## 3.7 Hybridization **Hybridization** is the concept of mixing atomic orbitals of different energies and shapes to produce a new set of equivalent orbitals, called **hybrid orbitals**. These new orbitals have the same energy and shape. During hybridization, electrons in the ground state are often promoted to an excited state, which increases the number of unpaired electrons available for bonding. The energy required for this electron promotion is compensated by the energy released during bond formation. The type of hybridization depends on the atomic orbitals involved (s, p, d, etc.). This concept is crucial for explaining molecular geometry and bonding properties. ### Types of Hybridization There are several types of hybridization, including $sp$, $sp^2$, $sp^3$, $dsp^2$, $dsp^3$, and $d^2sp^3$. We will focus on the first three types. ### 3.7.1 sp Hybridization * **Definition:** The intermixing of one *s* and one *p* orbital to produce two new, equivalent *sp* hybrid orbitals. * **Character:** Each sp-hybridized orbital has 50% *s*-character and 50% *p*-character. * **Geometry:** The two *sp* hybrid orbitals are arranged linearly, with a bond angle of **$180^\circ$**. * **Bonding:** *sp* hybrid orbitals typically form sigma ($\sigma$) bonds. #### Examples of sp Hybridization **1. Beryllium Chloride ($BeCl_2$)** * **Central Atom:** Beryllium (Be) * **Ground State Configuration of Be:** $1s^2 2s^2$ * **Excited State:** To form two bonds, one $2s$ electron is promoted to an empty $2p$ orbital. * **Excited State Configuration of Be:** $1s^2 2s^1 2p^1$ * **Hybridization:** The $2s$ and one $2p$ orbital mix to form two *sp* hybrid orbitals. * **Bonding:** Each *sp* hybrid orbital of Beryllium overlaps with a *p* orbital of a Chlorine atom to form two Be-Cl sigma bonds. * **Resulting Shape:** The molecule is **linear** with a Cl-Be-Cl bond angle of **$180^\circ$**. **2. Acetylene ($C_2H_2$)** * **Central Atom:** Carbon (C) * **Ground State Configuration of C:** $1s^2 2s^2 2p_x^1 2p_y^1$ * **Excited State:** To form four bonds, a $2s$ electron is promoted to the empty $2p_z$ orbital. * **Excited State Configuration of C:** $1s^2 2s^1 2p_x^1 2p_y^1 2p_z^1$ * **Hybridization:** Each carbon atom undergoes *sp* hybridization, mixing its $2s$ and one $2p$ orbital to form two linear *sp* hybrid orbitals. Two *p* orbitals ($2p_y$ and $2p_z$) remain unhybridized on each carbon. * **Bonding:** 1. One *sp* orbital from each carbon overlaps to form a C-C sigma ($\sigma$) bond. 2. The other *sp* orbital on each carbon overlaps with the *1s* orbital of a hydrogen atom to form two C-H sigma ($\sigma$) bonds. 3. The two unhybridized *p* orbitals on each carbon overlap sideways to form two pi ($\pi$) bonds between the carbon atoms. * **Resulting Shape:** The molecule is **linear** with H-C-C and C-C-H bond angles of **$180^\circ$**. The C-C bond is a triple bond (one $\sigma$ and two $\pi$). ### 3.7.2 $sp^2$ Hybridization * **Definition:** The intermixing of one *s* and two *p* orbitals to produce three new, equivalent *$sp^2$* hybrid orbitals. * **Character:** Each $sp^2$-hybridized orbital has 33.3% *s*-character and 66.7% *p*-character. * **Geometry:** The three *$sp^2$* hybrid orbitals are arranged in a **trigonal planar** geometry, with bond angles of **$120^\circ$**. #### Examples of $sp^2$ Hybridization **1. Boron Trichloride ($BCl_3$)** * **Central Atom:** Boron (B) * **Ground State Configuration of B:** $1s^2 2s^2 2p^1$ * **Excited State:** A $2s$ electron is promoted to a $2p$ orbital to make three unpaired electrons available for bonding. * **Excited State Configuration of B:** $1s^2 2s^1 2p_x^1 2p_y^1$ * **Hybridization:** The $2s$ orbital and two $2p$ orbitals mix to form three *$sp^2$* hybrid orbitals. * **Bonding:** The three *$sp^2$* hybrid orbitals of Boron each overlap with a *p* orbital from a Chlorine atom, forming three B-Cl sigma ($\sigma$) bonds. * **Resulting Shape:** The molecule is **trigonal planar** with Cl-B-Cl bond angles of **$120^\circ$**. **2. Ethylene ($C_2H_4$)** * **Central Atom:** Carbon (C) * **Excited State Configuration of C:** $1s^2 2s^1 2p_x^1 2p_y^1 2p_z^1$ * **Hybridization:** Each carbon atom undergoes *$sp^2$* hybridization, mixing its $2s$ and two $2p$ orbitals. This leaves one unhybridized $2p$ orbital on each carbon. * **Bonding:** 1. One *$sp^2$* orbital from each carbon overlaps to form a C-C sigma ($\sigma$) bond. 2. The remaining two *$sp^2$* orbitals on each carbon overlap with *1s* orbitals of hydrogen atoms, forming four C-H sigma ($\sigma$) bonds. 