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.
|
Physical
Properties
Physical State at r.t.p |
|---|
|
Structures |
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Giant Metallic Lattice |
|
Solid Exception: Mercury is a liquid |
|
Giant Ionic Lattice |
Solid |
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Giant Covalent Molecule |
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Solid |
|
Simple Covalent Molecule |
|
Mostly liquids/gases, while for some cases, they are solids |
|
Physical
Properties
Melting and Boiling Point |
|---|
|
Structures |
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Giant Metallic Lattice |
|
High melting and boiling point
|
|
Giant Ionic Lattice |
High melting and boiling point
|
|
Giant Covalent Molecule |
|
High melting and boiling point
|
|
Simple Covalent Molecule |
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Low melting and boiling point
|
|
Physical
Properties
Solubility |
|---|
|
Structures |
|
Giant Metallic Lattice |
|
Insoluble in both polar & non-polar solvents |
|
Giant Ionic Lattice |
Depends on lattice energy and hydration energy |
|
Giant Covalent Molecule |
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Insoluble in all solvents |
|
Simple Covalent Molecule |
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solvents Depends on strength of solute-solute, solvent-solvent and solute-solvent interactions |
|
Physical
Properties
Electrical Conductivity |
|---|
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Structures |
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Giant Metallic Lattice |
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Good conductors in all states |
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Giant Ionic Lattice |
Good conductors in liquid/molten and aqueous state Non-conductors in solid state |
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Giant Covalent Molecule |
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Non-conductors except for graphite |
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Simple Covalent Molecule |
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Non-conductors in all states Some exceptions: molecules that ionise in solution |
|
Physical
Properties
Hardness |
|---|
|
Structures |
|
Giant Metallic Lattice |
|
Hard but malleable and ductile |
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Giant Ionic Lattice |
Hard and brittle |
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Giant Covalent Molecule |
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Hard except for graphite which is soft and slippery |
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Simple Covalent Molecule |
|
Soft |
| Physical Properties |
Structures | |||
|---|---|---|---|---|
| Giant Metallic
Lattice |
Giant Ionic
Lattice |
Giant Covalent
Molecule |
Simple Covalent
Molecule |
|
| Physical
State at r.t.p. |
Solid Exception: Mercury is a liquid |
Solid |
Solid |
Mostly liquids/gases, while for some cases, they are solids |
| Melting and
Boiling Point |
High melting and boiling point - Due to the breaking of strong metallic bonds |
High melting and boiling point - Due to the breaking of strong ionic bonds |
High melting and boiling point - Due to the breaking of strong covalent bonds |
Low melting and boiling point - Due to the breaking of weak intermolecular forces of attraction |
| Solubility |
Insoluble in both polar & non-polar solvents |
Depends on lattice energy and hydration energy |
Insoluble in all solvents |
Depends on strength of solute-solute, solvent-solvent and solute-solvent interactions |
| Electrical
Conductivity |
Good conductors in all states |
Good conductors in liquid/molten and aqueous state Non-conductors in solid state |
Non-conductors except for graphite |
Non-conductors in all states Some exceptions: molecules that ionise in solution |
| Hardness |
Hard but malleable and ductile |
Hard and brittle |
Hard except for graphite which is soft and slippery |
Soft |
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.
- 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.
Not far off, we have bonds which are still classified as covalent bonds, but with a degree of ionic character
Ionic character increases with the following factor:
A Common Trap:
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.
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.
Ionic Character in Covalent Bonds
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.
Why is that so?
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
The Valence Shell Electron pair Repulsion theory (VSEPR) theory is one of the most important concepts in chem bonding as this theory helps us to determine the shape and bond angles of covalent molecules.
The 3 basic principles that students should take note to remember 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.
