Chapter 03: Chemical Bonding

Chapter 03: Chemical Bonding

A-Level H2 Chemistry: Chapter 3
Chemical Bonding

Concise Summary of Properties

The following table is a summarised compilation of the physical properties of the different types of structures for easier revision.

You can think of chemical bonding as a spectrum, where covalent and ionic bonding are on both ends of the spectrum. Most compounds typically fall somewhere in between the two extreme types of bonding due to polarisation of either covalent bonds or ions.

The Grey Area in Chemical Bonding

A perfect covalent bond would involve the equal sharing of electrons between the two covalently-bonded atoms. Hence, polarisation of covalent bonds can be understood as the unequal sharing of electrons.
On the other extreme, a perfect ionic bond involves the total transfer of electrons from one ion to the other. Thus the polarisation of ions is the incomplete transfer of electrons.

Covalent Character in Ionic Bonds

Let’s dive a little deeper into ionic bonds that exhibit covalent character.
In such bonds, the cation (A⁺) attracts the negative charge of the anion (B⁻), hence distorting the electron cloud of the anion. As a result, electrons from the cation are not completely transferred to the anion.

Some factors lead to a higher degree of covalent character in ionic bonds.

• Cation with a small ionic radius with a high charge. It is said to have a high charge density with high polarising power.
• Anion with a large ionic radius. This means the electron cloud of the anion is bigger, thus easily polarisable.

Ionic Character in Covalent Bonds

Some bonds are classified as covalent but still have a degree of ionic character.

Difference in electronegativity. The larger the difference in negativity between the two atoms, the more likely that the shared pair(s) of electrons are pulled much closer to the more electronegative atom. As a result, there are partial charges in the molecule and it becomes polar. Hence the sharing of electrons between the two atoms is no longer equal.

AlF₃ and AlBr₃ may appear to be similar compounds with the same structure and bonding. Don’t be fooled, as AlF₃ has a giant ionic lattice structure with strong electrostatic forces of attraction between positively charged Al³⁺ and negatively charged F⁻ ions. Meanwhile, AlBr₃ has a simple molecular structure with instantaneous dipole-induced dipole forces of attraction between molecules.

Watch out for Al when you approach bonding and structure questions involving it. Due to the high charge density of Al³⁺, it easily polarises the electron cloud of Br⁻ anion to such a great extent that the bonding between the two ions becomes covalent. However, since the ionic radius of F⁻ is much smaller than Br⁻, it is less easily polarised by Al³⁺. Hence, the bonding between F⁻ and Al³⁺ remains ionic. As such, take note that due to the different bonding and structure in AlF₃ and AlBr₃, they exhibit different physical and chemical properties.

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory helps determine the shapes and bond angles of covalent molecules.

The three basic principles are:

• Electron pairs repel each other and arrange themselves as far apart as possible to maximise stability and minimise electrostatic repulsion
• Orbitals containing lone pairs are larger than those containing bond pairs/odd electrons hence they have greater repulsive effect than bond pairs. In other words, lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion
• Repulsion between electron pairs is increased by an increase in the electronegativity of the central atom.

Worked Example 1: 2020 SAJC H2 Chemistry Prelim Paper 3, Q5(c)(ii)

Platinum complexes have been researched extensively for their anti-cancer properties. One such platinum complex is cis-platin, which can be synthesised from PtCl₄. The central atom, platinum, in PtCl₄ is surrounded by six pairs of electrons arranged in an octahedral shape. An octahedral arrangement is shown in Fig. 5.1.

Two different molecular arrangements of PtCl₄ are possible. Draw clear diagrams of these two arrangements. Hence, apply the principles of the VSEPR theory to discuss the relative stabilities of these two possible arrangements.

Solution and Detailed Explanation:

Don’t be intimidated by the application context of the question. As long as you pick out the key information, it is solvable.

Part I of the question requires you to draw 2 possible arrangements for the compound PtCl₄. Since there are 4 Cl atoms surrounding the Pt central atom, there should be 4 sigma bond pairs and the remaining 2 electron pairs are lone pairs.
The two possible arrangements are shown in the diagram above.

Part II of the question requires you to discuss the relative stabilities of the above 2 arrangements drawn with the use of VSEPR theory. Recall the 3 basic principles of VSEPR theory and apply accordingly. Use the diagram drawn in Part I as a visual aid. The question also requires a comparison, so discuss the stability of both arrangements.

The square planar arrangement on the left is more stable.

Reason: For the one on the left, the two lone pairs of electrons are on opposite sides, 180° away, which minimises lone pair-lone pair repulsion.
For the one on the right, the two lone pairs of electrons are next to each other, 90° away, which experiences greater lone pair-lone pair repulsion.

Sigma and Pi Bonds

You may struggle to understand this portion of the chapter due to its abstract nature, as picturing the bonds in your head is not easy. With the help of some pictures we will guide you in understanding sigma and pi bonds better.

Many students tend to think that both the sigma and pi bonds are equal in strength.
But a pi bond is weaker than a sigma bond. This is because of the less effective side-on overlap compared with head-on overlap.

Dative Bonding

It is crucial that students remember the conditions for the formation of dative bonds. By using the ammonia-boron trifluoride complex as an example, students can check whether a compound can form a dative bond using the following steps.

Step 1: The donor atom should have a lone pair of electrons. In this case, the N atom is the donor atom: N has five valence electrons, three of which are used to bond with H atoms, leaving one lone pair.

Step 2: The acceptor atom must have an empty orbital in its valence shell. In this case, the B atom has an empty orbital.
Note: Acceptor atoms are electron-deficient species. More information about them appears below.

Step 3: The donor atom donates its lone pair to the acceptor atom’s empty valence orbital, represented by the arrow between N and B.

Ammonia–boron trifluoride complex

Students are often familiar with the octet rule: atoms tend to bond in ways that give them eight electrons in their valence shell, like a noble gas. There are exceptions:

1. Odd-Electron Species

• Such species are known as radicals and are very reactive due to their unpaired electrons.
• Common examples include NO, NO₂, and ClO₂.

2. Electron-Deficient Species

• Such species are molecules with fewer than eight valence electrons. They can form dative bonds with other molecules.
• Common examples include BF₃, AlCl₃, and BeCl₂.

3. Species with More Than Eight Valence Electrons

• Such species can have more than eight valence electrons around the central atom. They typically involve elements from Period 3 onward.
• Common examples include SF₆, PF₅, and XeF₄.

Intermolecular Forces

The following table is a summary of the intermolecular forces for simple covalent molecules:

• Polar molecules with a net dipole moment have both id-id and pd-pd attractions. Their boiling points are generally higher than those of comparable non-polar molecules when they have similar numbers of electrons.
• When intramolecular hydrogen bonds are present, intermolecular hydrogen bonding can be less extensive.