Intramolecular Bonds and Intermolecular Forces
This is part of preliminary HSC Chemistry course under the topic of Bonding
HSC Chemistry Syllabus
- Investigate the role of electronegativity in determining the ionic or covalent nature of bonds between atoms
- Investigate the differences between ionic and covalent compounds through:
- Examining the spectrum of bonds between atoms of varying degrees of polarity with respect to their constituent elements' positions on the periodic table.
- Explore the similarities and differences between the nature of intermolecular and intramolecular bonds and the strength of the forces associated with each, in order to explain the:
- Physical properties of elements
- Physical properties of compounds (ASCH020, ASCH055, ASCH058)
Intramolecular Bonds and Intermolecular Forces
In the periodicity section, we discussed what electronegativity was and what the general trend for electronegativity was on the periodic table. Electronegativity is a measurement of an atom’s tendency to attract an electron to form bonds and this property arises due the requirement for atoms to complete their valence shell in accordance with the octet rule.
The table below demonstrates the electronegativity values of elements on the periodic table. Notice that most group 18 noble gases do not have electronegativity values. This is because they have a complete valence shell and are mostly unreactive. The term that we give to describe the underactivity of a substances is “inert”. Krypton and Xenon in periods 4 and 5 respectively do have electronegativity values however due to the shielding effects which affect its outer shells, allowing it to form covalent bonds e.g., XeF3.
Intramolecular Bonds & Polarity
In Chemistry, an electric dipole moment is a measure of the polarity of a molecule. These dipole moments occur where there is a separation of electric charge between molecules causing an imbalance. Where this imbalance occurs, we classify molecules as being ‘polar’ and can use the terms polar and non-polar to label bonds
Intramolecular Bond Types:
- Non-polar Covalent
Non-polar covalent bonds occur where there is an equal sharing or symmetry between the two atoms of the electrons in the bond. Examples of covalent molecules containing non-polar bonds include the diatomic compounds `Cl_2`, `H_2`, `F_2`, `N_2`, `O_2`, and others such as `CH_4`
For a covalent bond to be classified as being non-polar, the electronegativity difference between atoms would have to less than 0.5.
- Polar Covalent
Polar covalent bonds occur when there is an unequal sharing of the electrons in the bond leading to an asymmetric molecule. The range for which covalent bonds are polar is quite large, ranging from 0.5 to almost 2.0.
These polar bonds occur in compounds including `NH_3` and `H_2O` which contain N–H bonds with an electronegativity difference of approximately 1.2. Examples range from barely polar bonds like the C-Cl bond with an electronegativity difference of 0.5 to H–F with an electronegativity difference of 1.8
Ionic bonds occur when electrons are completely transferred from one atom to another to form a cation and an anion. This is unlike the sharing of electrons which can be found in covalent bonds. Ionic bonds generally contain elements with an electronegativity difference greater than 2.0, and examples of substances which are made up of ionic bonds include NaCl and `MgCl_2`.
Types of Intermolecular Forces
Covalent and Ionic bonds are the forces which bind atoms together to form molecules and compounds. These bonds are a result of electrostatic attraction effects which exist between atoms that exist within molecules and thus they are called intramolecular bonds.
Dispersion forces, also known as London dispersion forces, are the weakest form of intermolecular force. These forces occur as a result of an uneven distribution in electrons around the nucleus and between neighbouring atoms. As a result of the uneven movement of electrons, one side of an atom can temporarily be more positive than the other size to create a temporary dipole (the phenomenon of there being an uneven distribution of electrons across a molecule). However because the electron movement within the molecule is random, the dipole constantly breaks and reforms while simultaneously inducing dipoles in adjoining molecules.
In general, the larger the molecule, the larger the dispersion force between it and its neighbours will be. The size of the dispersion force is also affected by the surface layer of the molecule. Dispersion forces are the intermolecular forces which bind the layers of graphite together in its covalent network structure.
Covalent substances often have dispersion forces as their main mode of intermolecular force. These types of substances usually demonstrate lower tensile strengths, boiling points, and melting points, characteristics of which can be attributed to the existence of weak dispersion forces. However, some covalent substances such as a range of several hydrocarbon fuels, can exist in the solid and liquid states due to the sheet amount of dispersion forces which exist within them
Dipole – Dipole Forces
Dipole – dipole forces are electrostatic ally attractive forces which bind molecules with permanent dipoles. Where permanent dipoles exist between multiple molecules, the negative ends attract the positive ends of the neighbouring molecule's dipoles. These forces are stronger than dispersion forces where molecules have similar weights.
The hydrogen bond can be thought of as being a special form of a dipole-dipole interaction because the atoms which the hydrogen bonds to (FON) are very electronegative.The electronegative atom attracts the electron in the hydrogen atom, creating a partial positive charge on the hydrogen atom and a partial negative charge on the electronegative atom. As a result of the large electronegativity difference, the electronegative atom also shares its own electrons with the positive hydrogen atoms. This consequently allows hydrogen bonds to have covalent intramolecular characteristics, hence the name of the intermolecular force being a hydrogen "bond"