Repulsion between electron pairs is increased by an increase in the electronegativity of the central atom.
| Total number of electron pairs | Number of bond pairs | Number of lone pairs | Shape/Bond angle | Example |
|---|---|---|---|---|
| 2 | 2 | 0 |
Linear 180°
|
N₂O |
| 3 | 3 | 0 |
Trigonal planar 120°
|
BF₃ |
| 3 | 2 | 1 |
Bent 110° - 120°
|
SO₂ |
| 4 | 4 | 0 |
Tetrahedral 109.5°
|
CH₄ |
| 4 | 3 | 1 |
Trigonal pyramidal 107°
|
SOCl₂ |
| 4 | 2 | 2 |
Bent 105°
|
H₂O |
| 5 | 5 | 0 |
Trigonal bipyramidal
|
PCl₅ |
| 5 | 4 | 1 |
See-saw 120°
|
SF₄ |
| 5 | 3 | 2 |
T-shaped 90°
|
ClF₃ |
| 5 | 2 | 3 |
Linear 180°
|
XeF₂ |
| 6 | 6 | 0 |
Octahedral 90°
|
SF₆ |
| 6 | 5 | 1 |
Square pyramidal 90°
|
BrF₅ |
| 6 | 4 | 2 |
Square planar 90°
|
XeF₄ |
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Total number of electron pairs: 2 Number of bond pairs: 2 Number of lone pairs: 0 Shape/Bond angle : Linear 180°
Example : N₂O |
|
Total number of electron pairs: 3 Number of bond pairs: 3 Number of lone pairs: 0 Shape/Bond angle : Trigonal planar 120°
Example :BF₃ |
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Total number of electron pairs: 3 Number of bond pairs: 2 Number of lone pairs: 1 Shape/Bond angle : Bent 110° - 120°
Example : SO₂ |
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Total number of electron pairs: 4 Number of bond pairs: 4 Number of lone pairs: 0 Shape/Bond angle : Tetrahedral 109.5°
Example : CH₄ |
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Total number of electron pairs: 4 Number of bond pairs: 3 Number of lone pairs: 1 Shape/Bond angle : Trigonal pyramidal 107°
Example : SOCl₂ |
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Total number of electron pairs: 4 Number of bond pairs: 2 Number of lone pairs: 2 Shape/Bond angle : Bent 105°
Example : H₂O |
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Total number of electron pairs: 5 Number of bond pairs: 5 Number of lone pairs: 0 Shape/Bond angle : Trigonal bipyramidal
Example : PCl₅ |
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Total number of electron pairs: 5 Number of bond pairs: 4 Number of lone pairs: 1 Shape/Bond angle : See-saw 120°
Example : SF₄ |
|
Total number of electron pairs: 5 Number of bond pairs: 3 Number of lone pairs: 2 Shape/Bond angle : T-shaped 90°
Example : ClF₃ |
|
Total number of electron pairs: 5 Number of bond pairs: 2 Number of lone pairs: 3 Shape/Bond angle : Linear 180°
Example : XeF₂ |
|
Total number of electron pairs: 6 Number of bond pairs: 6 Number of lone pairs: 0 Shape/Bond angle : Octahedral 90°
Example : SF₆ |
|
Total number of electron pairs: 6 Number of bond pairs: 5 Number of lone pairs: 1 Shape/Bond angle : Square pyramidal 90°
Example : BrF₅ |
![bond-angle](https://www.achieversdream.com.sg/wp-content/uploads/2021/06/Square-pyramidal-90.png")
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Total number of electron pairs: 6 Number of bond pairs: 4 Number of lone pairs: 2 Shape/Bond angle : Square planar 90°
Example : XeF₄ |
Worked Example 1:
[2020 SAJC H2 Chem Prelim Paper 3 Qn 5(c)(ii)]
Platinum complexes have been researched extensively due to its 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.
Therefore, the only 2 possible arrangements are:
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. Do note that the question also requires you to do a comparison, thus you must mention about the stability of BOTH drawn arrangements.
The square planar arrangement/the one 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
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.
| Sigma bond (σ) | Pi bond (π) |
|---|---|
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'side-by-side' (π) overlap 
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| Sigma bond (σ) |
|---|
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| Pi bond (π) |
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'side-by-side' (π) overlap'
|
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 as compared to the “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 few steps.
