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Flashcards in intermolecular Deck (24):
1

Define intermolecular force.

An intermolecular force is the electrostatic attraction between molecules. These forces are the key determinants of physical properties such as boiling point, vapour pressure, viscosity etc.

2

List the three intermolecular forces.

The three intermolecular forces are London dispersion forces, dipole-dipole forces, and hydrogen bonding.

3

Describe the London dispersion forces.

London dispersion forces are the attraction between molecules as a result of temporary dipoles.

The strength of such forces is related to how polarizable the electrons are, which in turn is related to how big the molecule is. We usually use molar mass as an indication of molecular sze. Therefore, we say London dispersion forces are related to a molecule's molar mass. The greater the molar mass, the stronger the London dispersion force.

4

Describe diplole-dipole forces.

Dipole-dipole forces are the attraction between molecules as a result of permanent dipoles.

That is, they are the attraction between polar molecules (or the attraction between the polar parts of the molecule). The more polar the molecule is, the stronger the dipole-dipole forces. Molecular polarity is determined by considering bond polarity and how these bonds are oriented as related to the molecular geometry.

5

Describe hydrogen bonding

Hydrogen bonding refers to forces of attraction between the nucleus of a hydrogen atom (which is covalently bonded to a nitrogen, oxygen, or fluorine) and a lone pair of electrons on a nitrogen, oxygen, or fluorine. Hydrogen bonding is an intermolecular force (weaker than a typical chemical bond) and can be considered as a very specific kind of dipole-dipole force.

Because the hydrogen is covalently bonded to the highly electronegative nitrogen, oxygen, or fluorine, the hydrogen has large delta positive charge.  As such, it is strongly attracted to the lone pair of electrons of nitrogen, oxygen, or fluorine on a different molecule.

6

Which intermolecular force is common to all molecules?

London dispersion forces exist for all molecules. Some molecules (e.g. H2) only experience London dispersion forces, whereas others (e.g. H2O), in addition to London dispersion forces, have also dipole-dipole forces and hydrogen bonding.

7

Is hydrogen bonding always the most important (i.e. strongest) intermolecular force for a molecule?

No, while hydrogen bonding is considered to be a "strong" intermolecular force, it may not always be the most important.

For a small molecule, like H2O, London dispersion forces are considered to be relatively weak, but the hydrogen bonding is quite significant.

Whereas for a large molecule like octan-1-ol, (CH3CH2CH2CH2CH2CH2CH2CH2OH), while hydrogen bonding (and dipole-dipole forces) are present, the most important intermolecular force would be London dispersion forces.

8

Define vapour pressure.

Vapour pressure is the pressure of a vapour in equilibrium with its liquid phase (at a particular temperature).

9

Define boiling point.

Boiling point is the temperature at which the vapour pressure is equal to the atmospheric pressure. Therefore a boiling point should specify both a temperature and an atmospheric pressure.

Normal boiling point is the boiling point when atmospheric pressure equals exactly 1 atm. At the precise boiling point both the liquid and gas phases of a chemical are in equilibrium.

10

Why is the boiling point related to the strength of intermolecular forces?

Although boiling point is defined as the temperature at which the vapour pressure is equal to atmospheric pressure, the boiling point also indicates at which temperature molecules have sufficient energy to break from the forces of attraction between molecules.

The stronger the intermolecular forces, the more energy that is required to break these forces of attraction, and the higher will be the boiling point.

11

Why is the vapour pressure related to the strength of intermolecular forces?

Vapour pressure is the pressure of a vapour in equilibrium with its liquid phase (at a particular temperature).

The stronger the intermolecular forces, the more energy required to break these forces of attraction and (at a particular temperature) the fewer molecules that will be in the gas phase. Fewer molecules in the gas phase indicates a lower vapour pressure.

12

Explain an important outcome of the fact that boiling point depends on atmospheric pressure.

An important outcome of the fact that boiling point depends on atmospheric pressure is that the boiling point of chemical can be changed.

For instance, by using a pressure cooker, we can artificially increase the atmospheric pressure inside the cooker and has a result have water boil at a temperature much higher than 100 °C. This allows food to cook more quickly. Alternatively, under conditions of low atmospheric pressure (at high elevations for example), water boils at temperature less than 100 °C.

13

Explain what information is available from a phase diagram.

A phase diagram for a chemical is a plot of atmospheric pressure versus temperature.

This plot shows under which conditions of pressure and temperature the chemical is in the solid, liquid and gas phase. Further, it shows under which conditions the these phases are in equilibrium. Therefore a trace at 1 atmospheric pressure shows the normal melting point and the normal boiling point.

The slope of the line between the solid and liquid phases also indicates which phase is more dense. For water, this line has a negative slope because the liquid phase for water is more dense. For most other chemicals, this line has a positive slope, which indicstes the solid phase is more dense.

14

Define triple point.

The triple point in a phase diagram represents the conditions of pressure and temperature where all three phases of matter (solid, liquid and gas) are in equilibrium - that is, they coexist.

15

Define sublimation.

Sublimation is the phase transition from solid to gas.

16

Define critical point.

The critical point (for a particular chemical) indicates the temperature above which the gaseous phase cannot condensed to a liquid through the application of pressure. Under such conditions (i.e. high pressure above the critical point) the chemical has both gas and liquid-like properties and is referred to as a "super-critical fluid".

17

Explain from a molecular viewpoint why the solid phase for water is less dense than its liquid phase.

The solid phase for water is less dense than the liquid phase because in the solid phase the water molecules are fixed in space (in the liquid phase the water molecules flow over one another). The most energetically favourable arrangement for the water molecules in the solid phase is for the water molecules to align in a way to maximize hydrogen bonding.

Such an alignment puts the water molecules in a more open crystal lattice (as compared to the liquid phase). For a given mass, this crystal lattice occupies a greater volume and therefore the solid phase is less dense.

18

Explain an important outcome of the fact that the solid phase for water is less dense than the liquid phase.

An important outcome of the fact that solid water is less dense is that when liquid water freezes, it occupies a greater volume. This is important in that we usually need to drain liquid water from pipes when there is a risk that the liquid water will freeze. Such freezing can cause the pipes to break, which will result in leaks and costly repairs.

19

These represent the electrostatic attraction between molecules. These forces are the key determinants of physical properties such as boiling point, vapour pressure, viscosity etc.

Intermolecular force

20

The pressure of a vapour in equilibrium with its liquid phase (at a particular temperature).

Vapour pressure

21

The temperature at which the vapour pressure is equal to the atmospheric pressure. Therefore this point should specify both a temperature and an atmospheric pressure.

 

Boiling point

22

In a phase diagram, this represents the conditions of pressure and temperature where all three phases of matter (solid, liquid and gas) are in equilibrium - that is, they coexist.

Triple point

23

The phase transition from solid to gas.

Sublimation

24

For a particular chemical, this indicates the temperature above which the gaseous phase cannot condensed to a liquid through the application of pressure. Under such conditions (i.e. high pressure above the critical point) the chemical has both gas and liquid-like properties and is referred to as a "super-critical fluid".

Critical point