Chapter 11 States of Matter; Liquids and Solids Dr. Peter Warburton [email protected]

Chapter 11States of Matter; Liquids and Solids

Dr. Peter [email protected]://www.chem.mun.ca/zcourses/1011.php

© Peter Warburton 2008All media copyright of their respective

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Gases

Gases are compressible fluids.

This behaviour arises because the gas molecules are in

constant random motion through mostly empty space.


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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 11.2


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Ideal gas law

We’ve treated gases in the past as behaving ideally, so that

PV = nRT

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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All molecules have intermolecular forces (IMFs)

Attractive IMFs cannot be overcome at sufficiently low temperatures (molecules

have low average kinetic energies).

Lower T!


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Changes in state

We can often change the physical state of a substance (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

phase.Since ALL of the the energy is involved in the phase change, the temperature MUST REMAIN 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 each molecule, on average, has more

kinetic energy that COULD ALLOW the molecule 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, then 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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Boiling point

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.


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Normal boiling point

The normal boiling point of a liquid is the temperature where

the vapour pressure IS THE SAME AS an external pressure of

exactly1 ATMOSPHERE.

1 atm = 760 mmHg = 101.325 kPa


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Freezing (or melting) point

The freezing point of a liquid is the temperature where 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.


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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.


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Intermolecular forces

The 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.


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Intermolecular forces

At everyday conditions, the forces between molecules tend to be

weakly attractive overall.

These intermolecular forces are generally called van der Waals

forces. However, vdW forces can be subdivided into two different groups.


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Dipole-dipole forces

We’ve already seen that some molecules have a permanent

molecular dipole.

Since full charges are not involved in molecular dipoles, these dipole-dipole

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


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Dipole-dipole forces

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

:

Cl-H


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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.”


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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.


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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.


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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.

Larger molecules and atoms are more polarizable, and generally

have larger London forces.


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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.

Stronger London forces

Weaker London forces


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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.


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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


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Hydrogen bonding (Figure 11.18)

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.


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Hydrogen bonding (Figure 11.18)

This trend is generally true, except for NH3, H2O, and HF because of the hydrogen bonds

(stronger than London forces) that can occur.


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IMFs and bonding at a glance


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Problem 11.37

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

BF3 CH3CHOHCH3 HI


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Problem 11.43


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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.


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Molecular 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.


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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.


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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.


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Diamond and graphite


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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 solids are conducting of

electricity and heat. Silver and iron are examples.


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Documents

Chapter 11 States of Matter; Liquids and Solids Dr. Peter Warburton [email protected]