8–1

Chapter 13

States of Matter; Liquids and Solids

8–2

Overview

• Changes of State– Phase transitions– Phase Diagrams

• Liquid State– Properties of Liquids; Surface tension and viscosity– Intermolecular forces; explaining liquid properties

• Solid State– Classification of Solids by Type of Attraction between Units– Crystalline solids; crystal lattices and unit cells– Structures of some crystalline solids– Calculations Involving Unit-Cell Dimensions– Determining the Crystal Structure by X-ray Diffraction

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Phase Transitions

• Melting: change of a solid to a liquid.• Freezing: change a liquid to a solid.

• Vaporization: change of a solid or liquid to a gas. Change of solid to vapor often called sublimation.

• Condensation: change of a gas to a liquid or solid. Change of a gas to a solid often called deposition.

H2O(s) H2O(l)

H2O(l) H2O(s)

H2O(l) H2O(g) or

H2O(s) H2O(g)

H2O(g) H2O(l) orH2O(g) H2O(s)

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Comparison of Gases, Liquids and Solids

– Gases are compressible fluids. Their molecules are widely separated.

– Liquids are relatively incompressible fluids. Their molecules are more tightly packed.

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

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Energy of Heat and Phase Change

• Heat of vaporization: heat needed for the vaporization of a liquid.

H2O(l) H2O(g) H = 40.7 kJ

• Heat of fusion: heat needed for the melting of a solid.

H2O(s) H2O(l) H = 6.01 kJ

• Temperature does not change during the change from one phase to another.

8–6

Start with a solution consisting of 50.0 g of H2O(s) and 50.0 g of H2O(l) at 0°C. Determine the heat required to heat this mixture to 100.0°C and evaporate half of the water. The Heat of Fusion of water is 333 j/g and its specific heat is 4.184 j/g oC.

8–7

Intermolecular Forces

• Intermolecular forces: attractions and repulsions between molecules that hold them together.

• Intermolecular forces (van der Waals forces) hold molecules together in liquid and solid phases.

• The strength of the intermolecular forces determines the physical properties of the substance.

• All forces are electrostatic in nature

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Ion – Dipole Force

Interaction between an ion and partial charges in a polar molecule.

Ion-Dipole forces are important in ionic substances in polar solvents and in hydrates.

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Dipole-dipole force

Attractive force between polar molecules with positive end of one molecule aligned with negative side of other.

Polar molecules attract one another when the partial positive charge on one molecule is near the partial negative charge on the other molecule.

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Hydrogen bonds (H-FON Bonds)

The interaction between hydrogen bound to an electronegative element (N, O, or F) and an electron pair from another electronegative element.

The hydrogen atom has no inner core of electrons, so the side of the atom facing away from the bond represents a virtually naked nucleus.

This positive charge is attracted to the negative charge of an electronegative atom in a nearby molecule.

Hydrogen bonding is the dominate force holding the two DNA molecules together to form the double helix configuration of DNA.

8–11

London Dispersion Forces

The motion of electrons within an atom or non-polar molecule can result in a transient dipole moment

The distribution of electrons around an individual atom, at a given instant in time, may not be perfectly symmetrical. Both electrons may be on one side of the nucleus.

The atom would have an apparent dipole moment at that instant in time (a transient dipole moment)

Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom.

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Dispersion Forces cont.

• Dispersion forces are also present between polar/nonpolar and polar/polar molecules (i.e. between all molecules)

8–13

• http://chemmovies.unl.edu/ChemAnime/attractive_forces.htm

8–14

Polarizability

Ease with which the electron cloud of some substance can be distorted by presence of some electric field (such as another dipolar substance).

Related to size of atom or molecule. Larger molecules tend to have greater polarizability because their electrons are further away from the nucleus and the number of electrons is greater.

Dispersion forces tend to increase with increasing molar mass.

8–15

Comparison of Energies for Intermolecular Forces

Interaction Forces Approximate Energy

Intermolecular

London 1 – 10 kJ

Dipole-dipole 3 – 4 kJ

Ion-dipole 5 – 50 kJ

Hydrogen bonding 10– 40 kJ

Chemical bonding

Ionic 100 – 1000 kJ

Covalent 100 – 1000 kJ

8–16

Determination of Intermolecular Forces

• Draw the Lewis Structure• Determine the molecular geometry• Determine if the molecule is polar or nonpolar

• Nonpolar molecule = dispersion forces• Polar molecule = dipole forces, dispersion forces• Polar molecules with hydrogen directly bonded to

fluorine, oxygen or nitrogen = hydrogen bonding, dipole forces, dispersion forces

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8–18

Properties of Liquids

• Vapor Pressure• Boiling Point• Surface Tension• Viscosity

8–19

Vapor Pressure

• In a sealed container, some of a liquid evaporates to establish a pressure in the vapor phase.

