Title: Liquids and Solids
1Liquids and Solids
- Gas
- low density
- high compressibility
- completely fills its container
- Solid
- high density
- only slightly compressible
- rigid
- maintains its shape
2Liquids and Solids
- Liquids
- properties lie between those of solids and gases
- H2O(s) --gt H2O(l) DHofus 6.02 kJ/mol
- H2O(l) --gt H2O(g) DHovap 40.7 kJ/mol
- large value of DHvap suggests greater changes in
structure in going from a liquid to a gas than
from a solid to liquid - suggests attractive forces between the molecules
in a liquid, though not as strong as between the
molecules of a solid
3Liquids and Solids
- Densities of the three states of water
- H2O(g) D 3.26 x 10-4g/cm3 (400oC)
- H2O(l) D 0.9971 g/cm3 (25oC)
- H2O(s) D 0.9168 g/cm3 (OoC)
- Similarities in the densities of the liquid and
solid state indicate similarities in the
structure of liquids and solids
4Intermolecular Forces
- Bonds are formed between atoms to form molecules
- intramolecular bonding (within the molecule)
5Intermolecular Forces
- The properties of liquids and solids are
determined by the forces that hold the components
of the liquid or solid together - may be covalent bonds
- may be ionic bonds
- may weaker intermolecular forces between molecules
6Intermolecular Forces
- During a phase change for a substance like water
- the components of the liquid or solid remain
intact - the change of state is due to the changes in the
forces between the components - e.g., H2O(s) --gt H2O (l) the molecules are still
unchanged during the phase change
7Dipole-Dipole Forces
- Polar molecules
- line up in an electric field
- positive end of molecule will line up with the
negative pole of the electric field while the
negative end of the molecule will line up with
the positive pole - can attract each other
- positive end of one molecule will attract the
negative end of another molecule
8Dipole-Dipole Forces
- Dipole-dipole forces
- about 1 as strong as covalent or ionic bonds
- become weaker with distance
- unimportant in the gas phase
9Hydrogen Bonding
- A particularly strong dipole-dipole force
- When hydrogen is covalently bonded to a very
electronegative atom such as N, O, or F - Very strong due to
- great polarity of the bond between H and the N, O
or F - close approach of the dipoles due to Hs small
size
10Hydrogen Bonding
- H-bonding has a very important effect on physical
properties - For example, boiling points are greater when
H-bonding is present
11London Dispersion Forces
- aka Van der Waals forces
- Nonpolar molecules must exert some kind of force
or they would never solidify
12London Dispersion Forces
- London dispersion forces (LDF)
- due to an instantaneous dipole moment
- created when electrons move about the nucleus
- a temporary nonsymmetrical electron distribution
can develop (I.e., all the electrons will shift
to one side of the molecule)
13London Dispersion Forces
- The instantaneous dipole moment can induce an
instantaneous dipole moment in a neighboring
molecule, which could induce another
instantaneous dipole moment in a neighboring
molecule, etc. (like a wave in the stands of a
football game)
14London Dispersion Forces
- The LDF is very weak and short-lived
- To form a solid when only LDF exists requires
very low temperatures - the molecules or atoms must be moving slowly
enough for the LDF to hold the molecules or atoms
together in a solid unit
15London Dispersion Forces
- Element Freezing Point (oC)
- Helium -269.7
- Neon -248.6
- Argon -189.4
- Krypton -157.3
- Xenon -111.9
16London Dispersion Forces
- Notice that as the MM of the noble gas increases,
the freezing point increases - This implies that the LDF between the atoms is
stronger as the MM increases - Large atoms with many electrons have an increased
polarizability (the instantaneous dipole would be
larger), resulting in a larger London Dispersion
Force between the atoms than between smaller atoms
17The Liquid State
- Properties of liquids
- low compressibility
- lack of rigidity
- high density (compared to gases)
18The Liquid State
- Surface Tension
- results in droplets when a liquid is poured onto
a surface - depends on IMFs
19The Liquid State
- Molecules at the surface experience an uneven
pull, only from the sides and below. Molecules
in the interior are surrounded by IMFs - Uneven pull results in liquids assuming a shape
with minimum surface area - Surface tension is a liquids resistance to an
increase in surface area. - Liquids with high IMFs have high surface
tensions
20The Liquid State
- Capillary Action
- Exhibited by polar molecules
- The spontaneous rising of a liquid in a narrow
tube - due to two different forces involving the liquid
21The Liquid State
- Cohesive forces - IMF between the liquid
molecules - Adhesive forces - forces between the liquid
molecules and the polar (glass) container - adhesive forces tend to increase the surface area
- cohesive forces counteract this
- Concave meniscus (water) - indicates adhesive
forces of water towards the glass is greater than
the cohesive forces between the water molecules. - Convex meniscus (nonpolar substances such as
mercury) shows cohesive forces is greater than
adhesive forces.
