Tetravalence of Carbon — Revision Notes

⚡ 30-Second Revision

Tetravalence: — Carbon forms 4 covalent bonds to achieve octet.\n- Valence Electrons: 4 ($2s^2 2p^2$).\n- Hybridization: Mixing of atomic orbitals to form new hybrid orbitals.\n - $sp^3$: 4 $\sigma$ bonds, 0 lone pairs. Geometry: Tetrahedral. Bond Angle: $109.5^\circ$. $s$-character: 25%. Example: $CH_4$.\n - $sp^2$: 3 $\sigma$ bonds, 0 lone pairs (1 double bond + 2 single bonds). Geometry: Trigonal Planar. Bond Angle: $120^\circ$. $s$-character: 33.3%. Example: $C_2H_4$.\n - $sp$: 2 $\sigma$ bonds, 0 lone pairs (1 triple bond + 1 single bond OR 2 double bonds). Geometry: Linear. Bond Angle: $180^\circ$. $s$-character: 50%. Example: $C_2H_2$, $CO_2$.\n- **Sigma ($\sigma$) Bond: Head-on overlap, strong, free rotation.\n- Pi ($\pi$) Bond:** Lateral overlap of unhybridized $p$-orbitals, weaker, restricted rotation.\n- Bond Counting: Single = 1 $\sigma$; Double = 1 $\sigma$ + 1 $\pi$; Triple = 1 $\sigma$ + 2 $\pi$.\n- Catenation: Carbon's ability to form long chains/rings with itself.

2-Minute Revision

Carbon's tetravalence, its ability to form four stable covalent bonds, is the cornerstone of organic chemistry. This property stems from its four valence electrons, which it shares to complete its octet.

To explain the observed molecular geometries, we use the concept of hybridization. Carbon can undergo $sp^3$, $sp^2$, or $sp$ hybridization.\n\n$sp^3$ hybridized carbons form four single bonds, resulting in a tetrahedral geometry with bond angles of $109.

5^\circ$(e.g., alkanes). These are all sigma ($\sigma$) bonds.\n\n$sp^2$hybridized carbons form one double bond and two single bonds, leading to a trigonal planar geometry with$120^\circ$ bond angles (e.

g., alkenes). A double bond consists of one $\sigma$ and one $\pi$ bond.\n\n$sp$ hybridized carbons form one triple bond and one single bond (or two double bonds), resulting in a linear geometry with $180^\circ$ bond angles (e.

g., alkynes, $CO_2$). A triple bond consists of one $\sigma$ and two $\pi$ bonds.\n\nUnderstanding these hybridization states, their associated geometries, bond angles, and the nature of sigma and pi bonds is crucial for predicting molecular structures, reactivity, and for accurately counting bonds in NEET questions.

5-Minute Revision

Carbon's tetravalence is its defining characteristic, allowing it to form four covalent bonds. This is due to its electronic configuration ($1s^2 2s^2 2p^2$), where it has four valence electrons. To achieve a stable octet, carbon shares these electrons.

The concept of hybridization explains how carbon forms these four bonds with specific geometries.\n\n1. **$sp^3$ Hybridization:** When carbon forms four single bonds (e.g., in methane, $CH_4$), one $s$ and three $p$ orbitals hybridize to form four equivalent $sp^3$ orbitals.

These point towards the corners of a tetrahedron, giving a bond angle of $109.5^\circ$. All bonds are sigma ($\sigma$) bonds, formed by head-on overlap. The $s$-character is 25%. Example: In ethane ($CH_3-CH_3$), both carbons are $sp^3$ hybridized, forming a C-C $\sigma$ bond and three C-H $\sigma$ bonds each.

\n\n2. **$sp^2$ Hybridization:** When carbon forms one double bond and two single bonds (e.g., in ethene, $C_2H_4$), one $s$ and two $p$ orbitals hybridize to form three $sp^2$ orbitals. These lie in a plane, forming a trigonal planar geometry with $120^\circ$ bond angles.

The remaining unhybridized $p$ orbital forms a pi ($\pi$) bond by lateral overlap. A double bond is one $\sigma$ and one $\pi$ bond. The $s$-character is 33.3%. Example: In propene ($CH_3-CH=CH_2$), the two carbons of the double bond are $sp^2$, while the methyl carbon is $sp^3$.

