Transcript Solids and Liquids IMF, Properties, Changes of State Go to this link

Solids and Liquids
IMF, Properties, Changes of State
Go to this link
http://www.quia.com/jg/455837.html
Liquids and Solids – The Condensed States
Gas
Liquid
Solid
Highly Compressible
Very slightly compressible
Least compressible
Low density
High density
High density
Fills container completely
Does not expand to fill
container – has definite
volume
Rigid and retains its volume
Assumes the shape of its
container
Assumes the shape of its
container
Retains its own shape
Rapid diffusion
Slow diffusion
Extremely slow diffusion –
only at it’s surface
Total disorder; particles
have freedom of motion
and are far apart from one
another
Disordered; particles are
free to move relative to
one another and are close
together
Ordered arrangement;
particles can vibrate but
remain fixed in position
and are close together
High expansion on heating
Low expansion on heating
Low expansion on heating
What causes a substance to be in one
state or another at room temp?
• All particles at room temperature have the
same kinetic energy
• Kinetic molecular theory
 according to the kinetic molecular theory, the
state of a substance at room temperature
depends on the strength of the attractions
between its particles
Intermolecular Forces
• Forces of attraction between neighboring
particles
• Much weaker than bonding forces
• Responsible for the state of the matter and
some physical properties
– e.g. The stronger the attractive forces, the higher
the melting and boiling points
• Intermolecular forces are involved in changes
of state
Strength of Inter vs. Intra
+
-
431 kJ/mol (to break bond)
16 kJ/mol (to separate molecules)
+
-
PowerPoint is posted on Ms. D’s blog
AT END OF CLASS (Warn me when we have 10 minutes left)
• Turn in lab. Place sheet in clear container, lab
notebook in the box.
• Take out homework
• Video consents?
Different Types of IMF
•
•
•
•
Dispersion forces
Dipole-dipole forces
Induced dipole forces
Hydrogen bonds
Dispersion Forces
• The motion of electrons can create an
instantaneous or temporary dipole on an
atom
– For example, if at any one time both of a helium
atom’s electrons are on the same side of the atom
at the same time
• A temporary dipole on one atom can cause, or
induce, a temporary dipole on an adjacent
atom
Dispersion Forces
• dispersion forces – forces of attraction
between induced dipoles
• Exist in all phases of matter
• These forces are found in ALL molecular
compounds
• These are the only kinds of forces that effect
nonpolar compounds.
• Increases with increasing molecular size and
mass
Dispersion Forces
Dispersion Forces
Recap
• Which of the following compounds will have
dispersion forces?
a. HF
b. H2O
c. CH4
d. CH3COOH
Recap
• Which of the following compounds will have
the greatest dispersion force between its
particles? Why?
a. HF
b. H2O
c. CH4
d. CH3COOH
Dipole-Dipole Forces
• Polar covalent molecules have a positive end and a
negative end (permanent dipoles)
• Dipole-dipole forces occur when the positive end of
one molecule is attracted to the negative end of
another
• Only effective when polar molecules are very close
together, but are present in all phases of matter
• For molecules of about the same size, dipole forces
increase with increasing polarity
Dipole-Dipole
Hydrogen Bonds
• a special type of dipole-dipole force
• occurs between molecules containing a hydrogen
atom bonded to a small, highly electronegative
atom (H-N, H-O & H-F)
• The small electronegative atom must have at least
one lone pair of electrons
• The hydrogen in one molecule will be attracted to
the electronegative atom in another molecule
• Strongest IMF
Hydrogen Bonding & Boiling Point
http://thestephenation.blogspot.com/2009/09/hydrogen-bonding.html
http://www.chem.ufl.edu/~itl/2045/lectures/lec_g.html
Properties of Water
• Density of ice is less than the density of
liquid water…
WHY??
You get to figure it out!
