Covalent Bond — Revision Notes

⚡ 30-Second Revision

Covalent Bond: — Sharing of electrons between atoms.
Octet Rule: — Atoms aim for 8 valence electrons (2 for H).
Types: — Single (1 shared pair), Double (2 shared pairs), Triple (3 shared pairs).
Bond Order: — Number of shared electron pairs.
Bond Length: — Distance between nuclei. Bond Order $\uparrow \implies$ Bond Length $\downarrow$.
Bond Energy: — Energy to break bond. Bond Order $\uparrow \implies$ Bond Energy $\uparrow$.
Sigma ($\sigma$) Bond: — Head-on overlap, stronger, allows rotation.
Pi ($\pi$) Bond: — Sideways overlap, weaker, restricts rotation.
Coordinate Bond: — One atom donates both electrons (e.g., $\text{NH}_4^+$).
Polarity: — Determined by electronegativity difference ($\Delta\text{EN}$). $\Delta\text{EN} \approx 0 \implies$ Nonpolar; $\Delta\text{EN} > 0 \implies$ Polar.
Molecular Polarity: — Vector sum of bond dipoles; depends on geometry (e.g., $\text{CO}_2$ is nonpolar, $\text{H}_2\text{O}$ is polar).
VSEPR Theory: — Electron pair repulsion determines molecular geometry and bond angles.
Hybridization: — Mixing of atomic orbitals to form new hybrid orbitals (sp, sp$^2$, sp$^3$, sp$^3$d, sp$^3$d$^2$). Steric Number = $\sigma$ bonds + lone pairs.

2-Minute Revision

Covalent bonds are formed by sharing electrons, primarily between non-metals, to achieve stable electron configurations, usually an octet. This sharing can be one (single), two (double), or three (triple) pairs, defining the bond order.

Higher bond order means shorter bond length and stronger bond energy. Single bonds are always sigma ($\sigma$) bonds, formed by head-on orbital overlap. Double bonds have one $\sigma$ and one pi ($\pi$) bond (sideways overlap), and triple bonds have one $\sigma$ and two $\pi$ bonds.

$\sigma$ bonds are stronger than $\pi$ bonds and allow free rotation, while $\pi$ bonds restrict it.

Bond polarity depends on the electronegativity difference ($\Delta\text{EN}$). If $\Delta\text{EN}$ is small, the bond is nonpolar; if significant, it's polar. Molecular polarity, however, also considers molecular geometry.

Symmetrical molecules with polar bonds can be nonpolar overall (e.g., $\text{CO}_2$). VSEPR theory is crucial for predicting molecular shapes and bond angles by minimizing electron pair repulsion. Hybridization (sp, sp$^2$, sp$^3$, etc.

) explains the formation of equivalent orbitals for bonding and is determined by the steric number (sigma bonds + lone pairs). Remember exceptions to the octet rule (e.g., $\text{BF}_3$, $\text{SF}_6$) and coordinate covalent bonds where one atom donates both electrons.

5-Minute Revision

A covalent bond is the sharing of electron pairs between atoms, typically non-metals, to achieve a stable electron configuration (octet or duplet). This sharing is quantified by bond order: 1 for single, 2 for double, 3 for triple bonds. A higher bond order implies greater electron density between nuclei, leading to a shorter bond length and a higher bond energy (stronger bond).

Covalent bonds are categorized into **sigma ($\sigma$) and pi ($\pi$) bonds**. Sigma bonds result from direct, head-on overlap of atomic orbitals (s-s, s-p, p-p) and are the strongest, allowing free rotation. Pi bonds result from sideways overlap of unhybridized p-orbitals, are weaker, and restrict rotation. A single bond is 1 $\sigma$. A double bond is 1 $\sigma$ + 1 $\pi$. A triple bond is 1 $\sigma$ + 2 $\pi$.

Coordinate covalent bonds are special cases where one atom donates both electrons to the shared pair (e.g., $\text{NH}_3$ donating to $\text{H}^+$ to form $\text{NH}_4^+$). Once formed, they are indistinguishable from regular covalent bonds.

Bond polarity is determined by the electronegativity difference ($\Delta\text{EN}$) between bonded atoms. A large $\Delta\text{EN}$ creates a polar bond with partial charges ($\delta^+$ and $\delta^-$).

If $\Delta\text{EN}$ is zero or very small, the bond is nonpolar. Molecular polarity depends on both bond polarity and molecular geometry. Symmetrical molecules (e.g., $\text{CO}_2$, $\text{CCl}_4$) can be nonpolar even if they contain polar bonds, as their bond dipoles cancel out due to symmetry.

Asymmetrical molecules (e.g., $\text{H}_2\text{O}$, $\text{NH}_3$) are polar.

VSEPR Theory (Valence Shell Electron Pair Repulsion) predicts molecular geometry. Electron pairs (bonding and lone pairs) around a central atom repel each other and arrange to minimize repulsion. For example, 4 electron domains lead to tetrahedral electron geometry, but if one is a lone pair (e.g., $\text{NH}_3$), the molecular geometry is trigonal pyramidal.

