CHAPTER 10. The forces with which molecules attract each other. Intermolecular forces are weaker...
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LIQUIDS & SOLIDS CHAPTER 10
CHAPTER 10. The forces with which molecules attract each other. Intermolecular forces are weaker than ionic or covalent bonds. Intermolecular forces are
The forces with which molecules attract each other.
Intermolecular forces are weaker than ionic or covalent bonds.
Intermolecular forces are responsible for the physical state of a
compound (solid, liquid or gas). TYPES OF INTERMOLECULAR FORCES
Dipole Interactions (between polar molecules) London Dispersion
Forces (between all molecules but mainly force between nonpolar
molecules and noble gases) Hydrogen Bonds (between molecules where
hydrogen is bonded to nitrogen, oxygen and fluorine) INTERMOLECULAR
FORCES
Slide 3
Dipole-dipole forces exist between neutral polar molecules.
Polar molecules need to be close together. Weaker than ion-dipole
forces. There is a mix of attractive and repulsive dipole-dipole
forces as the molecules tumble. If two molecules have about the
same mass and size, then dipole-dipole forces increase with
increasing polarity. DIPOLE-DIPOLE FORCES
Slide 4
Weakest of all intermolecular forces. It is possible for two
adjacent neutral molecules to affect each other. The nucleus of one
molecule (or atom) attracts the electrons of the adjacent molecule
(or atom).For an instant, the electron clouds become distorted.In
that instant a dipole is formed (called an instantaneous dipole).
Polarizability is the ease with which an electron cloud can be
deformed. The larger the molecule (the greater the number of
electrons) the more polarizable. London dispersion forces increase
as molecular weight increases. London dispersion forces exist
between all molecules. London dispersion forces depend on the shape
of the molecule. The greater the surface area available for
contact, the greater the dispersion forces. LONDON DISPERSION
FORCES
Slide 5
H-bonding requires H bonded to an electronegative element (most
important for compounds of F, O, and N). Electrons in the H-X (X =
electronegative element) lie much closer to X than H. H has only
one electron, so in the H-X bond, the + H presents an almost bare
proton to the - X. Therefore, H-bonds are strong. Special case of
dipole-dipole forces. By experiments: boiling points of compounds
with H-F, H-O, and H-N bonds are abnormally high. Intermolecular
forces are abnormally strong. HYDROGEN BONDING
Slide 6
Hydrogen bonds are responsible for: Ice Floating Solids are
usually more closely packed than liquids; Therefore, solids are
more dense than liquids. Ice is ordered with an open structure to
optimize H-bonding. Therefore, ice is less dense than water. In
water the H-O bond length is 1.0 . The OH hydrogen bond length is
1.8 . Ice has waters arranged in an open, regular hexagon. Each + H
points towards a lone pair on O.
Slide 7
EFFECTS OF INTERMOLECULAR FORCES ON PHYSICAL PROPERTIES
Slide 8
Viscosity Viscosity is the resistance of a liquid to flow. A
liquid flows by sliding molecules over each other. The stronger the
intermolecular forces, the higher the viscosity. Surface Tension
Bulk molecules (those in the liquid) are equally attracted to their
neighbors. VISCOSITY & SURFACE TENSION
Slide 9
Surface molecules are only attracted inwards towards the bulk
molecules. Therefore, surface molecules are packed more closely
than bulk molecules. Surface tension is the amount of energy
required to increase the surface area of a liquid. Cohesive forces
bind molecules to each other. Adhesive forces bind molecules to a
surface.
Slide 10
Slide 11
The stronger the intermolecular forces, the higher the boiling
point. For compounds with approximately the same molecular weight:
BOILING POINT
Slide 12
HYDROGEN BONDED TO GROUP 16 ELEMENTS: NOTICE THAT H2O HAS A
GREATER BOILING POINT THAN THE OTHER EVEN THOUGH IT HAS THE LOWEST
MOLECULAR MASS DUE TO THE HYDROGEN BONDING. THE OTHER HYDROGEN
COMPOUNDS EXPERIENCE DIPOLE-DIPOLE BONDING AND THE BOILING POINT
INCREASES WITH INCREASING MOLECULAR MASS. HYDROGEN BONDED TO GROUP
14 ELEMENTS: NOTICE THAT THESE MOLECULES WITH HYDROGEN ARE
EXPERIENCE LONDON DISPERSION FORCES AND ALL BOILING POINTS INCREASE
WITH INCREASE MOLECULAR MASS BOILING POINT TRENDS
Slide 13
Slide 14
Slide 15
A. Evaporation and Vapor Pressure Vaporization or evaporation
Endothermic
Slide 16
MOLECULES AT THE SURFACE HAVE LESS INTERMOLECULAR FORCES ON
THEM THAN THE MOLECULES BELOW THEM. THUS THEY CAN ESCAPE THE LIQUID
PHASE EASIER EVAPORATION
Slide 17
Explaining Vapor Pressure on the Molecular Level Some of the
molecules on the surface of a liquid have enough energy to escape
the attraction of the bulk liquid. These molecules move into the
gas phase. As the number of molecules in the gas phase increases,
some of the gas phase molecules strike the surface and return to
the liquid. After some time the pressure of the gas will be
constant at the vapor pressure.
