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Slide 1 Solon City Schools •Define polar nonpolar dipole dipole forces ion dipole forces

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Slide 1 Solon City Schools •Define polar nonpolar dipole dipole forces ion dipole forces Powered By Docstoc
					•Define: polar, nonpolar, dipole-dipole forces, ion-dipole forces, Hydrogen “bonding”, and London dispersion forces; sublimation, condensation, evaporation, freezing point, boiling point, vapor pressure, viscosity, surface tension, delta H of fusion, vaporization , and sublimation. •Distinguish between intermolecular and intramolecular attractions •Put a list of compounds in order of increasing melting point, boiling point, and vapor pressure •Use and label the parts of a phase diagram •Understand the expansion of ice due to hydrogen bonding •Use the Clausius-Clapeyron equation

Solid, Liquid, or Gas
Two factors determine whether a substance is a solid, a liquid, or a gas:
1. The kinetic energies of the particles (atoms, molecules, or ions) that make up a substance. Kinetic energy tends to keep the particles moving apart. 2. The attractive intermolecular forces between particles that tend to draw the particles together.

Types of Attractive Forces
There are several types of attractive intermolecular forces: 1. Hydrogen bonding 2. Dipole-dipole forces 3. Induced-dipole forces 4. London dispersion forces 5. Ionic Bonds
All of the intermolecular forces that hold a liquid together are called cohesive forces.

Hydrogen Bonding
Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen

Hydrogen bonding between ammonia and water

Hydrogen Bonding in DNA
Thymine hydrogen bonds to Adenine
H3C HO NH HO HO O P O O N O N N O O H2N N N O OH O P

OH OH

T

A

Hydrogen Bonding in DNA
Cytosine hydrogen bonds to Guanine
NH2 HO N HO HO O P O O N O H2N HN N O O N N O OH O P

OH OH

C

G

Dipole-Dipole Forces
•Dipole-dipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. •They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together (touching or almost touching).

An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole. •Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. •Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases.

Ion-Dipole Forces

Induced-Dipole Forces
Induced dipole forces result when an ion or a dipole induces a dipole in an atom or a molecule with no dipole. These are weak forces.

Ion-Induced Dipole Forces
An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

Dipole-Induced Dipole Forces
A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

London Dispersion Forces
The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. London forces increase with the size of the molecules.
Fritz London 1900-1954

London Dispersion Forces

London Forces in Hydrocarbons

Boiling point as a measure of intermolecular attractive forces

Relative Magnitudes of Forces
The types of bonding forces vary in their strength as measured by average bond energy.
Ionic bonds Strongest Weakest Hydrogen bonding (12-16 kcal/mol )

Ion-dipole interactions
Dipole-dipole interactions (2-0.5 kcal/mol) Ion induced dipole interactions

Induced Dipole-dipole interactions
London forces (less than 1 kcal/mol)

What Is a Liquid?
A liquid is a state of matter in which a sample of matter: •is made up of very small particles (atoms, molecules, and/or ions). •flows and can change its shape. •is not easily compressible and maintains a relatively fixed volume. The particles that make up a liquid: •are close together with no regular arrangement, •vibrate, move about, and slide past each other.

This bottle contains both liquid bromine [Br2(l), the darker phase at the bottom of the bottle] and gaseous bromine [Br2(g), the lighter phase above the liquid]. The circles show microscopic views of both liquid bromine and gaseous bromine.

More Properties of a Liquid
 Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals).

Capillary Action: Spontaneous rising of a liquid in a narrow tube.

Even More Properties of a Liquid
 Viscosity: Resistance to flow

High viscosity is an indication of strong intermolecular forces

Evaporation is the change of a liquid to a gas.

Evaporation

Microscopic Microscopic view of a liquid. view after evaporation.

When a liquid is heated sufficiently or when the pressure on the liquid is decreased sufficiently, the forces of attraction between molecules do not prevent them from moving apart, and the liquid evaporates to a gas. •Example: The sweat on the outside of a cold glass evaporates when the glass warms. •Example: Gaseous carbon dioxide is produced when the valve on a CO2 fire extinguisher is opened and the pressure is reduced.

Condensation
Condensation is the change from a vapor to a condensed state (solid or
liquid). When a gas is cooled sufficiently or, in many cases, when the pressure on the gas is increased sufficiently, the forces of attraction between molecules prevent them from moving apart, and the gas condenses to either a liquid or a solid. •Example: Water vapor condenses and forms liquid water (sweat) on the outside of a cold glass or can. •Example: Liquid carbon dioxide forms at the high pressure inside a CO2 fire extinguisher.

Microscopic view of a gas.

Microscopic view after condensation.

The vapor pressure of a liquid is the equilibrium pressure of a vapor above its liquid (or solid) The vapor pressure of a liquid varies with its temperature, as the following graph shows for water. The line on the graph shows the boiling temperature for water.

Vapor Pressure

As the temperature of a liquid or solid increases its vapor pressure also increases. Conversely, vapor pressure decreases as the temperature decreases.

