Covalent bonding and coordinate (dative covalent) bonding - Chemistry (9701) - Cambridge A-Level

* The content provided by thinka is generated by AI and may not always be accurate or up-to-date. Please use it as a supplementary resource and verify with official materials.

Covalent Bonding and Coordinate (Dative) Bonding

Welcome to the world of bonding! This chapter is absolutely fundamental because it explains how atoms stick together to form the molecules that make up everything around us—from the air we breathe to the DNA in our cells.
We're going beyond simple sharing to look at the geometry and strength of these bonds. Don't worry if concepts like orbital overlap sound complicated; we will break them down step-by-step!

1. Covalent Bonding: The Shared Partnership (Syllabus 3.4.1)

Covalent bonding occurs primarily between non-metal atoms. Unlike ionic bonding, where electrons are completely transferred, covalent bonding involves the sharing of electrons.

Definition of Covalent Bonding

The official definition you must know is:
Covalent bonding is the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.

How Covalent Bonds Form

When two atoms approach, their electron clouds overlap. The electrons in this shared region are attracted simultaneously to the positive nuclei of *both* atoms, effectively holding the atoms together.

Quick Tip: Covalent bonding allows atoms to achieve a stable electronic configuration, often filling their outer shell to form a stable octet (8 electrons).

Common Covalent Molecules and Bonding Patterns

You need to be able to describe the bonding (using dot-and-cross diagrams, as covered in Syllabus 3.7, or descriptions) in many simple molecules:

Single Bond (One shared pair):
Examples: Hydrogen ($\text{H}_2$), Chlorine ($\text{Cl}_2$), Hydrogen Chloride ($\text{HCl}$), Methane ($\text{CH}_4$), Ethane ($\text{C}_2\text{H}_6$).
Double Bond (Two shared pairs):
Examples: Oxygen ($\text{O}_2$), Carbon Dioxide ($\text{CO}_2$), Ethene ($\text{C}_2\text{H}_4$).
Triple Bond (Three shared pairs):
Examples: Nitrogen ($\text{N}_2$).

Did you know? The strength of the bond increases from single to double to triple bond because more electron pairs are shared, leading to a stronger attractive force between the nuclei and the electrons.

2. Expanding the Octet (Syllabus 3.4.1(b))

Most atoms are happiest when they have 8 electrons in their outer shell (the octet rule). However, atoms in Period 3 and below can sometimes exceed this limit, holding 10, 12, or even more electrons.

Why Can Period 3 Elements Expand Their Octet?

Elements like Sulfur (S) and Phosphorus (P) are located in Period 3 of the Periodic Table. This means that in addition to the 3s and 3p orbitals, they also have empty 3d orbitals available.

These empty d orbitals can be used to accommodate extra electrons when forming bonds, allowing the atom to bond with more than four surrounding atoms.

Examples of Expanded Octets (More than 8 valence electrons):

Sulfur Dioxide ($\text{SO}_2$): Sulfur uses more than 8 electrons in its bonding structure.
Phosphorus Pentachloride ($\text{PCl}_5$): Phosphorus forms five bonds, resulting in 10 electrons around the central P atom.
Sulfur Hexafluoride ($\text{SF}_6$): Sulfur forms six bonds, resulting in 12 electrons around the central S atom.

Key Takeaway: Octet expansion is only possible for elements in Period 3 or below because they have energetically accessible d orbitals available for bonding.

3. Coordinate (Dative Covalent) Bonding (Syllabus 3.4.1(c))

A dative covalent bond is a special type of covalent bond.

Definition of Dative Covalent Bonding

It is a bond where both shared electrons originate from the same atom. The atom providing the pair of electrons is the donor; the atom accepting the pair is the acceptor.

Analogy: Imagine ordinary covalent bonding is like two friends each contributing $1 to buy a $2 snack. Dative bonding is like one generous friend paying the full $2 for the snack, but they still share it equally.

Key Examples of Coordinate Bonding:

1. The Ammonium Ion ($\text{NH}_4^{+}$):

Ammonia ($\text{NH}_3$) has a lone pair of electrons on the nitrogen atom. A hydrogen ion ($\text{H}^+$) is simply a proton with an empty valence shell.

The nitrogen atom donates its lone pair into the empty 1s orbital of the $\text{H}^+$ ion, forming a dative bond.
$\text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^{+}$
(In the context of the syllabus, this is exemplified by the reaction between ammonia and hydrogen chloride gases.)

2. Aluminium Chloride Dimer ($\text{Al}_2\text{Cl}_6$):

At high temperatures, aluminium chloride exists as monomeric $\text{AlCl}_3$. However, Aluminium is a Group 13 element, and in $\text{AlCl}_3$, it only has six valence electrons (it is electron deficient).

Two $\text{AlCl}_3$ molecules join together to form a dimer, $\text{Al}_2\text{Cl}_6$. Two chlorine atoms (from two separate $\text{AlCl}_3$ units) donate a lone pair to the electron-deficient aluminium atom in the other unit, forming two dative bonds that bridge the two aluminium atoms.

Quick Review: Look for a lone pair donor and an electron-deficient acceptor (often a metal ion or a molecule like $\text{AlCl}_3$) to identify dative bonds.