Slide 18
A. Evaporation and Vapor Pressure Amount of liquid first
decreases then becomes constant Condensation - process by which
vapor molecules convert to a liquid When no further change is
visible the opposing processes balance each other - equilibrium
Vapor Pressure
Slide 19
A. Evaporation and Vapor Pressure Vapor pressure - pressure of
the vapor present at equilibrium with its liquid Vapor Pressure
Vapor pressures vary widely - relates to intermolecular forces
Slide 20
BOILING OCCURS WHEN THE VAPOR PRESSURE OF THE LIQUID EQUALS THE
ATMOSPHERIC PRESSURE. LESS PRESSURE MEANS THAT THE MOLECULES CAN
ESCAPE EASIER AND THUS HAS A LOWER BOILING POINT Boiling Point and
Atmospheric Pressure
Slide 21
Vapor Pressure and Boiling Point Liquids boil when the external
pressure equals the vapor pressure. Temperature of boiling point
increases as pressure increases. Two ways to get a liquid to boil:
increase temperature or decrease pressure. Pressure cookers operate
at high pressure. At high pressure the boiling point of water is
higher than at 1 atm. Therefore, there is a higher temperature at
which the food is cooked, reducing the cooking time required.
Normal boiling point is the boiling point at 760 mmHg (1 atm).
Slide 22
Slide 23
The melting point is the temperature at which a solid is
converted to its liquid phase. In melting, energy is needed to
overcome the attractive forces in the more ordered crystalline
solid. The stronger the intermolecular forces, the higher the
melting point. Because ionic compounds are held together by
extremely strong interactions, they have very high melting points.
With covalent molecules, the melting point depends upon the
identity of the intermolecular force. For compounds of
approximately the same molecular weight: MELTING POINT
Slide 24
C. Energy Requirements for the Changes of State Molar heat of
fusion energy required to melt 1 mol of a substance Molar heat of
vaporization energy required to change 1 mol of a liquid to its
vapor
Slide 25
Generally heat of fusion (enthalpy of fusion) is less than heat
of vaporization: it takes more energy to completely separate
molecules, than partially separate them.
Slide 26
CALCULATING HEAT OF VAPORIZATION & VAPOR PRESSURE OF WATER
Ln(P vap ) =[ (-H vap /R)(1/T)] Textbook : pg 486
Slide 27
TO FIND THE BOILING POINT AT A DIFFERENT PRESSURE Ln (P1/P2) =
(H/R)[ (1/T2) (1/T1)] or Ln (P1/P2) = (H/R)[ (1/T2) (1/T1)] Ln
(P2/P1) = (H/R)[ (1/T1) (1/T2)] H is H vap
Slide 28
Slide 29
HEATING AND COOLING CURVES
Slide 30
Energy Changes Accompanying Phase Changes All phase changes are
possible under the right conditions. The sequence heat solid melt
heat liquid boil heat gas is endothermic. The sequence cool gas
condense cool liquid freeze cool solid is exothermic.
Slide 31
Plot of temperature change versus heat added is a heating
curve. During a phase change, adding heat causes no temperature
change. These points are used to calculate H fus and H vap.
Supercooling: When a liquid is cooled below its melting point and
it still remains a liquid. Achieved by keeping the temperature low
and increasing kinetic energy to break intermolecular forces.
Slide 32
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Slide 34
PHASE DIAGRAMS
Slide 35
Phase diagram: plot of pressure vs. Temperature summarizing all
equilibria between phases. Given a temperature and pressure, phase
diagrams tell us which phase will exist. Any temperature and
pressure combination not on a curve represents a single phase.
PHASE DIAGRAMS
Slide 36
Features of a phase diagram: Triple point: temperature and
pressure at which all three phases are in equilibrium.
Vapor-pressure curve: generally as pressure increases, temperature
increases. Critical point: critical temperature and pressure for
the gas. Melting point curve: as pressure increases, the solid
phase is favored if the solid is more dense than the liquid. Normal
melting point: melting point at 1 atm.
Slide 37
Critical Temperature and Pressure Gases liquefied by increasing
pressure at some temperature. Critical temperature: the minimum
temperature for liquefaction of a gas using pressure. Critical
pressure: pressure required for liquefaction.
Slide 38
Slide 39
Water: The melting point curve slopes to the left because ice
is less dense than water. Triple point occurs at 0.0098 C and 4.58
mmHg. Normal melting (freezing) point is 0 C. Normal boiling point
is 100 C. Critical point is 374 C and 218 atm. Carbon Dioxide:
Triple point occurs at - 56.4 C and 5.11 atm. Normal sublimation
point is -78.5 C. (At 1 atm CO 2 sublimes it does not melt.)
Critical point occurs at 31.1 C and 73 atm.