Vapor Pressure Revealed

•When a solid or a liquid evaporates to a gas in a closed container, the molecules cannot escape. •Some of the gas molecules will eventually strike the condensed phase and condense back into it. •When the rate of condensation of the gas becomes equal to the rate of evaporation of the liquid or solid, the amount of gas, liquid and/or solid no longer changes. •The gas in the container is in equilibrium with the liquid or solid.

Factors That Affect Vapor Pressure
•Surface Area: the surface area of the solid or liquid in contact with the gas has no effect on the vapor pressure. •Types of Molecules: the types of molecules that make up a solid or liquid determine its vapor pressure. If the intermolecular forces between molecules are:

substance diethyl ether
bromine ethyl alcohol water

vapor pressure at 25oC

0.7 atm
0.3 atm 0.08 atm 0.03 atm

Temperature Dependence of Vapor Pressures
• The vapor pressure above the liquid varies exponentially with changes in the temperature. • The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as:

1 ln P    C RT T

Hvap

Clausius – Clapeyron Equation
• A straight line plot results when ln P vs. 1/T is plotted and has a slope of Hvap/R. • Clausius – Clapeyron equation is true for any two pairs of points.

ln Pvap  

H vap R

1  C T

Write the equation for each and combine to get:

ln

Pvap@T 1 Pvap@T 2

H vap  1 1      R  T2 T1   

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).
E.g. 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. 334 K E.g.2. 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? 33.0 kJ/mol

A liquid boils at a temperature at which its vapor pressure is equal to the pressure of the gas above it. The lower the pressure of a gas above a liquid, the lower the temperature at which the liquid will boil.

Boiling

• As a liquid is heated, its vapor pressure increases until the vapor pressure equals the pressure of the gas above it. •In order to form vapor, the molecules of the liquid must overcome the forces of attraction between them. •The temperature of a boiling liquid remains constant, even when more heat is added. The boiling point of a liquid is the temperature at which its vapor pressure is equal to the pressure of the gas above it. The normal boiling point of a liquid is the temperature at which its vapor pressure is equal to one atmosphere (760 torr).

Factors That Affect the Boiling Point
•Pressure The following graph shows the boiling point for water as a function of the external pressure. The line on the graph shows the normal boiling point for water.

•Types of Molecules: the types of molecules that make up a liquid determine its boiling point. If the intermolecular forces between molecules are: •relatively strong, the boiling point will be relatively high. •relatively weak, the boiling point will be relatively low.

When a liquid is cooled, the average energy of the molecules decreases. At some point, the amount of heat removed is great enough that the attractive forces between molecules draw the molecules close together, and the liquid freezes to a solid.

Freezing

The temperature of a freezing liquid remains constant, even when more heat is removed. The freezing point of a liquid or the melting point of a solid is the temperature at which the solid and liquid phases are in equilibrium.

The rate of freezing of the liquid is equal to the rate of melting of the solid and the quantities of solid and liquid remain constant.
•Types of Molecules: the types of molecules that make up a liquid determine its freezing point. If the intermolecular forces between molecules are: •relatively strong, the freezing point will be relatively high. •relatively weak, the freezing point will be relatively low.

Factors That Affect Freezing Point

Solids

Types of Solids
 Amorphous solids: considerable disorder in their structures (glass).

Types of Solids

 Crystalline Solids: highly regular arrangement of their components

Representation of Components in a Crystalline Solid

Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

Bragg’s Law

xy + yz = n


and

xy + yz = 2d sin

n = 2d sin 

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.

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

Crystal Structures - Cubic

a a
Simple

a
Face-Centered Body-Centered

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 halfatoms in each face give 4 atoms per unit cell.

Crystal Structures - Monoclinic

c

b
a

b

Simple

End Face-Centered

Crystal Structures - Tetragonal

c a a
Simple

Body-Centered

Crystal Structures - Orthorhombic

c

a

b

Simple

End Face-Centered

Body Centered

Face Centered

Crystal Structures – Other Shapes

a

c

a
a a a

c b

120o

a

a

b g
a

Rhombohedral

Hexagonal

Triclinic

Packing in Metals

Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

Closest Packing Holes

Metal Alloys
 Substitutional Alloy: some metal atoms replaced by others of similar size. • brass = Cu/Zn

Metal Alloys
(continued)

 Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.

• steel = iron + carbon

Network Atomic Solids
Some covalently bonded substances DO NOT form discrete molecules.

Diamond, a network of covalently bonded carbon atoms

Graphite, a network of covalently bonded carbon atoms

Molecular Solids
Strong covalent forces within molecules Weak covalent forces between molecules

Sulfur, S8

Phosphorus, P4

Phase Transitions
• Melting: change of a solid to a liquid. • Freezing: change a liquid to a solid. • Vaporization: change of a liquid to a gas. • Condensation: change of a gas to a liquid. • Sublimation : Change of solid to gas • Deposition: Change of a gas to a solid.
H2O(s)  H2O(l) H2O(l)  H2O(s) H2O(l)  H2O(g) H2O(g)  H2O(l) H2O(s)  H2O(g) H2O(g)  H2O(s)

Water phase changes
Temperature remains constant during a phase change.

Energy

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.
E.g. 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. 130 kJ

Phase Diagrams
• 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.

Phase changes by Name

Water

Carbon dioxide

Carbon

Sulfur


				
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