3. The unhybridized *p* orbitals on each carbon overlap sideways to form one pi ($\pi$) bond. * **Resulting Shape:** The molecule is **trigonal planar** around each carbon, with H-C-H and H-C-C bond angles of approximately **$120^\circ$**. All six atoms lie in the same plane. ### 3.7.3 $sp^3$ Hybridization * **Definition:** The intermixing of one *s* and three *p* orbitals to produce four new, equivalent *$sp^3$* hybrid orbitals. * **Character:** Each $sp^3$-hybridized orbital has 25% *s*-character and 75% *p*-character. * **Geometry:** The four *$sp^3$* hybrid orbitals are arranged in a **tetrahedral** geometry, with bond angles of **$109.5^\circ$**. #### Examples of $sp^3$ Hybridization **1. Methane ($CH_4$)** * **Central Atom:** Carbon (C) * **Excited State Configuration of C:** $1s^2 2s^1 2p_x^1 2p_y^1 2p_z^1$ * **Hybridization:** The carbon atom undergoes *$sp^3$* hybridization, mixing its $2s$ and all three $2p$ orbitals. * **Bonding:** Each of the four *$sp^3$* hybrid orbitals on carbon overlaps with the *1s* orbital of a hydrogen atom, forming four C-H sigma ($\sigma$) bonds. * **Resulting Shape:** The molecule is **tetrahedral** with H-C-H bond angles of **$109.5^\circ$**. **2. Ammonia ($NH_3$)** * **Central Atom:** Nitrogen (N) * **Ground State Configuration of N:** $1s^2 2s^2 2p_x^1 2p_y^1 2p_z^1$ * **Hybridization:** Nitrogen undergoes *$sp^3$* hybridization. The $2s$ and three $2p$ orbitals mix to form four *$sp^3$* hybrid orbitals. * **Bonding & Structure:** * Three *$sp^3$* orbitals are half-filled and overlap with *1s* orbitals of three hydrogen atoms to form three N-H sigma ($\sigma$) bonds. * One *$sp^3$* orbital is completely filled and holds a **lone pair** of electrons. * **Resulting Shape:** The electron geometry is tetrahedral, but the molecular shape is **trigonal pyramidal**. The lone pair exerts greater repulsion than bonding pairs, compressing the H-N-H bond angle to **$107^\circ$**. **3. Water ($H_2O$)** * **Central Atom:** Oxygen (O) * **Ground State Configuration of O:** $1s^2 2s^2 2p_x^2 2p_y^1 2p_z^1$ * **Hybridization:** Oxygen undergoes *$sp^3$* hybridization, mixing its $2s$ and three $2p$ orbitals to form four *$sp^3$* hybrid orbitals. * **Bonding & Structure:** * Two *$sp^3$* orbitals are half-filled and overlap with *1s* orbitals of two hydrogen atoms to form two O-H sigma ($\sigma$) bonds. * Two *$sp^3$* orbitals are completely filled and hold **two lone pairs** of electrons. * **Resulting Shape:** The electron geometry is tetrahedral, but the molecular shape is **bent** or **V-shaped**. The two lone pairs exert strong repulsion, compressing the H-O-H bond angle to **$104.5^\circ$**. ### Possible Questions/Answers * **Q:** What is the molecular geometry of a molecule with an $sp^2$ hybridized central atom and no lone pairs? **A:** Trigonal planar, with bond angles of $120^\circ$. * **Q:** Why is the bond angle in water ($104.5^\circ$) smaller than the ideal tetrahedral angle ($109.5^\circ$)? **A:** The central oxygen atom has two lone pairs of electrons. The repulsion between these lone pairs is stronger than the repulsion between bonding pairs, which pushes the O-H bonds closer together and reduces the bond angle. * **Q:** How many sigma ($\sigma$) and pi ($\pi$) bonds are in an acetylene ($C_2H_2$) molecule? **A:** There are 3 sigma bonds (one C-C, two C-H) and 2 pi bonds (between the two carbons). ### Summary * **Hybridization** is the mixing of atomic orbitals to form new, identical hybrid orbitals for bonding. * The type of hybridization determines the molecular geometry and bond angles. * Lone pairs of electrons occupy hybrid orbitals and influence the final molecular shape by exerting stronger repulsive forces than bonding pairs. | Hybridization Type | Orbitals Mixed | Hybrid Orbitals Formed | Geometry | Bond Angle | Example | | :--- | :---: | :---: | :--- | :---: | :--- | | **$sp$** | one s, one p | 2 | Linear | $180^\circ$ | $BeCl_2, C_2H_2$ | | **$sp^2$** | one s, two p | 3 | Trigonal Planar | $120^\circ$ | $BCl_3, C_2H_4$ | | **$sp^3$** | one s, three p | 4 | Tetrahedral | $109.5^\circ$ | $CH_4, NH_3, H_2O$ | * **Significance:** Understanding hybridization is fundamental to predicting the three-dimensional shapes of molecules, which in turn dictates their physical and chemical properties, such as polarity, reactivity, and biological function.