chapter 10 liquids, solids, and intermolecular...
TRANSCRIPT
Chapter 10 Liquids, Solids, and
Intermolecular Forces
State of Matter Shape and volume Compressibility Ability to Flow
Solid Retains its own shape and volume very low none
LiquidConforms to shape of container, but not
to volumelow moderate
GasConforms to shape
and volume of container
high high
The Three Phases of Matter(A Macroscopic Comparison)
Particles packed close together and are fixed in position
(They may vibrate) Noncompressible
Retain shape and volume when placed in a new container
Do not flow
Molecular View of Phases of Matter
Particles closely packed Particles have some ability to move around
Noncompressible Take the shape of their
container Flow, but don’t have
enough freedom to escape or expand
Molecular View of Phases of Matter
Particles have complete freedom of motion
There is a large amount of space between the
particles
Molecular View of Phases of Matter
Kinetic – Molecular Theory
What state a material is in depends largely on two major factors:
1. the amount of kinetic energy the particles possess 2. the strength of attraction between the particles
These two factors are in competition.
Gas phase particles
Attractive Forces
Particles are attracted to each other by electrostatic forces.
The strength of the attractive forces varies.
The strength of the attractive forces depends on the kind(s) of particles.
The stronger the attractive forces between the particles, the more they resist moving.
The Two Condensed States of Water
Phase Changes
Gases can be condensed.
The amount of kinetic energy the particles have determines the state of matter.
Solids melt when heated.
Liquids boil when heated.
Liquids can be condensed.
Special Properties of Liquids
Surface tension
Viscosity
Capillary Action
Surface Tension
Surface tension -a property of liquids that results from the tendency of liquids to minimize their surface area
To minimize their surface area, liquids form drops that are spherical.
Surface Tension
Viscosity, the resistance of a liquid to flow.
Larger intermolecular attractions ⬇
larger viscosity
Higher temperature ⬇
reduced viscosity
Capillary Action
Capillary action - the ability of a liquid to flow up a thin tube against the influence of gravity
Capillary action is the result of two forces working in conjunction,
adhesive forces attract the outer liquid molecules to the tube’s surface
cohesive forces hold the liquid molecules together
Capillary Action
Adhesive forces pull the surface liquid up the side of a tube, and the cohesive forces pull the interior liquid with it.
The liquid rises up the tube until the force of gravity counteracts the capillary action forces.
The narrower the tube diameter, the higher the liquid will rise up the tube.
Vaporization
Molecules in a liquid are constantly in motion; some molecules have more kinetic energy than others.
If these high energy molecules are at the surface, they may have enough energy to overcome the attractive forces
Therefore – the larger the surface area, the faster the rate of evaporation
This will allow them to escape the liquid and become a vapor.
Dynamic Equilibrium
In a closed container, once the rates of vaporization and condensation are equal, the total amount of vapor and liquid will not change.
Evaporation and condensation are still occurring, but because they are opposite processes, there is no net gain or loss of either vapor or liquid.
Dynamic Equilibrium
When two opposite processes reach the same rate so that there is no gain or loss of material,
We call it a dynamic equilibrium This does not mean there are equal amounts of vapor and liquid –
it means that they are changing by equal amounts
Effect of Intermolecular Attraction on Evaporation and Condensation
Weaker attractive forces ➡ less energy needed to vaporize
Weaker attractive forces ➡ more energy will need to be removed from the vapor molecules before condensation
Weak attractive forces ➡ the faster the evaporation
Liquids that evaporate easily are said to be volatile.
Liquids that do not evaporate easily are called nonvolatile.
The amount of heat energy required to vaporize one mole of the liquid is called the heat of vaporization, ΔHvap, or the enthalpy of vaporization.
Energetics of Vaporization
Calculate the amount of heat needed to vaporize 90.0 g of C3H7OH at its boiling point (ΔHvap = 39.9 kJ/mol)
g mol kJ
39.9 kJ + C3H8O (liquid) C3H8O (gas)
Calculate the mass of water that can be vaporized with 155 kJ of heat at 100 ° (ΔHvap = 40.7 kJ/mol)
kJ mol H2O g H2O
40.7 kJ + H2O (liquid) H2O (gas)
What happens when you heat up a liquid ??
Boiling As a liquid is heated, its temperature rises and the molecules move past each other more vigorously.