Structure of Ice Activity
• In ice, hydrogen bonding causes hexagonal
structures to form
• Prevents other molecules from getting inside the
rings
Arrangement of molecules in
liquid water
Arrangement of molecules in
ice
Water
• unexpectedly high boiling point causes it to be a
liquid at room temp
• other hydrogen compounds are corrosive gases at
room temp
• can absorb or release relatively large quantities of
heat without large temp changes
• has relatively high surface tension
• has very high heat of vaporization
• called the universal solvent b/c it can dissolve so
many things
IMF Summary
Properties of Liquids
• Only slightly compressible; not a discernable
difference when compressed
• Have much greater densities than their vapors
• Fluidity: ability to flow
– Liquids can diffuse through one another, but at a
much slower rate than gases
Properties of Liquids
• Physical properties are determined mainly by the
nature and strength of IMF present between
molecules
• Viscosity: resistance to flow
– Determined by:
• The stronger the attractive forces, the higher
the viscosity
• The larger the particles, the higher the viscosity
• Increases as temp decreases
high viscosity = thick/slow flowing
low viscosity = thin/fast flowing
Properties of Liquids
• Surface Tension: the imbalance of forces at
the surface of a liquid
• The uneven forces make the surface behave as
if it has a tight film stretched across it
• The stronger the intermolecular forces, the
higher the surface tension
Surface Tension
Properties of Liquids
•
Surfactants: compounds that lower the
surface tension of water
– Frequently added to detergents
•
Capillary action: movement of a liquid
through narrow spaces
Properties of Solids
• Have extremely strong intermolecular forces
in order for solids to have definite shape and
volume
• Particle arrangement causes solids to almost
always have higher densities than liquids
• Ice is an exception: it expands when it freezes
because of the way the particles arrange
themselves during the freezing process
Properties of Solids
• Particle arrangements cause different types of
solids:
– Crystalline solids
– Amorphous solids
Crystalline Solids
• particles exist in a highly ordered repeating
pattern
– Precious stones, sugar, Ionic solids – salts, atomic
- Metallic solids
• 7 principle crystal patterns
• atoms, ions, or molecules arranged in an orderly,
geometric, 3-D structure
• Smallest arrangement of repeated crystal is called a
unit cell
Examples of how particles can be arranged in a cubic crystal.
Amorphous Solid
• Solid arrangement of particles lacking
a regular repeating pattern
• Like liquids that have been cooled to
such a low temp that their viscosity
becomes very high
• glass, rubber, wax, tar
• Particles are trapped in a disordered
arrangement that is characteristic of
liquids
• get softer over a wide range of
temperatures before melting
–Molecular such as sucrose or ice whose
constituent particles are molecules held together
by the intermolecular forces.
Molecular Solids
• Type of Particles: atoms or molecules
• Held together by dispersion forces, dipoledipole forces or hydrogen bonds
• Most are NOT solids at room temperature
• Poor conductors of heat and electricity (don’t
contain ions)
• Soft w/ low to moderate melting points
• Examples are sucrose, ice, most organics
Sodium chloride
Cupric chloride
Ionic Solids
• Type of particles: cations & anions
• Type and ratio of ions determine the shape of
the crystaline structure
• The network of attractions that extend
through an ionic compound gives these
compounds their high melting points and
hardness
• Hard and brittle, poor electrical and thermal
conductors in solid state
• Examples are salts (NaCl, KBr, MgSO4)
Covalent Network Solids-Atomic
• Atoms that can form multiple covalent bonds
• Properties: very hard, very high melting point,
often poor thermal and electrical conductors
• Most allotropes exist in this form
– Allotropes are forms of the same element that
have different bonding patterns of arrangement
• Examples include diamonds and graphite,
silicon, quartz (SiO2)
Graphite
Diamond
Covalent network solids such as quartz where
atoms are held together by 3-D networks of
covalent bonds. Here the hexagonal pattern of Si
(violet) and O (red) atoms in structure matches
the hexagonal crystal shape
Buckminster fullerene
Carbon microtubules
Gold
Copper
Silver
Metallic Solids - Atomic
• Consist of positive metal ions surrounded by a
sea of mobile electrons
• Mobile electrons make metals malleable,
ductile, and good conductors of heat and
electricity
• Also possible to form alloys by this type of
bonding
substitutional vs. interstitial
Why malleable and ductile?
Pure Solids
Amorphous Solids:
- any solid without a definite
lattice ( a well organized
pattern or structure)
ex. glass
Ionic Solids:
- any solid made from
metal/nonmetal containing
compounds.
ex. NaCl, KBr, Ca3(PO4)2
- generally known as “salts”
- held together as a solid through
ionic bonds
Crystalline Solids:
- any solid with a definite lattice (a
well organized pattern or structure
- a solid with any one of 3 unit cells
(simple cubic, face-centered cubic, or
body-centered cubic)
Molecular Solids:
- any solid made from
nonmetal (covalent)
compounds
ex. sugar, sand, fats,
plastics
- held together as a solid
through dipole-dipole and/or
dispersion forces
Metalic Solids:
- any solid created from
combining metal elements
ex. Na, Ag, brass, bronze,
steel
- generally known as “alloys”
- held together as a solid
through delocalized covalent
bonds
Atomic Solids:
- a solid created from one
type of element
Network Solids:
- any solid created from combining
nonmetal elements
ex. diamond; graphite; red,
white or black phosphorus
- generally known as “allotropes”
- held together as a solid through
covalent bonds.