4. Delving Deeper: Orbital Overlap and Hybridisation (Syllabus 3.4.2)

Covalent bonds don't just "appear"; they are formed by the overlap of atomic orbitals. This overlap results in two distinct types of covalent bonds: sigma ($\sigma$) and pi ($\pi$) bonds.

4.1. Sigma ($\sigma$) Bonds

A sigma bond ($\sigma$) is formed by the direct (head-on) overlap of orbitals between the bonding atoms.

This can be s-s overlap (e.g., in $\text{H}_2$).
This can be s-p overlap (e.g., in $\text{HCl}$).
This can be head-on p-p overlap (e.g., in $\text{Cl}_2$).
All single bonds are sigma bonds.
Sigma bonds are very strong because the electron density is concentrated directly along the axis linking the two nuclei.

4.2. Pi ($\pi$) Bonds

A pi bond ($\pi$) is formed by the sideways overlap of adjacent p orbitals.

The overlap occurs above and below the line of the sigma ($\sigma$) bond.
Pi bonds are found in multiple bonds (double and triple bonds).
A double bond consists of one $\sigma$ bond and one $\pi$ bond.
A triple bond consists of one $\sigma$ bond and two $\pi$ bonds.
Pi bonds are generally weaker than sigma bonds because the sideways overlap is less effective than head-on overlap.

4.3. Hybridisation (Syllabus 3.4.2(c))

When atoms form bonds, their original s and p orbitals often mix to create new, identical orbitals called hybrid orbitals. This allows for maximum overlap and minimizes repulsion, which dictates the molecule's shape (covered further in Syllabus 3.5).

$\text{sp}^3$ hybridisation: Mixing one s orbital and three p orbitals to form four equivalent $\text{sp}^3$ orbitals.
Found in: Single bonds (e.g., Methane $\text{CH}_4$, Ethane $\text{C}_2\text{H}_6$). All bonds are $\sigma$.
$\text{sp}^2$ hybridisation: Mixing one s orbital and two p orbitals to form three equivalent $\text{sp}^2$ orbitals, leaving one unhybridised p orbital.
Found in: Double bonds (e.g., Ethene $\text{C}_2\text{H}_4$). The double bond contains one $\sigma$ bond (from $\text{sp}^2$ overlap) and one $\pi$ bond (from unhybridised p orbital overlap).
$\text{sp}$ hybridisation: Mixing one s orbital and one p orbital to form two equivalent $\text{sp}$ orbitals, leaving two unhybridised p orbitals.
Found in: Triple bonds (e.g., Nitrogen $\text{N}_2$, Hydrogen Cyanide $\text{HCN}$). The triple bond contains one $\sigma$ bond (from $\text{sp}$ overlap) and two $\pi$ bonds (from unhybridised p orbital overlap).

Don't worry if this seems tricky at first! The key is remembering the relationship:
Single bond $\rightarrow$ 1 $\sigma$ bond $\rightarrow$ $\text{sp}^3$
Double bond $\rightarrow$ 1 $\sigma$ bond, 1 $\pi$ bond $\rightarrow$ $\text{sp}^2$
Triple bond $\rightarrow$ 1 $\sigma$ bond, 2 $\pi$ bonds $\rightarrow$ $\text{sp}$

5. Bond Energy and Bond Length (Syllabus 3.4.3)

The properties of covalent molecules, especially their reactivity, are closely related to how strong and how long their bonds are.

5.1. Definitions

Bond Energy (or Bond Enthalpy)

Definition: The energy required to break one mole of a particular covalent bond in the gaseous state.

Bond breaking is an endothermic process (requires energy input), so bond energies are always given as positive values ($\Delta H$ positive).

Note: Many bond energy values are averages, as the energy required to break a specific bond (e.g., a C-H bond) might vary slightly depending on the molecule it is in.

Bond Length

Definition: The internuclear distance of two covalently bonded atoms. This is the distance between the centres of the two atoms.

5.2. Comparing Reactivity

We use bond energy and bond length to compare the reactivity of molecules:

Bond Strength and Length Relationship: Generally, shorter bonds are stronger bonds (require more energy to break).
Multiple Bonds: Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.
Reactivity: Molecules containing weaker bonds (low bond energy) are usually more reactive because less energy is needed to start a reaction (i.e., breaking the existing bonds).

Example: The carbon-carbon double bond in ethene ($\text{C}_2\text{H}_4$) is made up of a strong $\sigma$ bond and a weaker $\pi$ bond. The relatively weak $\pi$ bond is easily broken, making ethene much more reactive than ethane ($\text{C}_2\text{H}_6$), which only contains strong $\sigma$ bonds.

Key Takeaway: Shorter, stronger bonds typically lead to less reactive molecules, while longer, weaker bonds increase reactivity.

Chapter Summary: Covalent and Dative Bonds

We learned that covalent bonds involve sharing electrons, defined by the electrostatic attraction between nuclei and the shared pair. Period 3 elements can be rebels and expand their octet using d-orbitals.
We also saw that dative bonding is unique because one atom donates both electrons.
Finally, we explored bond types: $\sigma$ bonds (head-on overlap, strong) form the core, and $\pi$ bonds (sideways overlap, weaker) appear in multiple bonds, often requiring orbital hybridisation ($\text{sp}^3, \text{sp}^2, \text{sp}$) to maximise bonding efficiency. These structures dictate the bond energy and length, which in turn determine chemical reactivity. Good job on tackling these core concepts!