Intermolecular Forces Topic 4.3

Intermolecular Forces Intramolecular forces – refer to the forces that hold atoms together within molecules or formula units; a chemical bond. Intermolecular forces- refer to the forces between particles in a substance. These forces are very weak compared to intramolecular forces. For example, 927 kJ of energy is required to decompose one mole of water vapor into H and O atoms. However, only 40.7 kJ are required to convert one mole of liquid water into steam at 100 o C.

Intermolecular Forces

If it were not for the existence of intermolecular attractions, condensed phases (liquids and solids) would not exist. These are the forces that hold the particles close to one another in liquids and solids. Changing the state of a substance from solid to liquid or liquid to gas involves separating particles by overcoming the forces between them. The stronger the intermolecular forces, the more energy needed to separate the particles in substance, and so the higher the substance’s melting and boiling points

Intermolecular Forces

We will consider three types of intermolecular forces: – Van der Waals’ forces – Dipole-Dipole attraction – Hydrogen bonding

Van der Waal’s forces Non-polar molecules such as chlorine (Cl 2 ) have no separation of charges within their bonds because the electrons are shared equally. However, because electron density at any one moment may be greater over one atom than the other, some separation of charge may occur, this is known as a temporary or instantaneous dipole. This instantaneous dipole may influence the electron distribution in the neighboring molecule by causing an induced dipole.

Van der Waals’ Forces

As a result, weak forces of attraction, known as van der Waals’ forces, will occur between opposite ends of these two temporary dipoles in the molecules. Van der Waals’ forces are present between all types of molecules, but are the only kind of intermolecular forces present among symmetrical non-polar substances and monatomic species such as noble gases.

Van der Waals’ Forces They are the weakest form of intermolecular forces. Their strength increases as the number of electrons within a molecule increases, that is, with increasing molecular mass. Substances that are held together by van der Waals’ forces generally have low melting and boiling points, because relatively little energy is required to break the forces and separate the molecules from each other.

Van der Waals’ Forces Boiling point data show how the strength of van der Waals’ forces increases with increasing molecular mass ElementMrMr Boiling Point/ o CState at room temperature F2F Gas Cl Gas Br Liquid I2I Solid

Dipole-Dipole Attraction Molecules such as HCl have a permanent separation of charge within their bonds as a result of the difference in electronegativity of the atoms. The Cl end has a partial negative charge (δ - ) while the H end has a partial positive charge(δ + ). This is known as a permanent dipole, and results in opposite charges in neighboring molecules attracting each other, generating a force known as a dipole-dipole attraction.

Dipole-Dipole Attraction

The strength of the dipole-dipole attraction depends on the degree of polarity within the bond. It will decrease as the degree of polarity with the molecule decreases. For example, it will decrease in strength in the order: HCl>HBr>HI. Dipole-dipole forces are stronger than van der Waals’ forces, consequently, the melting and boiling points of polar compounds are higher than those of non-polar substances of comparable molecular mass.

Hydrogen Bonding When a molecule contains hydrogen covalently bonded to a very electronegative atom (fluorine, nitrogen, or oxygen), these molecules are attracted to each other by a strong type of intermolecular force called hydrogen bonding. The hydrogen bond is in essence a particular case of dipole-dipole attraction.

Hydrogen Bonding

The large electronegativity difference between hydrogen and the bonded fluorine, oxygen, or nitrogen, results in the electron pair being pulled away from the hydrogen. Given its small size and the fact that it has no other electrons to shield the nucleus, the hydrogen now exerts a strong attractive force on a lone pair of electrons in the electronegative atom of a neighboring molecule. This is the hydrogen bond.

Hydrogen Bonding

Hydrogen bonds are the strongest form of intermolecular attraction. Consequently, they cause boiling points of substances that contain them to be significantly higher than would be predicted from their molecular weight.

Hydrogen Bonding In all four groups there is an observable trend of boiling point increasing down the group as molecular weight increases. The anomalies are NH 3, HF and H 2 O which have much higher boiling points than expected from their molecular weight. T his can only be explained by the presence of hydrogen bonding in these molecules

Hydrogen Bonding Likewise, when we compare the boiling points of some organic molecules that similar or equal molecular weights, we find a higher value where hydrogen bonding occurs between the molecules. For example: CH 3 -O-CH 3 CH 3 -CH 2 -O-H MW= 46 Does not form hydrogen bonding Boiling point = -23 o C Forms hydrogen bonds Boiling point = 79 o C

Relative Strength of Intermolecular Forces H-bonding > dipole-dipole > London With molecules possessing similar molecular weights, the molecule having the stronger IMF will usually have the larger bp/mp. This is especially true when the IMF is H-bonding.

Relative Strength of Intermolecular Forces With molecules in which one of the atoms has been changed with another atom within the same group, the molecule having the larger formula weight will usually have the larger bp/mp, except when the IMF is H-bonding. Thus, NH 3 (IMF – H-bonding) will have a higher bp/mp than PH 3 (IMF- dipole-dipole) even though NH 3 has a smaller molecular weight. However, AsH 3 will have a higher bp/mp than PH 3 because both have the same IMF, but AsH 3 has a larger molecular weight.

Hydrogen Bonding Water can form up to four hydrogen bonds because it has two lone pairs on the oxygen atom. Liquid water will contain fewer than this number, but in the solid form, ice, each water molecule forms 4 hydrogen bonds. The result is a tetrahedral arrangement that holds the molecules a fixed distance apart, forming a fairly open structure which is less dense than liquid water, allowing ice to float in water.

Hydrogen Bonding