Hybridization explains the formation of equivalent bonds and observed geometries. Atomic orbitals (s, p, d) mix to form hybrid orbitals (sp, sp$^2$, sp$^3$, sp$^3$d, sp$^3$d$^2$). The steric number (number of sigma bonds + number of lone pairs) determines hybridization. For example, a steric number of 4 implies sp$^3$ hybridization. Remember exceptions to the octet rule, such as electron-deficient molecules (e.g., $\text{BF}_3$) and expanded octets (e.g., $\text{SF}_6$).

Worked Example: Determine the hybridization and geometry of $\text{H}_2\text{O}$.

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Lewis Structure: — Oxygen is central, bonded to two H atoms, and has two lone pairs.
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Sigma Bonds: — Two O-H single bonds = 2 $\sigma$ bonds.
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Lone Pairs: — Two lone pairs on Oxygen.
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Steric Number: — 2 ($\sigma$ bonds) + 2 (lone pairs) = 4.
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Hybridization: — Steric number 4 $\implies$ $\text{sp}^3$ hybridization.
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Electron Geometry: — 4 electron domains $\implies$ tetrahedral.
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Molecular Geometry: — 2 bonding pairs + 2 lone pairs $\implies$ bent (V-shaped).

Prelims Revision Notes

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Definition: — Covalent bond = sharing of electrons. Formed between non-metals.
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Octet Rule: — Atoms achieve 8 valence electrons (H achieves 2). Exceptions: Electron-deficient (e.g., $\text{BF}_3$), Expanded octet (e.g., $\text{PCl}_5$, $\text{SF}_6$), Odd-electron (e.g., $\text{NO}_2$).
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Types of Bonds:

* Single Bond: 1 shared pair, 1 $\sigma$ bond. * Double Bond: 2 shared pairs, 1 $\sigma$ + 1 $\pi$ bond. * Triple Bond: 3 shared pairs, 1 $\sigma$ + 2 $\pi$ bonds.

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**Sigma ($\sigma$) vs. Pi ($\pi$) Bonds:**

* **$\sigma$:** Head-on overlap (s-s, s-p, p-p), stronger, allows free rotation. * **$\pi$:** Sideways overlap (p-p), weaker, restricts rotation.

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Coordinate Covalent Bond (Dative Bond): — Both shared electrons from one atom (donor, with lone pair) to another (acceptor, with empty orbital). Example: $\text{NH}_4^+$, $\text{H}_3\text{O}^+$.
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Bond Parameters:

* Bond Order: Number of bonds between atoms. Higher bond order $\implies$ shorter bond length, higher bond energy. * Bond Length: Distance between nuclei. $\text{Single} > \text{Double} > \text{Triple}$. * Bond Energy: Energy to break bond. $\text{Triple} > \text{Double} > \text{Single}$. * Bond Angle: Determined by VSEPR and hybridization.

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Electronegativity (EN): — Atom's ability to attract shared electrons.
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Bond Polarity:

* Nonpolar: $\Delta\text{EN} \approx 0$ (equal sharing, e.g., $\text{H}_2$, $\text{Cl}_2$). * Polar: $\Delta\text{EN} > 0$ (unequal sharing, partial charges, e.g., $\text{HCl}$, $\text{H}_2\text{O}$). Creates a bond dipole.

1
Molecular Polarity:

* Net dipole moment = vector sum of bond dipoles. * Symmetrical molecules with polar bonds can be nonpolar (e.g., $\text{CO}_2$, $\text{CCl}_4$). * Asymmetrical molecules are polar (e.g., $\text{H}_2\text{O}$, $\text{NH}_3$).

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VSEPR Theory (Valence Shell Electron Pair Repulsion):

* Electron pairs (bonding + lone pairs) repel and arrange to minimize repulsion. * **Steric Number (SN) = ($\sigma$ bonds) + (Lone pairs)**. * SN 2: Linear (e.g., $\text{BeCl}_2$, $\text{CO}_2$) * SN 3: Trigonal Planar (e.

g., $\text{BF}_3$, $\text{SO}_3$) * SN 4: Tetrahedral (e.g., $\text{CH}_4$, $\text{NH}_4^+$) * SN 5: Trigonal Bipyramidal (e.g., $\text{PCl}_5$) * SN 6: Octahedral (e.g., $\text{SF}_6$) * Lone pairs cause greater repulsion, distorting ideal bond angles (LP-LP > LP-BP > BP-BP).

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Hybridization: — Mixing of atomic orbitals to form new hybrid orbitals.

* SN 2 $\implies$ sp (Linear) * SN 3 $\implies$ sp$^2$ (Trigonal Planar) * SN 4 $\implies$ sp$^3$ (Tetrahedral) * SN 5 $\implies$ sp$^3$d (Trigonal Bipyramidal) * SN 6 $\implies$ sp$^3$d$^2$ (Octahedral)

Vyyuha Quick Recall

To remember the order of bond parameters (length, energy) with increasing bond order:

Bond Order Increases, Bond Length Decreases, Bond Energy Increases.

B.O.I. B.L.D. B.E.I. (Pronounced: 'Boy, B.L.D. B.E.I.!')

This helps recall that as you go from single to double to triple bonds, the bond gets shorter and stronger.