Group 8A/18 Solids:
- any solid created when
heavy noble gas elements
solidify
ex. Rn, Xe, Kr
- held together as a solid
through dispersion forces
Endothermic
• Always involve a
change in energy
• Energy is needed
either to
overcome or
form attractive
forces between
particles
Exothermic
Phase Changes
Vaporization
• The change of state from a liquid to a gas
• Vapor – refers to the gaseous state of a
substance that is normally a liquid or a solid at
room temp
• Two methods of vaporization:
– Evaporation
– Boiling
Evaporation
• Occurs at the surface of a liquid
• Occurs b/c molecules close to the surface have
enough energy to overcome the attractions of
neighboring molecules and escape
• Slower molecules stay in the liquid state
• Rate of evaporation increases as temp increases
• volatile – evaporates easily, molecules don’t exert a
very strong attractive force upon one another
• evaporative cooling – molecules with higher kinetic
energy escape, the avg. kinetic energy of the
remaining molecules decrease resulting in lower
temperature.
Vapor Pressure
• vapor pressure - the pressure of the vapor resulting
from evaporation of a liquid (or solid) above a
sample of the liquid (or solid) in a closed container
• Increases steadily as temperature increases
• The gas in the container is in equilibrium with the
liquid or solid.
Boiling
• Occurs within the liquid
• rapid vaporization of liquid
• boiling point: temp at which vapor pressure equals
atmospheric pressure
• heat of vaporization – the amount of heat required
to vaporize a given amount of liquid
• Liquids with strong intermolecular attractions have
high heats of vaporization
• Although energy is added, temp remains constant
during the phase change
Condensation
• Change of a gas to a liquid
• Molecules of vapor can return to the liquid state by
colliding with the liquid surface
– The particles become trapped by the intermolecular
attractions of the liquid
• Rate of condensation increases as the # of vapor
particles increases
• When the rate of vaporization and rate of
condensation are equal, a state of dynamic
equilibrium is reached (liquid-vapor equilibrium)
Melting and Freezing
• Melting point/freezing point: temp at which solid and
liquid forms exist in equilibrium
• requires smaller potential energy changes than
vaporization
• Particles are about the same distance apart in the solid
and liquid forms
• not affected significantly by a change in external pressure
• Heat of fusion – ΔHfus amount of heat required to
convert a solid to a liquid
– depends on the strength of attractive forces between molecules
When the rate of freezing is the same as the rate of
melting, the amount of ice and the amount of water won't
change. The ice and water are said to be in dynamic
equilibrium with each other. The ice is melting, and the
water is freezing, but both are occurring at the same rate,
so there is no net change in either quantity.
Sublimation and Deposition
• Sublimation: solid goes directly to a gas
• Deposition is the reverse process
• solids exert vapor pressure
– tends to be much lower than liquid vapor pressure
• Solids with high vapor pressure sublime
relatively easily
• Solids without strong attractive forces sublime
readily, mostly molecular solids
Heating Curves
• Graphic illustrations of phase changes
• Plot of temp of a sample as a function of time
• Notice temp remains constant during phase
changes while amount of energy varies
Temperature
Heating Curve
VAPORIZATION
CONDENSATION
MELTING
FREEZING
TIME or ENERGY
Heating Curve of Water
A: Rise in temperature as ice absorbs heat.
B: Absorption of heat of fusion.
C: Rise in temperature as liquid water absorbs heat.
D: Water boils and absorbs heat of vaporization.
E: Steam absorbs heat and thus increases its temperature.
The above is an example of a heating curve. One could reverse the process, and
obtain a cooling curve. The flat portions of such curves indicate the phase changes.
Phase Diagrams
• Diagram that relates the states of a substance to
temp and pressure
• State depends on temp and pressure
• 2 states can exist simultaneously at certain temps
and pressures
• Triple point: the temp and pressure when all three
states exist at the same time
• Critical point – the temp and pressure combination
at which a gas form of a substance is converted to a
liquid
Triple Point
TRIPLE POINT - The temperature and pressure at which the solid,
liquid, and gas phases exist simultaneously.
CRITICAL POINT - The temperature above which a substance will
always be a gas regardless of the pressure.
NOTE:
•The line between the solid and liquid phases is a curve of all the
freezing/melting points of the substance.
•The line between the liquid and gas phases is a curve of all the boiling
points of the substance.
Freezing Point - The temperature at which the solid and liquid phases
of a substance are in equilibrium at atmospheric pressure.
Boiling Point - The temperature at which the vapor pressure of a liquid
is equal to the pressure on the liquid.
Normal (Standard) Boiling Point - The temperature at which the
vapor pressure of a liquid is equal to standard pressure (1.00 atm = 760
mmHg = 760 torr = 101.325 kPa)