states of matter: liquids and solids chapter 14. chapter 112 copyright © by houghton mifflin...

States of Matter: Liquids and

Solids

Chapter 14

Chapter 11 2Copyright © by Houghton Mifflin Company. All rights reserved.

States of Matter

Comparison of gases, liquids, and solids.

Gases are compressible fluids, completely fill a container, and molecules are widely separated (low density).

Liquids are relatively incompressible fluids, molecules are more tightly packed.

Solids are nearly incompressible and rigid. Their molecules or ions are in close contact and do not move.


Changes of State

A change of state or phase transition is a change of a substance from one state to another.

solid

liquid

gas

melting freezing

condensationboiling

sublimation(see Figure 11.3)

condensation or

deposition


Properties of Liquids; Surface Tension and Viscosity

The molecular structure of a substance defines the intermolecular forces (forces that occur between molecules) holding it together. Intramolecular = bonds

Many physical properties of substances are attributed to their intermolecular forces.

These properties include vapor pressure and boiling point.



Surface tension

A molecule within a liquid is pulled in all directions, whereas a molecule on the surface is only pulled to the interior.

As a result, there is a tendency for the surface area of the liquid to be minimized.



Surface tensionThis explains why falling raindrops are nearly spherical, minimizing surface area.

In comparisons of substances, as intermolecular forces between molecules increase, the apparent surface tension also increases.


14.1 Intermolecular Forces; Explaining Liquid Properties

Viscosity is the resistance to flow exhibited by all liquids and gases.

Viscosity can be illustrated by measuring the time required for a steel ball to fall through a column of the liquid. Even without such measurements, you know that syrup has a greater viscosity than water.

In comparisons of substances, as intermolecular forces increase, viscosity usually increases.


Figure 11.20:Comparison of the viscosities of two liquids. Photo courtesy of James Scherer.

Return to Slide 30


Intermolecular Forces; Explaining Liquid Properties

Many of the physical properties of liquids (and certain solids) can be explained in terms of intermolecular forces, the forces of attraction between molecules.

Three types of forces are known to exist between neutral molecules (ionic – not neutral).

1. London (or dispersion) forces

2. Dipole-dipole forces

3. Hydrogen bonding


London Forces

London Dispersion forces are the weak attractive forces resulting from instantaneous dipoles that occur due to the distortion of the electron cloud surrounding a molecule.

London forces increase with molecular weight. The larger a molecule, the more easily it can be distorted to give an instantaneous dipole. (More e-’s available, stronger bond.)

All covalent molecules exhibit some London force.


Dipole-Dipole Forces

Polar covalent molecules (dipole moment) can attract one another through dipole-dipole forces.

The dipole-dipole force is an attractive intermolecular force resulting from the tendency of polar molecules to align themselves positive end to negative end. As distance between molecules increases, dipole-dipole force decreases.

H Cl H Cl


Hydrogen bonding occurs in substances containing hydrogen atoms bonded to certain very electronegative atoms. Strong dipole-dipole forces occur between H and a strong electronegative atom (N, O, F).

Does H2S have hydrogen bonding?

No

H2O?

Yes

HYDROGEN BONDING


Identify the major attracting force.

HF

Hydrogen bonding

N2

London dispersion

PCl3

Dipole-dipole


14.2 Water and its phase changes

Water is the most important substance.

Crucial for sustaining the reactions in our bodies.

Oceans moderate Earth’s temperatures

Water cools automobile engines and power plants.


Phase Diagrams – Heating/cooling curve

A phase diagram is a graphical way to summarize the conditions under which the different states of a substance are stable. Phase changes are physical not chemical.

Water (at 1 atm) is a liquid between 0o C and 100oC.

Normal boiling point of water is 100oC.

Normal freezing point of water is 0oC.


Figure 11.9: Heating curve for water.

Return to Slide 9


14.3 Energy requirements for the changes of states.

Molar heat of fusion = energy required to melt 1 mole of a substance. (ice = 6.02 KJ/mol)

Molar heat of vaporization = energy required to change 1 mole of a liquid to its vapor. (water = 40.6 KJ/mol)


A Problem to Consider

The heat of vaporization of ammonia is 23.4 kJ/mol. How much heat is required to vaporize 1.00 kg of ammonia?

First, we must determine the number of moles of ammonia in 1.00 kg (1000 g).

33

33

3 NH mol 8.58NH g 0.17NH mol 1

NH g 10 .001


Then we can determine the heat required for vaporization.

kJ 10 1.38 kJ/mol 23.4 NH mol 8.58 33


Problem

Calculate the energy required to vaporize 35.0 g of water at 100°C. The molar hear ot vaporization is 40.6 kJ/mol.

78.9 kJ


Calculating energy changesQ = s x m x ΔT

Q = energy required, s = specific heat capacity, m = mass, ΔT = change in temp.Calculate the energy required to heat 22.5 g of liquid water at 0 °C and change it to steam at 100 °C. The specific heat capacity of liquid water is 4.18 J/g°C. The molar hear of vaporization of water is 40.6 kJ/mol.60.1 kJ


14.4 Evaporation and vapor pressure

The vapor pressure of a liquid depends on intermolecular forces. When the intermolecular forces in a liquid are strong, you expect the vapor pressure to be low.

As intermolecular forces increase, vapor pressures decrease.


Vapor Pressure

Liquids are continuously vaporizing.

Evaporation = a cooling process (endothermic)

Condensation = vapor molecules forming a liquid

Equilibrium = equal evaporation and condensation.

The vapor pressure of a liquid is the partial pressure of the vapor over the liquid, measured at equilibrium at a given temperature.


Vapor Pressure

The vapor pressure of a liquid depends on its temperature.

As the temperature increases, the kinetic energy of the molecular motion becomes greater, and vapor pressure increases.

Liquids and solids with relatively high vapor pressures at normal temperatures are said to be volatile.


Figure 11.7: Variation of vapor pressure with temperature.

Return to Slide 5


14.5 Boiling point and vapor pressure

The normal boiling point is related to vapor pressure and is lowest for liquids with the weakest intermolecular forces.

When intermolecular forces are weak, little energy is required to overcome them. Consequently, boiling points are low for such compounds.


Boiling Point

The temperature at which the vapor pressure of a liquid equals the pressure exerted on the liquid is called the boiling point.

As the temperature of a liquid increases, the vapor pressure increases until it reaches atmospheric pressure.

At this point, stable bubbles of vapor form within the liquid. This is called boiling.

The normal boiling point is the boiling point at 1 atm.


The higher the boiling point, the stronger the intermolecular forces.


Which has the lowest boiling point?


Freezing Point

The temperature at which a pure liquid changes to a crystalline solid, or freezes, is called the freezing point.

The melting point is identical to the freezing point and is defined as the temperature at which a solid becomes a liquid.

Unlike boiling points, melting points are affected significantly by only large pressure changes.

states of matter: liquids and solids chapter 14. chapter 112 copyright © by houghton mifflin...

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