Intermolecular Forces Exam Qs Flashcards
Water might be expected to have a lower boiling temperature than hydrogen sulfide but it actually has a higher boiling temperature.
Comment on this statement by referring to the intermolecular forces in both these substances.
A detailed description of how the intermolecular forces arise is not required.
(4)
- LOW BP as water has fewer electrons than hydrogen sulfide
- Water has weaker London forces
- HIGH BP occurs as water has hydrogen bonding
- Hydrogen bonding is STRONGER than London forces so MORE energy is needed to break these forces
Nitrogen trichloride, NCl3, has a boiling temperature of 344 K, and nitrogen trifluoride, NF3, has a boiling temperature of 144 K.
Explain this difference in boiling temperatures, by referring to all the intermolecular forces present.
(5 marks)
…………………………………………………………………………………………
- London forces are greater in NCl3
- as NCl3 has more electrons as Cl has more electrons than the F atom
- NF3 has stronger permanent dipole-dipole interactions than NCl3
- as F is more electronegative than Cl
- MORE energy needed to overcome intermolecular forces between NCl3 molecules
Explain the boiling temperatures increase from chlorine to iodine
Cl= -35
Br= 59
I = 184
- number of electrons increase down the group (17 to 53)
- strength of the instantaneous induced dipole dipole forces increase or London forces AND MORE energy is required
The compounds hydrogen fluoride, water and methane, all have simple molecular structures, but they have significantly different boiling temperatures.
Discuss the reasons for the differences in the boiling temperatures of the three compounds, using the data in the table and the Pauling electronegativity values in the Data Booklet.
CH4 = -161.5
H2O= 100
HF= 19.5
ALL HAVE 10 ELECTRONS
Same number of electrons = SAME LONDON FORCES
LARGE electronegativity for HF and H2O than CH4 (small) - include electronegativity values!!
ONLY WEAK London forces in CH4
Hydrogen bonding in HF and H2O (NOT in CH4)
MORE (DOUBLE the number) hydrogen bonds in H2O than HF
MORE energy needed to break stronger intermolecular forces
Explain why ethene has a boiling temperature of −104 °C, whereas ethanol has a boiling temperature of 78 °C. (3)
Ethanol has hydrogen bonding, London forces and dipole-dipole
Ethene ONLY has weak London forces
MORE energy needed to break the stronger intermolecular forces in alcohols
This question is about the solubility of some alcohols.
The table shows the solubility in water of the first six alcohols in a homologous series.
Methanol, Ethanol and propan-1-ol= soluble in all proportions
But an-1-ol= 632
Pentan-1-ol= 22
Hexan-1-ol= 5.9
Explain why methanol and water are ‘soluble in all proportions’.
Methanol hydrogen bonds to water
At least one lone pair on an oxygen (at 180 degrees)
Strength of intermolecular forces between methanol and water are approx the same
Explain why 2,2-dimethylpropane has a much lower boiling temperature than its isomer pentane.
Detailed descriptions of the forces involved are not required.
Branching results in weaker London forces
Due to LESS surface area/points of contact
The intermolecular attractions between halogen molecules are London forces.
(i) Describe how London forces form between halogen molecules.
UNEVEN distribution of ELECTRONS
Results in an instantaneous/ temporary dipole (on first molecule)
INDUCES a second dipole on another molecule
Explain why bromine has a higher boiling temperature than chlorine. (2)
Cl= -34
Br= 59
Bromine has MORE electrons than chlorine
Bromine has stronger London forces
More energy needed to overcome forces between Br molecules
Methanol, CH3OH, is miscible with water in all proportions. Sodium chloride is much less soluble in methanol than in water.
Explain these statements using your knowledge of the interactions between solutes and solvents.
You must use diagrams to illustrate your answers.
(6)
- M2 and M5 for diagrams (hydrogen bonding between methanol and water) and of one ion
M1- hydrogen bonding between water (solvent) and methanol (solute)
M4- hydration of Na+ and Cl-
M6- ionic bonding is strong than the bonding between sodium or chloride ions and methanol
(i) The O–H bond in water is polar because, when compared with the hydrogen atom, the oxygen atom has
A a higher mass number
B a larger atomic radius
C greater electronegativity
D more electrons
C
Explain why hydrogen bonding causes ice to be less dense than liquid water. (2)
more open/ more space between molecules = LESS dense
Due to 3D structure in Ive
Hydrogen bonds are longer than covalent
In many swimming pools, sodium chlorate(I) has replaced chlorine gas as a disinfectant. Sodium chlorate(I) is an ionic compound. It is very soluble in water.
(i) Describe, using diagrams to illustrate your answer, the interactions between each of the ions and the solvent when sodium chlorate(I) dissolves in water.
BOTH Na and Cl having more than one water molecule around each ion
- show polarisation S- (dipole on oxygen) or S+ (hydrogen)
Ethanol is very soluble in water whereas chloroethane is almost insoluble in water.
Explain this observation by comparing the types of intermolecular forces formed by each of these molecules with water.
- ethanol forms hydrogen bonds with water
- chloroethane forms PERMANENT dipole dipole attractions and LONDON forces
Nitrogen forms several hydrides. In addition to ammonia, NH3, it forms hydrazine, N2H4, in which the two nitrogen atoms are covalently bonded together.
Hydrazine is very soluble in water.
Explain, using a labelled diagram and naming the relevant intermolecular interactions, why
hydrazine is very soluble in water. (3)
- hydrogen bonding between water and hydrazine (1)
- diagram showing correct hydrogen bond between correct atoms ( e.g.N to H)
- lone pair on either nitrogen or oxygen and a bond angle of 180 (1)
