Chapter 11: Liquids, Solids, and Intermolecular Forces

Chapter 11: Liquids, Solids, and Intermolecular Forces


11.1 Climbing Geckos and Intermolecular Forces (Suggested Reading) 11.2 Solids, Liquids and Gases: A Molecular Comparison [11.1]

11.3 Intermolecular Forces: The Forces that Hold Condensed Phases Together [11.4]

11.5 Vaporization and Vapor Pressure (Excluding the Clausius-Clapeyron Equation) [11.2]

11.6 Sublimation and Fusion [11.2]

11.7 Heating Curve for Water [11.2]

11.11 Crystalline Solids: The Fundamental Types [11.5]States of Matter; Liquids and Solids, Paul G. Mezey

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States of Matter; Liquids and Solids, Paul G. Mezey

.Gases are compressible fluids.

This behaviour arises because the gas molecules are in constant random motion through mostly empty space.

Gases



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Liquids

Liquids are incompressible fluids.

This behaviour arises because the molecules of the liquid are in

constant random motion

without much empty space to move around in.



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Solids

Solids are incompressible and rigid.

This behaviour arises because the molecules of the solid can only vibrate because they have almost no empty space to move into.



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Figure



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Ideal gas lawWe might already know from earlier studies that gases behaving ideally obey the law

PV = nRT

where P, V, n, R, and T stand for pressure, volume, number of moles, gas constant, and absolute temperature, respectively.

However, this gas law ASSUMES that molecules:

1) Have NO SIZE (no repulsive intermolecular forces)

2) DO NOT have attractive intermolecular forces



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All molecules have intermolecular forces (IMFs)

Repulsive IMFs (real molecules have size) will limit the compressibility of a

group of molecules.

SQUEEZE!



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Changes in state

We can often change the physical state

of a substance (in a process called a phase transition) by

changing the temperature and/or pressure

by suitable amounts.



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Phase transitions:

From solid to liquid is fusion.

Change in enthalpy (or heat) of fusion is Hfus

From liquid to gas is vaporization.

Change in enthalpy of vaporization is Hvap

From solid to gas is sublimation.

Change in enthalpy of sublimation is Hsub



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Phase transitions



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Enthalpy of phase transitions

The enthalpy change of a phase transition tells us how much heat must be added to (or removed from) a substance in a given phase so it changes its phase.

Since ALL of the the energy is involved in the phase change, the temperature REMAINS CONSTANT during the change.



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Enthalpy of phase transitions



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Vapour pressure

Vapour pressure is the partial pressure

(the part of the total pressure that comes from a given substance)

of the vapour (gas)

above the liquid (or solid) phase

measured

At EQUILIBRIUM

at a GIVEN TEMPERATURE



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Equilibrium and vapour pressure

Vapour pressure should be measured when the vapour pressure

has STOPPED CHANGING.

This equilibrium means that the rate of molecules leaving the liquid (or solid) phase is BALANCED EXACTLY by the rate of the molecules in the vapour joining the liquid (or solid) phase.



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Equilibrium and vapour pressure



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Temperature and vapour pressure

The vapour pressure depends on how easily molecules can overcome attractive intermolecular forces that keep it in the liquid (or solid) phase.

Higher temperatures mean molecules, on average, have more kinetic energy that COULD ALLOW the molecules to “escape” from the attractive forces.



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.Since all groups of molecules AT THE SAME TEMPERATURE have the SAME AVERAGE KINETIC ENERGY, molecules that have WEAKER intermolecular forces are MORE VOLATILE (have GREATER vapour pressures at the GIVEN temperature) than molecules with STRONGER intermolecular forces.




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In this Figure, the MOST VOLATILE (weakest IMFs) liquid is on the left,

while the LEAST VOLATILE (strongest IMFs) is on the right.




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The boiling point of a liquid is the temperature where the vapour pressure IS THE SAME AS the external pressure.

Therefore, if the external pressure changes, the boiling point temperature ALSO changes.

Boiling Point



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The normal boiling point of a liquid is the temperature where the vapour pressure IS THE SAME AS the “normal” external pressure: exactly 1 ATMOSPHERE.

1 atm = 760 mmHg = 101.325 kPa

Normal Boiling Point



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The freezing point of a liquid is the temperature at which the phase transition from liquid to solid occurs.

Since melting (fusion) is the exact opposite transition (solid to liquid), the freezing point and melting point are IDENTICAL.