Once the temperature reaches the boiling point, the molecules have sufficient energy to overcome the attractions that hold them in contact with other
molecules and the liquid boils.
ΔT
What happens when you heat up a solid ??
Melting As a solid is heated, its temperature rises and the
molecules vibrate more vigorously.
Once the temperature reaches the melting point, the molecules have sufficient energy to overcome some of the attractions that hold them in position and the solid melts.
ΔT
The amount of heat energy required to melt one mole of the solid is called the Heat of Fusion, ΔHfus or the enthalpy of fusion
Generally much less than ΔHvap
Energetics of Melting
How much heat energy is required to raise the temperature of 1.0 mol of water
from -25ºC to 125ºC ??
Quantitative Aspects of Phase Changes
Within a phase, a change in heat is accompanied by a change in temperature which is associated with a change in average kinetic energy of the molecules.
q = ( )(molar heat capacity)(∆T)quantity
ofmatter
J = g x x ºCJg ºC
Quantitative Aspects of Phase Changes
During a phase change, a change in heat occurs at a constant temperature, which is associated with a change in average rotational and translational energy of the molecules, as the average distance between molecules changes.
q = ( )(enthalpy of the phase change)quantity
ofmatter
kJ = mol x kJmol
Heating Curve of Water
Segment 1
Heating 1.00 mole of ice at −25.0 °C up to the melting point, 0.0 °C
q = mass x Cs x ΔTmass of 1.00 mole of ice = 18.0 g Cs = 2.09 J/g·°C
Segment 2
Melting 1.00 mole of ice at the melting point, 0.0 °Cq = n·ΔHfus
n = 1.00 mole of ice ΔHfus = 6.02 kJ/mol
Segment 3
Heating 1.00 mole of water at 0.0 °C up to the boiling point, 100.0 °C
q = mass x Cs x ΔTmass of 1.00 mole of water = 18.0 g Cs = 4.18 J/g·°C
Segment 4
Boiling 1.00 mole of water at the boiling point, 100.0 °C
q = n·ΔHvapn = 1.00 mole of water ΔHvap = 40.7 kJ/mol
Segment 5
Heating 1.00 mole of steam at 100.0 °C up to 125.0 °Cq = mass x Cs x ΔT
mass of 1.00 mole of water = 18.0 g Cs = 2.01 J/g·°C
0.941 6.02 7.5240.7
0.904 56.085
56.1 kJ
Heating Curve of Water
Attractive Forces
Particles are attracted to each other by electrostatic forces
The strength of the attractive forces depends on the kind(s) of particles
The stronger the attractive forces between the particles, the more they resist moving
The strength of the attractions between particles of a substance determines its state.
Kinds of Attractive Forces
Hydrogen Bonds between Molecules An especially strong dipole–dipole attraction resulting
from the attachment of H to an extremely electronegative atom
Dispersion Forces between Molecules Temporary polarity in molecules due to
unequal electron distribution
Dipole–Dipole Attractions between Molecules Permanent polarity in molecules due to their structure
Ion–Dipole Attractions - Not Intermolecular Between mixtures of ionic compounds and polar
compounds (esp. aqueous solutions)
Some molecules are considered nonpolar because of the atoms which they contain and the
arrangement of these atoms in space.
CH4 BH3 C2H2 CO2
Nonpolarizedelectronclouds
But these molecules can all be “condensed.”
Origin of Instantaneous Dipoles
δδ-δδ+
The δδ- charge repels electrons.
The δδ+ charge attracts electrons.
Size of the Induced DipoleThe magnitude of the induced dipole depends on several factors:
Polarizability of the electrons
Volume of the electron cloud
larger molar mass ⇒ more electrons ⇒ larger electron cloud ⇒ increased polarizability ⇒ stronger attractions
Larger molecules have more
electrons, leading to increased
polarizability.
Size of the Induced DipoleShape of the molecule
more surface-to-surface contact ⇒ larger induced dipole
⇒ stronger attraction
Molecules that are flat have more surface
interaction than spherical ones.
Gas Radius Molar Mass B.P.(K)
He 31 4 4.2
Ne 38 20 27
Ar 71 40 87
Kr 88 84 120
Xe 108 131 165
Rn 120 222 211
Effect of Molecular Sizeon Magnitude of Dispersion Force
As the molar mass increases, the number of
electrons increases. Therefore, the strength of
the dispersion forces increases.