22The Liquid State
- Viscosity
- Measure of a liquids resistance to flow
- Depends on strength of IMFs between liquid
molecules - molecules with large IMFs are very viscous
- Large molecules that can get tangled up with each
other lead to high viscosity
23The Liquid State
- So what does a liquid look like?
- A liquid contains many regions where the
arrangements of the components are similar to
those of a solid - There is more disorder in a liquid than in a
solid - There is a smaller number of regions in a liquid
where there are holes present
24Types of Solids
- Ways to classify solids
- Crystalline vs. Amorphous Solid
- Crystalline solids
- regular arrangement of components
- positions of components represented by a lattice
- unit cell - smallest repeating unit of the
lattice
25Types of Solids
- three common unit cells exist
- simple cubic
- body centered cubic
- face centered cubic
26Types of Solids
- Amorphous Solids
- noncrystalline
- glass is an example
- disorder abounds
27Types of Solids
- X-ray diffraction
- used to determine the structures of crystalline
solids - diffraction occurs when beams of light are
scattered from a regular array of points - obtain a diffraction pattern
- Bragg equation nl 2d sinq
28Types of Solids
- Where n is an integer
- l is the wavelength of the x-rays
- d is the distance between the atoms
- q is the angle of incidence and reflection
- Use x-ray diffraction to determine bond lengths,
bond angles, determine complex structures, test
predictions of molecular geometry
29Types of Solids
- Example
- x-rays of wavelength 1.54 A were used to analyze
an aluminum crystal. A reflection was produced
at q 19.3 degrees. Assuming n 1, calculate
the distance d between the planes of atoms
producing the reflection. - (D 2.33 A)
30Types of Solids
- Types of Crystalline Solids
- Ionic Solids (e.g. NaCl)
- Molecular Solids (e.g. C6H12O6)
- Atomic Solids which include
- Metallic Solids
- Covalent Network Solids
31Types of Solids
- Classify solids according to what type of
component is found at the lattice point (of a
unit cell) - Atomic Solids have atoms at the lattice points
- Molecular Solids have discrete, relatively small
molecules at the lattice points - Ionic solids have ions at the lattice points
32Types of Solids
- Different bonding present in these solids results
in dramatically different properties - Element (atomic solid) M.P. (oC)
- Argon -189
- C(diamond) 3500
- Cu 1083
33Structure and Bonding in Metals
- Properties of Metals
- high thermal conductivity
- high electrical conductivity
- malleability (metals can be pounded thin)
- ductility (metals can be drawn into a fine wire)
- durable
- high melting points
34Structure and Bonding in Metals
- Properties are due to the nondirectional covalent
bonding found in metallic crystals - Metallic crystal
- contains spherical atoms packed together
- atoms are bonded to each other equally in all
directions
35Structure and Bonding in Metals
- Closest Packing
- most efficient arrangement of these uniform
spheres - Two possible closest packing arrangements
- Hexagonal Closest Packed Structure
- Cubic Closest Packed Structure
36Structure and Bonding in Metals
- Hexagonal Closest Packed Structure (hcp)
- aba arrangement
- First Layer
- each sphere is surrounded by six other spheres
37Structure and Bonding in Metals
- Second Layer
- the spheres do not lie directly over the spheres
in the first layer - the spheres lie in the indentations formed by
three spheres - Third Layer
- the spheres lie directly over the spheres in the
first layer
38Structure and Bonding in Metals
- Cubic Closest Packed Structure (ccp)
- abc arrangement
- First and Second Layers are the same as in
hexagonal closest packed structure - Third Layer
- the spheres occupy positions such that none of
the spheres in the third layer lie over a sphere
in the first layer
39Structure and Bonding in Metals
- Finding the net number of spheres in a unit cell
- important for many applications involving solids
- (when I figure it out, Ill let you knowor when
it shows up on the ACS or AP testthen Ill
figure it out!)
