Inter-molecular forces Flashcards
What are inter-molecular forces? Name their types from weakest to strongest.
They are weak attraction forces that occur between molecules.
- London forces (instantaneous dipole-induced dipole)
- Permanent dipole- permanent dipole interactions
- Hydrogen bonds
In which type of atoms/molecules do london forces occur? Explain the formation of london forces.
They occur in all atoms and molecules: polar or non-polar.
1) Electrons in charge clouds are always moving very quickly. At any particular moment, the electrons in an atom are likely to be more to one side than the other.
At this moment, the atom would have a temporary (or instantaneous) dipole.
2) This dipole can induce another temporary dipole in the opposite direction on a neighbouring atom. The two dipoles are then attracted to each other.
3) The second dipole can induce yet another dipole in a third atom.
4) Because the electrons are constantly moving, the dipoles are being created and destroyed all the time. Even though the dipoles keep changing, the overall effect causes the molecules to be attracted to each other.
What do london forces cause in molecules (what structure does it produce)?
What is the effect of stronger london forces on molecules?
London forces hold molecules in a lattice. For example, I2 molecules are held together in a lattice by london forces. This forms a simple molecular structure.
Stronger london forces mean high melting and boiling points.
What affects the strength of london forces, and in turn the melting and boiling point of the corresponding molecule?
The number of electrons affects the strength of London forces.
- Larger molecules have larger electron clouds, meaning stronger London forces.
- Molecules with greater surface areas also have stronger London forces because they have a bigger exposed electron cloud.
Explain:
1- the trends in boiling temperatures of alkanes with increasing chain length
2- the effect of branching in the carbon chain on the boiling temperatures of
alkanes
1- The longer the carbon chain, the stronger the London forces- this is because there is more points of contact between adjacent molecules where electrons interact and so more london forces exist.
So, as as the molecules get longer, it gets harder to separate them because it takes more energy to overcome the larger no. of london forces.
The higher the mass of the molecule, the higher the number of electrons per molecule and so the instantaneous and induced dipoles also increase, so the more the london forces.
All of these cause high boiling temperatures in molecules.
2- Branched-chain alkanes can’t pack closely together and their molecular surface contact is small compared to straight chain alkanes of similar molecular mass. So fewer London forces can form. (Less points of contact between adjacent molecules)
(see cambridge tb)
In what type of molecules does dipole-dipole bonding occur? Does it occur along with london?
It occurs in polar molecules. Yes, where LDF and dipole dipole forces occur at the same time, there are higher boiling/melting points.
Explain dipole-dipole bonding.
When dipole-dipole weaker than london?
When dipoles are aligned correctly in 2 molecules, there will be a favourable interaction and the 2 molecules attract each other.
However, sometimes for example in a liquid, the random arrangement of molecules does not allow for the alignment of the dipoles to produce a favourable interaction. As a result, the average dipole-dipole strength is weaker than the strength of london forces.
Conditions for the formation of a hydrogen bond?
1- The atom the hydrogen is bonded to must be more electronegative than hydrogen.
2-Hydrogen bonding is most significant when Hydrogen is bonded to the following elements- Oxygen, Fluorine and Nitrogen. This is because they are small and highly electronegative. Hydrogen bonding is not confined to these elements.
L.O: understand the interactions in molecules, such as H2O, liquid NH3 and liquid HF, which give rise to hydrogen bonding
Draw the hydrogen bonding in these molecules.
How do u predict whether hydrogen bonding occurs in molecules other than the ones mentioned above?
CGP for drawings
You predict whether hydrogen bonding occurs in a molecule by checking for OH and NH groups.
Explain:
1- the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons
2- the trends in boiling temperatures of the hydrogen halides HF to HI
1- They have additional hydrogen bonding compared to alkanes with a similar no. of electrons. This increases the energy required to separate the molecules of alcohols, which increases their boiling points/lowers their volatility.
2- In HF, hydrogen bonding occurs, which is the strongest intermolecular force. This allows for it to have a very high boiling temperature. From HCl to HI, there is a steady increase in boiling temperature. This is because the number of electrons per molecule increases, increasing the strength of London forces.
Define enthalpy change of vapourisation and how can we use it to determine the strength of intermolecular forces?
It is a measure of the amount of energy that is required to completely separate the molecules of a liquid and convert it to a gas at the same temperature.
The greater the enthalpy change of vapourisation, the greater the forces of attraction between molecules.
Explain the factors that influence the choice of solvents, including:
1- water, to dissolve some ionic compounds, in terms of the hydration of the ions
2- water, to dissolve simple alcohols, in terms of hydrogen bonding
3- water, as a poor solvent for compounds (to include polar molecules such as halogenoalkane), in terms of inability to form hydrogen bonds
4- non-aqueous solvents, for compounds that have similar intermolecular forces to those in the solvent
1- Water is used to dissolve ionic compounds. This is due to the energy supplied by the hydration of ions to completely or partially separate the ions in the compound.
S- ends of water molecules attract the positive ions and remove them from the lattice. They then become surrounded by water molecules. The interaction between water molecules and ions is called ion-dipole interaction.
S+ ends attract negative ions and then become surrounded by water molecules when in solution. The ion becomes hydrogen bonded to the water molecule.
2- Alcohols contain an OH group so they can form hydrogen bonds with water.
Alcohols become less soluble in water as the carbon chain length increases because london forces slowly become the more predominant IMF.
3- For some molecules, including non-polar and some polar, water is a poor solvent. This is because they either do not form hydrogen bonds with water, or the hydrogen bonds they form are not strong enough to replace the bonds in water molecules. *
4- Non-polar substances such as ethene have London forces between their molecules.
They form similar bonds with non-polar solvents such as hexane — so they tend to dissolve in them.
Water molecules are attracted to each other more strongly than they are to non-polar molecules such as iodine — so non-polar substances don’t tend to dissolve easily in water.
Like dissolves like (usually) — substances usually dissolve best in solvents with similar inter-molecular forces.
Explain the following anomalous properties of water resulting from hydrogen bonding:
1- its high melting and boiling temperature when compared with similar molecules
2- the density of ice compared to that of water
1- Water molecules can form upto 4 hydrogen bonds per molecule. This means it has extensive hydrogen bonding, and this makes the strength of HB in water very high. This means it has a very high boiling and melting point compared to similar molecules with the same no. of electrons or that also have HB. Ex. HF and NH3.
2- The density of ice at 0 degrees celcius is much lower than the density of liquid water at the same temperature.
The water molecules in ice are arranged in rings of six held together by hydrogen bonds. This is a simple molecular structure that creates large areas of open space inside the rings. This makes it have a low density.
When melted, some of these hydrogen bonds are broken and the ring structure is destroyed which leaves a much smaller distance between molecules, increasing the density.