Step 1: The donor atom should have a lone pair of electrons (In this case, the N atom would be the donor atom due to having 2 unused electrons - N has 5 electrons in total in which 3 are used to undergo covalent bonding with H atoms, leaving a lone pair of electron)
Step 2 : The acceptor atom must have an empty orbital in its valence shell (In this case, the B atom has an empty orbital - meaning there is no electrons found in that orbital)
Note: Acceptor atoms are electron deficient species (More information regarding
electron deficient species can be found below)
Step 3 : The donor atom then donates the lone pair of electrons to fill the empty valence orbital of acceptor atom (represented by the arrow between the N and B atom)
Ammonia-borontrifluoride complex
Students are often familiar with the octet rule, where it states that atoms have a tendency to have 8 electrons in their valence shell and carry out bonding in such a way that fulfills the octet rule in order to achieve the electronic configuration of a noble gas.
However, many students tend to forget that there are EXCEPTIONS to this rule, mainly:
1. Odd electron species
- Such species are known as radicals and are very reactive due to their unpaired electrons.
- Common examples include NO, NO₂, ClO₂
2. Electron deficient species
- Such species are molecules with fewer than 8 valence electrons. They hence have the potential to form dative bonds with other molecules
- Common examples include BF₃, AlCl₃, BeCl₂
3. Species with more than 8 valence electrons
- Such species have the ability to expand their octet configuration and are elements found in period 3 onward. This is since period 3 elements and onwards have energetically accessible empty d-orbitals to accommodate more than 8 electrons.
- Common examples are SF₆, PF₅, XeF₄
Intermolecular Forces
The following table is a summary of the intermolecular forces for simple covalent molecules:
| Type of Intermolecular Forces |
|---|
|
Predominant Type of IMF Present |
|
Instantaneous dipole-induced dipole (id-id) attractions (weakest) |
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Between all molecules (both non-polar and polar) |
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Permanent dipole-permanent dipole (pd-pd) attractions |
Between polar molecules (Tip: molecules with net dipole moment) |
|
Hydrogen Bonding (strongest) |
|
Between a molecule with H bonded to F, O, or N and a molecule with lone pair of electrons on F, O, or N |
| Type of Intermolecular Forces |
|---|
|
Factors Affecting IMF |
|
Instantaneous dipole-induced dipole (id-id) attractions (weakest) |
|
1. No. of electrons - The more electrons, the larger the atom’s electron cloud and the more easily the electron cloud is polarised,resulting in stronger id-id attractions
2. Surface area - straight-chain (more open structure) vs branched chain molecules (cluttered) Id-id attractions are more extensive for a straight-chain molecule due to its larger surface area. |
|
Permanent dipole-permanent dipole (pd-pd) attractions |
Magnitude of dipole moment. |
|
Hydrogen Bonding (strongest) |
|
1. No. of hydrogen bonds per molecule (which depends on the number of lone pairs of electrons on the F, O or N atom.The more lone pair the electronegative atom has, the stronger the hydrogen bonds.) 2. Electronegativity Difference. |
| Type of Intermolecular Forces | Instantaneous dipole-induced dipole (id-id) attractions (weakest) | Permanent dipole-permanent dipole (pd-pd) attractions | Hydrogen Bonding (strongest) |
|---|---|---|---|
| Predominant Type of IMF Present | Between all molecules (both non-polar and polar) | Between polar molecules (Tip: molecules with net dipole moment) | Between a molecule with H bonded to F, O, or N and a molecule with lone pair of electrons on F, O, or N |
| Factors Affecting IMF | 1. No. of electrons
- The more electrons,
the larger the atom’s
electron cloud and
the more easily the
electron cloud is
polarised,resulting in
stronger id-id
attractions
(Note: Do not use higher Mr although it implies a larger number of electrons. Use ‘number of electrons’.) 2. Surface area - straight-chain (more open structure) vs branched chain molecules (cluttered) Id-id attractions are more extensive for a straight-chain molecule due to its larger surface area. |
Magnitude of dipole moment. | 1. No. of hydrogen
bonds per
molecule (which
depends on the
number of lone
pairs of electrons
on the F, O or N
atom.The more
lone pair the
electronegative
atom has, the
stronger the
hydrogen bonds.)
2. Electronegativity Difference. |
Special Things to Note:
- Do take note that polar molecules (molecules with net dipole moment) consist of both id-id and pd-pd attractions. The boiling point of polar molecules is only higher than non-polar molecules when they have a comparable number of electrons.
- When intramolecular hydrogen bonds are present, the intermolecular hydrogen bonds in the molecules are less extensive.
Continue Reading on Chapter 4: The Gaseous State