation of triple bonds and linear molecular geometries. By mixing one s and one p orbital, carbon atoms create two sp hybrid orbitals while leaving two p orbitals unhybridized. This arrangement leads to the formation of one sigma bond and two pi bonds, resulting in the characteristic triple bond and linear structure of ethyne. This concept is fundamental to understanding the bonding and geometry of molecules with triple bonds and plays a significant role

Comparing Hybridization Types and Their Effects

Hybridization is a fundamental concept in organic chemistry that explains the bonding behavior of atoms, particularly carbon. The three main types of hybridization - sp3, sp2, and sp - play crucial roles in determining molecular structure and properties. Understanding these hybridization types and their effects on bond angles, molecular shapes, and bond strengths is essential for predicting and explaining molecular behavior.

Sp3 hybridization occurs when one s orbital and three p orbitals combine to form four equivalent hybrid orbitals. This type of hybridization is characteristic of single-bonded carbon atoms, such as those found in methane (CH4). In sp3 hybridization, the bond angles are approximately 109.5 degrees, resulting in a tetrahedral molecular shape. The bonds formed by sp3 hybrid orbitals are typically single bonds, which are relatively weaker compared to double or triple bonds.

Sp2 hybridization involves the mixing of one s orbital and two p orbitals, creating three hybrid orbitals. This hybridization is common in molecules with double bonds, like ethene (C2H4). The bond angles between sp2 hybrid orbitals are about 120 degrees, leading to a trigonal planar geometry. The remaining unhybridized p orbital is perpendicular to this plane and participates in pi bonding. Molecules with sp2 hybridization often exhibit stronger bonds due to the presence of both sigma and pi bonds.

Sp hybridization is the combination of one s orbital and one p orbital, resulting in two hybrid orbitals. This type of hybridization is found in molecules with triple bonds, such as acetylene (C2H2). The bond angle between sp hybrid orbitals is 180 degrees, creating a linear molecular shape. Sp hybridized atoms form the strongest bonds among the three types due to the increased s-character and the presence of two pi bonds in addition to the sigma bond.

Comparing the bond strengths, sp hybridized bonds are generally the strongest, followed by sp2, and then sp3. This trend is due to the increasing s-character in the hybrid orbitals, which allows for greater orbital overlap and stronger bonding. The bond lengths follow an inverse relationship, with sp hybridized bonds being the shortest and sp3 the longest.

Examples of molecules exhibiting sp3 hybridization include methane (CH4), ethane (C2H6), and other alkanes. These molecules have tetrahedral geometry and form only single bonds. Sp2 hybridization can be observed in ethene (C2H4), benzene (C6H6), and carbonyl compounds like formaldehyde (CH2O). These molecules have planar structures and contain at least one double bond. Acetylene (C2H2), carbon dioxide (CO2), and hydrogen cyanide (HCN) are examples of molecules with sp hybridization, characterized by their linear shape and triple bonds.

Hybridization theory is a powerful tool for predicting and explaining molecular properties. It helps in understanding the three-dimensional structure of molecules, which is crucial for determining their reactivity and behavior in chemical reactions. For instance, the tetrahedral arrangement of sp3 hybridized carbon explains the stability of alkanes and their resistance to many chemical reactions. The planar structure of sp2 hybridized molecules like benzene elucidates their unique aromatic properties and reactivity patterns.

Moreover, hybridization theory aids in explaining the differences in bond strengths and lengths among various types of chemical bonds. This knowledge is invaluable in predicting the stability of molecules and their tendency to undergo certain reactions. For example, the stronger bonds in sp hybridized molecules make them more reactive in addition reactions compared to sp2 or sp3 hybridized molecules.

The concept of hybridization also helps in understanding the electronic distribution within molecules, which is essential for predicting their polarity and intermolecular interactions. This, in turn, affects properties such as boiling points, solubility, and reactivity. For instance, the tetrahedral arrangement of sp3 hybridized molecules like methane results in a symmetrical distribution of charge, making the molecule non-polar. In contrast, the trigonal planar structure of sp2 hybridized molecules can lead to polar bonds and overall molecular pol

Conclusion: The Importance of Hybridization in Chemistry

Atomic orbital hybridization is a fundamental concept in understanding chemical bonding and molecular structure. The introduction video provides a crucial foundation for grasping this complex topic. Hybridization explains how atoms can form stable bonds by mixing atomic orbitals, resulting in new hybrid orbitals with different shapes and energies. This process is essential in organic molecules, where carbon atoms can form various types of bonds. The importance of hybridization extends beyond simple molecules, playing a vital role in the structure and properties of complex organic compounds. By understanding hybridization, chemists can predict molecular geometries, bond angles, and reactivity. As you continue your chemistry journey, exploring hybridization in more intricate complex organic compounds will deepen your understanding of chemical bonding and molecular behavior. This knowledge is invaluable for fields such as biochemistry, materials science, and drug design. Remember, mastering hybridization opens doors to comprehending the fascinating world of molecular architecture and chemical reactions.

Mixing Atomic Orbitals - Hybridization

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

Prerequisite Topics

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.

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:



• 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.






• 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.