Topic 7 Intermolecular Forces Flashcards
Describe the nature of London forces.
Due to random electron density fluctuations, the electron density cloud of a molecule can become temporarily asymmetrical.
This creates a temporary dipole in the molecule which induces a dipole in an adjacent that is always aligned with the first dipole.
This alignment ensures that whenever this occurs, the instantaneous dipole and the induced dipole will always be attracted to each other.
The strength of London forces increases as the number of electrons per molecule increases because the fluctuations in electron density increase.
Describe the nature of permanent dipole-permanent dipole interactions.
If the molecules are polar, they will interact with each other.
If their dipoles are aligned, they will attract each other.
If their dipoles are not aligned, they will repel each other.
Describe the nature of hydrogen bonding.
For this bonding to occur, a hydrogen atom must be bonded to highly electronegative element ( N, O, or F).
The highly electronegative element must also contain a lone pair of electrons.
The δ+ of the hydrogen atom on one molecule will be attracted to the lone pair on the highly electronegative element of another molecule, which forms a hydrogen bond.
The more electronegative the element is, the stronger the hydrogen bond will be.
State the bond angle of hydrogen bonds.
180 degrees
State how many hydrogen bonds there are per molecule in water, ammonia, and hydrogen fluoride.
Water forms 4 hydrogen bonds per molecule.
Ammonia and hydrogen fluoride form 2 hydrogen bonds per molecule.
Explain why London forces are usually more significant than permanent dipole-permanent dipole interactions.
The induced and instantaneous dipoles in London forces are always aligned to attract each other.
However, permanent dipoles are not always aligned and repel each other a considerable amount of the time.
Explain why does hydrogen fluoride have a higher boiling point than ammonia.
Both molecules have 10 electrons each so their London forces will be nearly equal.
Also both molecules form two hydrogen bonds per molecule.
However, fluorine is more electronegative than nitrogen, so it forms stronger hydrogen bonds requiring more energy to break.
Explain why water has a higher boiling point than both ammonia and hydrogen fluoride.
All three molecules have 10 electrons each so their London forces will be nearly equal.
The strength of the hydrogen bonds in water is greater than ammonia but less than hydrogen fluoride.
This is because oxygen’s electronegativity is higher than nitrogen but less than fluorine.
However, water forms 4 hydrogen bonds per molecule while hydrogen fluoride and ammonia only form 2.
Therefore, the net intermolecular force in water requires more energy to break.
Describe and explain the anomalous physical properties of water.
- It has a much higher melting and boiling point compared to similar molecules.
This is because water forms hydrogen bonds while most molecules do not. - Ice has a lower density than liquid water.
This is because when water freezes, hydrogen bonds form between the molecules in such a way to create hexagonal rings.
This structure is very open and contains a lot of empty space, decreasing its density.
Describe and explain the trends in the melting and boiling points of alkanes.
The melting and boiling points of alkanes increase with chain length as the number of electrons per molecules increases, so the strength of the london forces increases.
The melting and boiling points of alkanes decrease with more branching as the points of contact to form london forces between molecules decreases.
Explain why alcohols are less volatile that alkanes with a similar number of electrons.
The london forces will be about the same.
However, since alcohols contain an OH functional group, they form hydrogen bonds while alkanes cannot.
The hydrogen bonds are strong and increase the energy required to break the net intermolecular forces.
Describe and explain the trend of boiling temperatures of hydrogen halides HF to HI.
The B.P. decreases from HF to HCl significantly and the it starts increasing.
The decrease is because HF forms strong hydrogen bonds while the other hydrogen halides do not.
The b.p. starts increasing because the number of electrons per molecule increases, increasing the strength of the London forces.
