Chemistry : Bond polarity and intermolecular forces

Bond polarity

Electronegativity describes the tendency of an atom or a functional group to attract e^- (or electron density) towards itself and thus the tendency to form negative ions. The tendency of electronegativity across the periodic table is that it decrease down the period and increase along the group. The most electronegative atom is F, O and N, the least electronegative atom is alkali metal (Fr, Cs,…), so they are often referred as electropositive. Noble gas has electronegativity zero since it's assumed to be can't form bonds.

The trend across the group can be explained by increasing in atomic mass causes increase in p⁺, so that the attraction on e^- bond pair increases.

The trend across the period can be explained by the increase in e^- shell cause distance between nucleus and the e^- bond pair increase, then the attraction decreases.

Concept of e^- cloud: e^- moving around the atom rather staying at a certain position. A electron cloud show the distribution that e^- has a certain probability to appear there.

Consider a bond pair, X-Y. If X has a higher electronegativity, the e^- bond pair is drawn towards X, then the electron cloud of the molecule is distorted to X (since e^- is more likely to be appeared there), and X receives a partial negative charge, denoted by δ-, while Y receives a partial positive charge, denoted by δ+. The larger difference in electronegativity between X and Y, the higher partial charge received.

1)       Pure ionic bond (electronegativity diff more than 2.0), the e^- bond pair is greatly drawn to the electronegative atom, and the it receives a whole positive / negative charge. e.g., NaCl (Na is electropositive while Cl is very electronegative)

2)       Ionic bond with covalent character: the e^- bond pair is greatly drawn, but the charge is not separated greatly as pure ionic bond. e.g., MgCl₂

3)       Covalent bond with ionic character (polar covalent bond): the e^- bond pair is slightly goes near to the electronegative atom. e.g., AlCl₃, HCl, H₂O

4)       Pure covalent bond (electronegativity diff less than 0.5): the e^- bond pair is nearly equally shared. E.g. C-H bond, all diatomic molecules like H₂, Cl₂.

Properties concerning different bonds:

-          Pure ionic substances matter dissolves in water and pH is still 7.

-          Ionic bond with covalent character / covalent bond with ionic character: also dissolves in water, but pH<7, due to dissociation. (Attracted by δ-O)

Dipole moment μ=Qd with unit D, where Q is the partial charge and d is the bond length. A larger diff in electronegativity mean larger Q and hence larger dipole moment. Dipole moment can be represented by  in the structural formula by the following arrow, pointing from δ+ to δ-. For example, the dipole moment in HCl is 1.1D.

The net dipole moment of a molecule is given by the vector sum of all dipole moment exist in the molecule. Note that even there's dipole moment in the molecule, the dipole moments can eliminate each other and cause no net dipole moment. For example, the dipole moment in CO₂ is cancelled out each other, and it has no net dipole moment, similar to BF₃ and CCl₄. We can call these molecules as non-polar.

When there's a net dipole moment like NH₃, CHCl₃, H₂O, SO₂, we call the molecules polar.

By the rule of like dissolve like, those polar molecules are very soluble in water while non-polar molecules like CO₂ are very insoluble in polar solvents like water.

In a polar molecule, the e^- cloud is permanently distorted, and there's a partial charge. We say that a permanent dipole is formed.

In a non-polar molecule, the e^- sometimes maybe distorted a bit (significantly less than permanent distortion), an instantaneous δ+ and δ- is created at the moment, and an instantaneous dipole is formed.

When a molecule with instantaneous/permanent dipole moves to a non-polar molecules, the e^- clouds in the non-polar molecules may be distorted due to attraction/repulsion of charge from the dipole, and create δ+ and δ- in the non-polar molecules, then induced dipole is formed.

Intermolecular forces (often referred as wan der Waals' force)

1)       Instantaneous dipole-induced dipole attraction, is the weakest among the three types of intermolecular forces but it appears in all molecules.

2)       Permanent dipole-induced dipole attraction appears in mixture or solution where non-polar molecules are polarized by polar molecules through induction.

3)       Permanent dipole-permanent dipole attraction, strength about 1/100 of covalent bond

Hydrogen bond is a strong attractive force between H attached to an electronegative atom and the lone pair on another electronegative atom. Usually the electronegative atom is F, O or N.

e.g., HF, H₂O, CH₃OH and NH₃. Exceptional case: CHCl₃ since there're 3 Cl atoms. The hydrogen bond is very strong (1/10 of covalent bond) since the e^- is attracted by the electronegative atom in a great extent, then the proton of H is "naked" and has a very strong positive charge, which attracts the lone pairs. This explains the unusual high m.p./b.p. of hydrides of F, O and N than other atoms in group 5,6,7.

Factors affecting strength of intermolecular forces:

1)       Polarizability: For a larger molecules/atoms, the attraction on outermost electron is weaker, then the e^- cloud is more easily to be polarized. For example, m.p./b.p. of group 0 increases down the group since the size of e^- cloud increase, then polarizability increases, then the instantaneous dipole-induced dipole attraction increase.

2)       Shape of molecule: A straight chain allows greater contact among molecules, and hence greater attraction. E.g. b.p. of CH₃(CH₂)₃CH₃ is greater than CH₃C(CH₃)₂CH₃ though their molecular formula is the same (C₅H₁₂).

Water and ice: water attains highest density at 4˚C because ice adopts the open caged structure so that each H₂O molecules forms 4 hydrogen bonds, and hence it has a very low density. Water starts to adopt the open cages structure when it's cooled under 4˚C, so the density increases from 0˚C to 4˚C. When it's heated over 4˚C, molecules gains K.E. so that the average separation increase, and hence lower density.

Other properties relating to intermolecular forces

1)       Surface tension is caused by intermolecular force acting in all direction on the molecules, allowing it can form a tighten elastic surface. Liquids with stronger intermolecular forces tends to have greater surface tension.

2)       Viscosity depends on the strength of intermolecular forces and the tendency of molecules to become entangled with each other. Usually molecules with longer and branched chain is entangled easier. For example, propan-1,2,3-triol (glycerol) has a significantly higher viscosity than ethanol because of its branched shape and high ability to form hydrogen bonds. Yet ethanol has a higher viscosity than propanone since propanone can't form hydrogen bond.

3)       Melting and boiling point relates to the strength of intermolecular forces. For example, the boiling point of C₃H₈ < CH₃OCH₃ (polar) < CH₃OH (hydrogen bond).

4)       Ease of evaporation: molecules in liquids have to do work against atmospheric pressure and intermolecular force to vaporize. Stronger intermolecular force reduces the ease of vaporization.

5)       Solubility in water: molecules that can form hydrogen bond is greatly soluble since forming of hydrogen bond helps the dissolving process.

Polar molecules which do not form hydrogen bonds are fairly soluble since less hydrogen bond can be formed between them.

Non-polar molecules is insoluble in water due to difference in strength of intermolecular forces between water and the non-polar molecules.