40Structure and Bonding in Metals
- Examples of metals that are ccp
- aluminum, iron, copper, cobalt, nickel
- Examples of metals that are hcp
- zinc, magnesium
- Calcium and some other metals can go either way
41Structure and Bonding in Metals
- Some metals, like the alkali metals are not
closest packed at all - may be found in a body centered cubic (bcc) unit
cell where there are only 8 nearest neighbors
instead of the 12 in the closest packed structures
42Bonding Models for Metals
- The model must account for the typical physical
properties of metals - malleability
- ductility
- efficient and uniform conduction of heat and
electricity in all directions - durability of metals
- high melting points
43Bonding Models for Metals
- To account for these physical properties, the
bonding in metals must be - strong
- nondirectional
- It must be difficult to separate atoms, but easy
to move them (as long as the atoms stay in
contact with each other
44Bonding Models for Metals
- Electron Sea Model (simplest picture)
- Positive Metal ions (Metal cations) are
surrounded by a sea of valence electrons - the valence electrons are mobile and loosely held
- these electrons can conduct heat and electricity
- meanwhile, the metal ions can move around easily
45Bonding Models for Metals
- Band Model or Molecular Orbital (MO) model
- related to the electron sea model
- more detailed view of the electron energies and
motions
46Bonding Models for Metals
- MO model
- electrons travel around the metal crystal in
molecular orbitals formed from the atomic
orbitals of the metal atoms - In atoms like Li2 or O2, the space between the
energies of the molecular orbitals is relatively
wide (big energy difference between the orbitals)
47Bonding Models for Metals
- However, when many metal atoms interact, the
molecular orbital energy levels are very close
together - Instead of separate, discrete molecular orbitals
with different energies, the molecular orbitals
are so close together in energies, that they form
a continuum of levels, called bands
48Bonding Models for Metals
- Core electrons of metals are localized
- the core electrons belong to a particular metal
ion - The valence electrons of metals are delocalized
- the valence electrons occupy partially filled,
closely spaced molecular orbitals
49Bonding Models for Metals
- Thermal and Electrical conductivity
- metals conduct heat and electricity because of
highly mobile electrons - electrons in filled molecular orbitals get
excited (from added heat or electricity) - these electrons move into higher energy, empty
molecular orbitals
50Bonding Models for Metals
- Conduction electrons
- the electrons that can be excited to empty MOs
- Conduction bands
- the empty MOs that can accept the conducting
electrons
51Metal Alloys
- Alloy
- a substance that contains a mixture of elements
and has metallic properties - Metals can form alloys due to the nature of their
structure and bonding
52Metal Alloys
- Two types of alloys
- Substitutional alloy
- host metal atoms are replaced by other metal
atoms of similar size - ex brass is an alloy of zinc and copper
- sterling silver - silver and copper
- pewter - tin and copper
- solder - lead and tin
53Metal Alloys
- Interstitial Alloys
- formed when some of the holes in the closest
packed structure are filled with smaller atoms - ex steel is an alloy with carbon filling the
interstices of an iron crystal
54Metal Alloys
- Presence of interstitial atoms changes the
properties of the host metal - Iron - soft, ductile, malleable
- Steel - harder, stronger, less ductile than pure
iron - due to directional bonds between carbon and iron
- More carbon, harder steel
55Covalent Network Solids
- Covalent Network Solids
- Macromolecule
- A giant molecule containing numerous covalent
bonds holding atoms together - Properties
- brittle
- do not conduct heat or electricity
- very high melting points
56Covalent Network Solids
- Typical Covalent Network Solids
- Diamond (Cdia) and Graphite (Cgraphite)
- Diamond
- each C atom is covalently bonded to four other C