The stronger the attractive forces
between the molecules, the
higher the boiling point.
Properties of Straight Chain AlkanesNonPolar Molecules
Effect of Molecular Shapeon Size of Dispersion Force
n-pentane molar mass=72.15
b.p = 36.1 ºC
2-methylbutane molar mass=72.15
b.p = 27.9 ºC
2,2-dimethylpropane molar mass=72.15
b.p = 9.5 ºC
A larger surface-to-surface contact between molecules results in stronger dispersion force attractions and a
higher boiling point.
Kinds of Attractive Forces
Hydrogen Bonds between Molecules An especially strong dipole–dipole attraction resulting
from the attachment of H to an extremely electronegative atom
Dispersion Forces between Molecules Temporary polarity in molecules due to
unequal electron distribution
Dipole–Dipole Attractions between Molecules Permanent polarity in molecules due to their structure
Ion–Dipole Attractions - Not Intermolecular Between mixtures of ionic compounds and polar
compounds (esp. aqueous solutions)
Some molecules are inherently polar because of the atoms which they contain and the
arrangement of these atoms in space.
H2O NH3 CH2O HCl
δ− δ+ A crude representation of a polar molecule
Dipole–Dipole Attractions
Polar molecules have a permanent dipole because of bond polarity and shape
1) dipole moment 2) as well as the always present induced dipole
The permanent dipole adds to the attractive forces between the molecules
Name Formula Molar mass Structure Structure b.p. m.p.
formaldehyde CH2O 30.03 -19.5º -92º
ethane C2H6 30.07 -88º -172º
H
H
H
H
HH
C
C
Effect of Dipole–Dipole Attraction on Boiling and Melting Points
Determine if dipole–dipole attractions occur between CH2Cl2 molecules
Lewis
Structure
Bond
Polarity
Molecule
Polarity Formula
Cl—C 3.0−2.5 = 0.5
polar
C—H 2.5−2.1 = 0.4
nonpolar 4 bonding
areasno lone pairs tetrahedral
shape
polar molecule; therefore dipole–dipole attractions
do exist
Hydrogen Bonding
When a very electronegative atom is bonded to hydrogen, it strongly pulls the bonding electrons toward it:
O─H, N─H, F─H
Because hydrogen has no other electrons, when its electron is pulled away, the nucleus becomes deshielded, exposing the H proton.
The exposed proton acts as a very strong center of positive charge.
H-Bonding in Water
Name Formula Molar mass Structure Structure b.p. m.p.
ethanol C2H6O 46.07 78.2º -114.1º
dimethyl ether C2H6O 46.07 -22º -138.5º
Effect of Hydrogen-Bonding on Boiling and Melting Points
H-Bonds
Very strong intermolecular attractive forces
Stronger than dipole–dipole or dispersion forces
Substances that can hydrogen bond will have higher boiling points and melting points than similar substances that cannot.
But hydrogen bonds are not nearly as strong as chemical bonds.
One of these compounds is a liquid at room temperature (the others are gases). Which one and why?
MM = 30.03PolarNo H-Bonds
MM = 34.03PolarNo H-Bonds
MM = 34.02PolarH-Bonds
Because only hydrogen peroxide has the additional very strong H-bond additional attractions, its intermolecular attractions will be the strongest. We therefore expect hydrogen peroxide to be the liquid.
-19ºC -78ºC +150ºC b.p.
All Molecules
Polar Molecules
Molecules containing O-H, N-H, or F-H
Bonds
Dispersion forces
Dipole forces
H-bonding
Hierarchy of Intermolecular Forces
Comparison of Intermolecular Forces
H2, b.p. -253ºCweak attractions between molecules
HCl, b.p. -85ºCstrong attractions between
molecules
HF, b.p. +20 ºCvery strong attractions
between molecules
Dispersion forces:
Dipole-dipole forces:
Hydrogen bonding:
Ion–Dipole Attractions - Not Intermolecular Between mixtures of ionic compounds and polar
compounds (esp. aqueous solutions)
0.05-40.0 kJ/mol
5-25 kJ/mol
10-40 kJ/mol
40-600 kJ/mol
Non-Bonding (Inter-Molecular) Forces
Bonding Molecular Forces
Types of Crystalline Solids
Types of Crystalline Solids