• Vapor pressure: partial pressure of the vapor over the liquid measured at equilibrium and at some temperature.

• Dynamic equilibrium

8–20

• http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/flash.mhtml

8–21

Temperature Dependence of Vapor Pressures

• The vapor pressure above the liquid varies exponentially with changes in the temperature.

• The difference between the various liquids is due to the strength of the intermolecular forces

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Clausius – Clapeyron Equation

• A straight line plot results when ln P vs. 1/T is plotted and has a slope of Hvap/R.

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Clausius-Clapeyron Equation

• The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as:

8–24

• Clausius – Clapeyron equation is true for any two pairs of points.

• Write the equation for each:

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Clausius – Clapeyron Equation

• Combine to get:

8–26

Using the Clausius – Clapeyron Equation

• Boiling point the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere.

• Normal boiling point the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm).

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• Determine normal boiling point of chloroform if its heat of vaporization is 31.4 kJ/mol and it has a vapor pressure of 190.0 mmHg at 25.0°C.

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• The normal boiling point of benzene is 80.1°C; at 26.1°C it has a vapor pressure of 100.0 mmHg. What is the heat of vaporization?

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Boiling Points vs. Molecular Weight

8–30

Properties of Liquids

• Surface tension: the energy required to increase the surface area of a liquid by a unit amount.

• Surface tension: The net pull toward the interior of the liquid makes the surface tend to as small a surface area as possible and a substance does not penetrate it easily.

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• Viscosity: a measure of a liquid’s resistance to flow.

• Viscosity is related to mobility of a molecule (proportional to the size and types of interactions in the liquid).

• Viscosity decreases as the temperature increases since increased temperatures tend to cause increased mobility of the molecule.

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Structure of Solids

• Types of solids:– Crystalline – a well defined arrangement of atoms; this

arrangement is often seen on a macroscopic level.• Ionic solids – ionic bonds hold the solids in a regular

three dimensional arrangement.• Molecular solid – solids like ice that are held together

by intermolecular forces.• Covalent network – a solid consists of atoms held

together in large networks or chains by covalent networks.

• Metallic – similar to covalent network except with metals. Provides high conductivity.

– Amorphous – atoms are randomly arranged. No order exists in the solid.

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Unit Cells in Crystalline Solids

• Metal crystals made up of atoms in regular arrays – the smallest of repeating array of atoms is called the unit cell.

• There are 14 different unit cells that are observed which vary in terms of the angles between atoms some are 90°, but others are not.

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• Length of sides a, b, and c as well as angles vary to give most of the unit cells.

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• http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/flash.mhtml

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Cubic Unit Cells in Crystalline Solids

• Simple-cubic shared atoms are located only at each of the corners. 1 atom per unit cell.

• Body-centered cubic 1 atom in center and the corner atoms give a net of 2 atoms per unit cell.

• Face-centered cubic corner atoms plus half-atoms in each face give 4 atoms per unit cell.

8–37

Calculations involving the Unit Cell

• The density of a metal can be calculated if we know the length of the side of a unit cell.

• The radius of an metal atom can be determined if the unit cell type and the density of the metal known

8–38

– Relationship between length of side and radius of atom:

• Simple 2r = l;

• FCC:

• BCC l4

3r

l42

r

8–39

Polonium crystallizes according to the primitive cubic structure. Determine its density if the atomic radius is 167 pm.

8–40

Calculate the radius of potassium if its density is 0.8560 g/cm3 and it has a BCC crystal structure.

8–41

Phase Diagrams

• Graph of pressure-temperature relationship; describes when 1,2,3 or more phases are present and/or in equilibrium with each other.

• Lines indicate equilibrium state two phases.

• Triple point- Temp. and press. where all three phases co-exist in equilibrium.

• Critical temp.- Temp. where substance must always be gas, no matter what pressure.

• Critical pressure- vapor pressure at critical temp.• Critical point- point where system is at its critical pressure and

temp.