atoms in a tetrahedral arrangement - sp3 hybridization of the C atoms
- Using MO model, diamond is a nonconductor due to
the large space between the empty MOs. - Electrons cannot be transferred easily to empty
MOs
57Covalent Network Solids
- Graphite
- slippery, black, and a conductor
- different bonding than diamond
- there are layers of sp2 hybridized C atoms in
fused six member rings - the layers are held loosely with weak LDFs
- graphite is slippery due to these weak LDFs
between layers
58Covalent Network Solids
- Graphite
- since the C atoms are sp2 hybridized, there is
one 2p orbital left - the 2p orbitals form p molecular orbitals above
the plane of the rings - the electrons are delocalized in these p
molecular orbitals - these delocalized electrons allow for electrical
conductivity
59Covalent Network Solids
- Convert graphite to diamonds
- apply pressure150,000 atm at 2800oC
- requires such high pressure and temperature to
completely break the bonds in graphite and
rearrange them to yield diamond
60Covalent Network Solids
- Silicon
- makes up many compounds found in the earths
crust - silicongeology as carbonbiology
- Even though silicon and carbon are in the same
family, the structures of silicon and carbon
compounds are very different
61Covalent Network Solids
- Carbon compounds usually contain long chains with
C-C bonds - Silicon compounds usually contain chains with
Si-O bonds
62Covalent Network Solids
- Silica
- Empirical formula - SiO2
- sand, quartz are composed of SiO2
- Si is the center of a tetrahedron, forming single
bonds with four oxygen atoms, which are shared by
other Si atoms - A covalent network solid like diamond
63Covalent Network Solids
- Silicates
- related to silica
- found in most rocks, soils, and clays
- based on interconnected SiO4 tetradera
- unlike silica, silicates contain silicon-oxygen
anions - silicates need positive metal cations to balance
the negative charge
64Covalent Network Solids
- Glass
- an amorphous solid
- formed when silica is heated and cooled rapidly
- more closely resembles a viscous solution than a
crystalline solid - adding different substances to the melted silica
results in different properties for the glass
65Covalent Network Solids
- Add B2O3 to produce glass for labware (pyrex)
- very little expansion or contraction with large
temperature changes - Add K2O to produce a very hard glass that can be
ground for eyeglasses or contacts
66Semiconductors
- Silicon is a semiconductor
- gap between filled and empty MOs is smaller than
the gap found in diamond (a nonconductor) - a few electrons can get excited and cross the gap
in silicon - at higher temperatures, more electrons can get
across, so conductivity increases at higher
temperatures
67Semiconductors
- Enhance conductivity of semiconductors by doping
the crystal with other atoms
68Semiconductors
- N - type semiconductor - dope Si with atoms with
more valence e-s (e.g. with As) - the extra electrons from As can conduct an
electric current
69Semiconductors
- analogy Given a row in a movie theater filled
with people. Each person has a bag of popcorn.
One person has two bags of popcorn. Passing one
bag of popcorn (the extra electron) down the row
is like electricity being conducted in an n-type
semiconductor
70Semiconductors
- p-type semiconductor - dope Si with atoms with
less valence e-s (e.g. with B) - Bs three valence e- leave a hole in an MO.
- Another e- could move into the hole, but it would
leave another hole for another electron to fill
71Semiconductors
- Analogy In a movie theater, a row of seats is
filled, except for one seat. One person could
get up out of his seat and move into the empty
seat. The next person could then move into the
newly emptied seat, and so on - the p in p-type refers to the positive hole
formed with a missing valence electron
72Types of Solids
- Ionic Solids
- between positive and negative ions
- held by ionic bonds
- electrostatic forces between oppositely charged
ions
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