Liquids and Solids• Gas
– low density – high compressibility – completely fills its container
• Solid– high density– only slightly compressible– rigid– maintains its shape
Liquids and Solids• Liquids
– properties lie between those of solids and gases• H2O(s) --> H2O(l)
kJ/mol
• H2O(l) --> H2O(g) Hovap = 40.7 kJ/mol
– large value of Hvap 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
Liquids 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
Intermolecular Forces
• Bonds are formed between atoms to form molecules– intramolecular bonding (within
the molecule)
Intermolecular 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
Intermolecular 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) --> H2O (l) …the molecules are still unchanged during the phase change
Dipole-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
Dipole-Dipole Forces
• Dipole-dipole forces – about 1% as strong as covalent or
ionic bonds– become weaker with distance– unimportant in the gas phase
Hydrogen 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 H’s
small size
Hydrogen Bonding
• H-bonding has a very important effect on physical properties– For example, boiling points are
greater when H-bonding is present
London Dispersion Forces• aka Van der Waals forces• Nonpolar molecules must exert
some kind of force or they would never solidify
London 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)
London 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)
London 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
London Dispersion Forces
• Element Freezing Point (oC) Helium -269.7 Neon -248.6 Argon -189.4
Krypton -157.3 Xenon -111.9
London 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
The Liquid State
• Properties of liquids– low compressibility– lack of rigidity– high density (compared to gases)
The Liquid State
• Surface Tension– results in droplets when a liquid
is poured onto a surface– depends on IMF’s
The Liquid State
– Molecules at the surface experience an uneven pull, only from the sides and below. Molecules in the interior are surrounded by IMF’s• 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 IMF’s have high
surface tensions
The 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
The 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.
The Liquid State
• Viscosity– Measure of a liquid’s resistance to
flow– Depends on strength of IMF’s
between liquid molecules• molecules with large IMF’s are very
viscous• Large molecules that can get tangled up
with each other lead to high viscosity
The 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
Types 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
Types of Solids
• three common unit cells exist– simple cubic– body centered cubic– face centered cubic
Types 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: n = 2d sin
Types of Solids
• Where n is an integer is the wavelength of the x-rays• d is the distance between the atoms 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
Types of Solids• Example:• x-rays of wavelength 1.54 A were
used to analyze an aluminum crystal. A reflection was produced at = 19.3 degrees. Assuming n = 1, calculate the distance d between the planes of atoms producing the reflection.
• (D = 2.33 A)
Types 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
Types 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
Types of Solids
• Different bonding present in these solids results in dramatically different properties
• Element (atomic solid) M.P. (oC)Argon -189C(diamond) 3500Cu 1083
Structure 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
Structure 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
Structure 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
Structure and Bonding in Metals
• Hexagonal Closest Packed Structure (hcp)– aba arrangement– First Layer
• each sphere is surrounded by six other spheres
Structure 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
Structure 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
Structure 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, I’ll let you know…
or when it shows up on the ACS or AP test…then I’ll figure it out!)
Structure 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
Structure 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
Bonding 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
Bonding 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
Bonding 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
Bonding 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
Bonding 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)
Bonding 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
Bonding 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
Bonding 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
Bonding Models for Metals
• Conduction electrons– the electrons that can be excited to
empty MO’s
• Conduction bands– the empty MO’s that can accept the
conducting electrons
Metal 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
Metal 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
Metal 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
Metal 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
Covalent 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
Covalent 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 MO’s.
– Electrons cannot be transferred easily to empty MO’s
Covalent 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
LDF’s• graphite is slippery due to these weak
LDF’s between layers
Covalent Network Solids
• Graphite– since the C atoms are sp2 hybridized,
there is one 2p orbital left– the 2p orbitals form molecular
orbitals above the plane of the rings– the electrons are delocalized in these
molecular orbitals• these delocalized electrons allow for
electrical conductivity
Covalent Network Solids
• Convert graphite to diamonds– apply pressure…150,000 atm at
2800oC– requires such high pressure and
temperature to completely break the bonds in graphite and rearrange them to yield diamond
Covalent Network Solids
• Silicon– makes up many compounds found in
the earth’s crust– silicon:geology as carbon:biology– Even though silicon and carbon are in
the same family, the structures of silicon and carbon compounds are very different
Covalent Network Solids
• Carbon compounds usually contain long chains with C-C bonds
• Silicon compounds usually contain chains with Si-O bonds
Covalent 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
Covalent 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
Covalent 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
Covalent 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
Semiconductors• Silicon is a semiconductor
– gap between filled and empty MO’s 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
Semiconductors• 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
Semiconductors
• 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
Semiconductors• p-type semiconductor - dope Si
with atoms with less valence e-’s (e.g. with B)– B’s 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
Semiconductors
• 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
Types of Solids
• Ionic Solids– between positive and negative ions– held by ionic bonds
•electrostatic forces between oppositely charged ions
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