\n\n3. **$sp$ Hybridization:** When carbon forms one triple bond and one single bond (e.g., in ethyne, $C_2H_2$), or two double bonds (e.g., in carbon dioxide, $CO_2$), one $s$ and one $p$ orbital hybridize to form two $sp$ orbitals.

These are linear, with a $180^\circ$ bond angle. The two remaining unhybridized $p$ orbitals form two $\pi$ bonds. A triple bond is one $\sigma$ and two $\pi$ bonds. The $s$-character is 50%. Example: In $C_2H_2$, both carbons are $sp$ hybridized.

\n\nKey Takeaways for NEET:\n* Counting Bonds: Every single bond is a $\sigma$ bond. A double bond is one $\sigma$ and one $\pi$. A triple bond is one $\sigma$ and two $\pi$. This is a common MCQ type.

\n* Hybridization and Geometry: Memorize the direct correlation between hybridization ($sp^3, sp^2, sp$) and geometry (tetrahedral, trigonal planar, linear) and bond angles ($109.5^\circ, 120^\circ, 180^\circ$).

\n* **$s$-character:** Higher $s$-character means greater electronegativity, shorter bond length, and increased acidity of C-H bonds (e.g., terminal alkynes).

Prelims Revision Notes

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Carbon's Valence: — Carbon is tetravalent, meaning it forms four covalent bonds. This is due to its four valence electrons ($2s^2 2p^2$). It achieves a stable octet by sharing these electrons.\n2. Hybridization Types:\n * **$sp^3$ Hybridization:** Occurs when carbon forms four single bonds. One $s$ and three $p$ orbitals mix. \n * **Number of $\sigma$ bonds:** 4\n * **Number of $\pi$ bonds:** 0\n * Geometry: Tetrahedral\n * Bond Angle: $109.5^\circ$\n * **$s$-character:** 25%\n * Example: Methane ($CH_4$), Ethane ($CH_3-CH_3$)\n * **$sp^2$ Hybridization:** Occurs when carbon forms one double bond and two single bonds.\n * **Number of $\sigma$ bonds:** 3\n * **Number of $\pi$ bonds:** 1\n * Geometry: Trigonal Planar\n * Bond Angle: $120^\circ$\n * **$s$-character:** 33.3%\n * Example: Ethene ($C_2H_4$), Benzene ($C_6H_6$)\n * **$sp$ Hybridization:** Occurs when carbon forms one triple bond and one single bond, or two double bonds.\n * **Number of $\sigma$ bonds:** 2\n * **Number of $\pi$ bonds:** 2 (for triple bond) or 2 (for two double bonds)\n * Geometry: Linear\n * Bond Angle: $180^\circ$\n * **$s$-character:** 50%\n * Example: Ethyne ($C_2H_2$), Carbon Dioxide ($CO_2$)\n3. **Sigma ($\sigma$) and Pi ($\pi$) Bonds:**\n * **$\sigma$ bond:** Formed by head-on overlap of orbitals. Stronger, allows free rotation. Present in all single bonds, and as the first bond in double/triple bonds.\n * **$\pi$ bond:** Formed by lateral overlap of unhybridized $p$ orbitals. Weaker, restricts rotation. Present as the second bond in a double bond, and the second and third bonds in a triple bond.\n4. Counting Bonds:\n * Single bond: 1 $\sigma$\n * Double bond: 1 $\sigma$ + 1 $\pi$\n * Triple bond: 1 $\sigma$ + 2 $\pi$\n5. Catenation: Carbon's unique ability to form strong bonds with other carbon atoms, leading to long chains, branched structures, and rings. This is a direct consequence of its tetravalence and small size.\n6. **Importance of $s$-character:** Higher $s$-character implies greater electronegativity of the carbon atom, shorter and stronger bonds, and increased acidity of C-H bonds (e.g., terminal alkynes are acidic due to $sp$ hybridized carbon).

Vyyuha Quick Recall

To remember hybridization, geometry, and angles: \n\nSingle bonds $\rightarrow$ SP3 $\rightarrow$ Tetrahedral $\rightarrow$ 109.5\nDouble bond $\rightarrow$ SP2 $\rightarrow$ Trigonal Planar $\rightarrow$ 120\nTriple bond $\rightarrow$ SP $\rightarrow$ Linear $\rightarrow$ 180\n\nThink: Simple SP3 Takes 109.5; Double SP2 Tries 120; Triple SP Loves 180.