Freezing (or melting) point



. Freezing and boiling points

Since the normal freezing and normal boiling points of a PURE substance are fixed properties, measuring them is an easy and useful first step in identifying unknown compounds.



. Intermolecular forces

The primary forces between molecules are electrostatic. They depend on charges (like charges repel, opposite charges attract), and the distance between the charges.

Larger charges (like those found on ions) and smaller distances between molecules tend to lead to stronger forces.



. Intermolecular forces

At ordinary conditions, the forces between molecules tend to be weakly attractive overall.

These intermolecular forces are generally called van der Waals forces.

These vdW forces can be subdivided into two groups.



. Dipole-dipole forces

We’ve already seen that some molecules have a permanent molecular dipole.

Since full charges are not involved in molecular dipoles, the dipole-dipole intermolecular interactions are

relatively weak as compared to ionic bonds, where full charges are involved.



. Dipole-dipole forces

The larger the dipole moment (the molecules are more polar), the stronger the IMFs tend to be.

:

Cl-H



. London (dispersion) forces

Nonpolar molecules still interact with each other despite their lack of permanent dipoles.

In a nonpolar molecule, “on average,” the electrons do not prefer one part of the molecule over the other. However, at any given instant, they might not be evenly distributed and so the molecule ends up having a “temporary dipole.”




When a molecule with a temporary dipole comes close to another molecule, the electrons of the second molecule will try to move away from the negative partial charge of the first molecule, leading to the second molecule having a temporary induced dipole.




London forces tend to be very weak because the partial charges tend to be small and fleeting.

However, ALL chemical species have London forces between them.

Variations in the strength of London forces depend on two factors.




Polarizability – is the ability for electrons to move freely within the molecule. The more freely electrons can move, the larger the induced dipole can be.




Shape – The shape of a molecule plays a part in determining how the electrons can move in a molecule.

More compact shapes are usually more symmetrical and allow less contact between molecules. They generally have smaller induced dipoles with weaker London forces.



. Hydrogen bonding

A hydrogen bond is an attractive interaction between a hydrogen atom bonded to a very electronegative atom (O, N, and F), and an unshared electron pair on another electronegative atom.



. Hydrogen bonding

Hydrogen bonds are really just a very special case of dipole-dipole forces.

H-F, H-O, and H-N bonds are very polar (larger partial charges)

Also, because the hydrogen is very small, it is possible for another molecule to approach it very closely (short distance).

Hydrogen bonds are relatively strong intermolecular forces



.Hydrogen bonding

With London forces, boiling points will increase with molecular size (polarizability). We EXPECT boiling points to follow the trend

CH4 < SiH4 < GeH4 < SnH4

and so on across the periodic table.



.Hydrogen bonding

This trend is generally true, except for NH3, H2O, and HF because of the hydrogen bonds (stronger than London forces) that can occur.



. IMFs and bonding at a glance



. Problem

For each of the following substances, list the kinds on intermolecular forces expected:

BF3 CH3CHOHCH3 HI



. Problem



. Classification of solids

Ionic solids are formed by regular arrangements of cations (positive) and anions (negative). Because of the strong ionic bonds involved, the melting points are usually high, and the solids are brittle and hard.

NaCl is an example.



. Classification of solidsMolecular solids are separate molecules held together through the mainly weaker dispersion, dipole-dipole and hydrogen bond intermolecular forces.

Weaker forces mean lower melting points and softer consistency.

Also, since electrons can’t easily move from one molecule to another, the solids are nonconducting of electricity.

Ice and sugar are examples.




Covalent network solids are formed by a large array of atoms. These solids are formed from repeating units, but are almost better considered to be “one large molecule”.

If the arrays are three-dimensional, the solids are rigid, and so the melting points are usually high, and the solids are hard but usually not brittle.

Quartz and diamond are examples.




Covalent network solids can also have structures of loosely held rigid sheets or chains.

Graphite can be used as a lubricant because the sheets of carbon can easily slide past each other.

Asbestos forms chains.



. Diamond and graphite




Metallic solids are large arrays of metal nuclei found in an “electron sea”.

Since the forces vary widely amongst metals we see widely variable melting points and hardnesses.

However, since electrons can move from one atom to another, the metallic solids are good conductors of electricity and heat.

Silver and iron are examples.



. Classification of solidsClassification of solids



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Chapter 11: Liquids, Solids, and Intermolecular Forces

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Transcript of Chapter 11: Liquids, Solids, and Intermolecular Forces