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A Review of Molecular Geometry a Precursor to Understanding Intermolecular Forces. http://www2.gasou.edu/chemdept/general/molecule/quiz/frame4b.htm. http://www2.gasou.edu/chemdept/general/molecule/quiz/frame4a.htm. Intermolecular forces. Molecular attractions.
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A Review of Molecular Geometry a Precursor to Understanding Intermolecular Forces
Attractions between molecules are classified as
The relative size of these interactions is important so the relative effects are understood.
Relative strengths for the different interactions are listed here.
Hydrogen bonding >
400 kcal >
12-16 kcal >
2-0.5 kcal >
less than 1 kcal
Clearly normal covalent bonds are almost 40 times the strength of hydrogen bonds. Covalent bonds are almost 200 times the strength of dipole-dipole forces, and more than 400 times the size of London forces.
What is required of a molecule or molecules to allow for Dipole-Dipole forces?
Dipole-dipole forces - only polar covalent molecules have the ability to form dipole-dipole attractions between molecules. Polar covalent molecules act as little magnets, they have positive ends and negative ends which attract each other.
London forces exist in nonpolar molecules.
These forces result from temporary charge imbalances. The temporary charges exist because the electrons in a molecule or ion move randomly in the structure. The nucleus of one atom attracts electrons form the neighboring atom. At the same time, the electrons in one particle repel the electrons in the neighbor and create a short lived charge imbalance.
These temporary charges in one molecule or atom attract opposite charges in nearby molecules or atoms. A local slight positivecharged+ in one molecule will be attracted to a temporary slight d-negative charge in a neighboring molecule.
London forces - all molecules have the capability to form London forces. These are solely dependent on the surface area and the polarizability of the surface of the molecule. These are the only types of forces that non-polar covalent molecules can form. They result from the movement of the electrons in the molecule which generates temporary positive and negative regions in the molecule.
Melting Points and Boiling Points of Similar Substances with Increasing Formula Weights
As the size of the halogens increases, the melting and boiling points increase. The energy required to move and separate the molecules from one another increases as the size of the molecules increases. If it takes more energy to separate the molecules, the attractions between molecules must be greater. The types of intermolecular forces responsible for the increase in melting points and boiling points of these non-polar covalent compounds are called London forces or dispersion forces.
Hydrogen bonding is a unique type of intermolecular molecular attraction. There are two requirements.
The first is a covalent bond between a H atom and either F, O, or N. (These are the three most electronegative elements.)
The second is an interaction of the H atom in this kind of polar bond with a lone pair of electrons on a nearby atom of F,O, or N.
The normal boiling point for water is 100 degrees Celsius. (The graph below has some artistic problems locating 100.) The observed bp is high compared to the expected value. The predicted bp from the trend of boiling points for H2Te, H2Se, H2S and H2O is very low. If the trend continued the predicted boiling point would be below -62 oC. The "anomalous" boiling point for water is the result of hydrogen bonding between water molecules.
methyl ether, CH3OCH3
Hydrogen peroxide, H2O2
methyl alcohol, CH3OH
Which of the following molecules display hydrogen bonding?
Answer: The hydrogen peroxide and methyl alcohol have hydrogen bonding between molecules.
Possible combinations where hydrogen bond can exist. The first entry shows the covalent bond to the O or N atom. These atoms form two and three covalent bonds. The single covalent bond between O,N,F is shown and the dashed line shows the hydrogen bond. NOTICE the H atom is attracted to a lone pair on the nearby N,O,F atom.
A covalent bond between -O-H ---- :O-
A covalent bond between -N-H----- :O-
A covalent bond between F-H ------ :O-
A covalent bond between -O-H ---- :N-
A covalent bond between -N-H---- :N-
A covalent bond between F-H ----- :N-
A covalent bond between -O-H ----- :F-
A covalent bond between -N-H ---- :F-
A covalent bond between F-H ------ :F-
Summary on hydrogen bonding
Hydrogen bonding is responsible for the expansion of water when it freezes. The water molecules in the solid have tetrahedral spatial arrangement for the two lone pairs and two single bonds radiating out from the oxygen. The lone pairs on the "O" atoms are attracted to nearby water molecules through hydrogen bonds. A cage like structure results. The cage has an hexagon shaped opening.
Melting Points and Boiling Points of Substances with Similar Formula Weights
All the substances in this table have similar formula weights thus they have similar London forces. If the only attractions between substances have to do with size, then they should have similar melting points and boiling points. They do not. Let us look more closely at the nature of the substance to see if we can relate the structure of the material with its properties.
Fluorine and nitrogen monoxide are similar in size and thus have similar London forces. Fluorine is a non-polar covalent molecule while nitrogen monoxide is a polar covalent molecule - it has a positive and a negative end, like a magnet. Since nitrogen monoxide has the higher melting point and boiling point, it must have the stronger intermolecular forces. Given the same size, polar covalent molecules must have stronger forces of attraction than non-polar covalent molecules. These forces of attractions are called dipole-dipole forces.
Nitrogen monoxide and methanol are similar in size and thus have similar London forces. Nitrogen monoxide and methanol are polar covalent molecules and thus have dipole-dipole forces. Since methanol has the higher melting point and boiling point, it must have the stronger intermolecular forces. The difference in these molecules is the presence of a certain extremely polar bond present in methanol that is not present in nitrogen monoxide. This is the oxygen - hydrogen bond.
Oxygen is more electronegative than hydrogen and pulls the electron density in the oxygen - hydrogen bond towards it. This leaves very little electron density around the hydrogen since hydrogen has no core electrons. The part of hydrogen directed away from the oxygen - hydrogen bond has very little electron density shielding the nucleus. Thus that part of the hydrogen nucleus which is exposed can interact with the non-bonding electrons on another methanol molecule. This interaction of a non-bonding pair with a hydrogen attached to an electronegative element such as oxygen is called a hydrogen bond.