Liquids and Solids

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Liquids and Solids

• Gas
• low density
• high compressibility
• completely fills its container
• Solid
• high density
• only slightly compressible
• rigid
• maintains its shape

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Liquids 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

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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

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Intermolecular Forces

• Bonds are formed between atoms to form molecules
• intramolecular bonding (within the molecule)

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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

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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) --gt H2O (l) the molecules are still
  unchanged during the phase change

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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

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Dipole-Dipole Forces

• Dipole-dipole forces
• about 1 as strong as covalent or ionic bonds
• become weaker with distance
• unimportant in the gas phase

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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 Hs small
  size

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Hydrogen Bonding

• H-bonding has a very important effect on physical
  properties
• For example, boiling points are greater when
  H-bonding is present

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London Dispersion Forces

• aka Van der Waals forces
• Nonpolar molecules must exert some kind of force
  or they would never solidify

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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)

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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)

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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

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London Dispersion Forces

• Element Freezing Point (oC)
• Helium -269.7
• Neon -248.6
• Argon -189.4
• Krypton -157.3
• Xenon -111.9

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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

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The Liquid State

• Properties of liquids
• low compressibility
• lack of rigidity
• high density (compared to gases)

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The Liquid State

• Surface Tension
• results in droplets when a liquid is poured onto
  a surface
• depends on IMFs

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The 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

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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

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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.

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The 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

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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

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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

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Types of Solids

• three common unit cells exist
• simple cubic
• body centered cubic
• face centered cubic

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Types of Solids

• Amorphous Solids
• noncrystalline
• glass is an example
• disorder abounds

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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 nl 2d sinq

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Types 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

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Types 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)

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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

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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

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Types 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

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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

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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

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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

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Structure and Bonding in Metals

• Hexagonal Closest Packed Structure (hcp)
• aba arrangement
• First Layer
• each sphere is surrounded by six other spheres

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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

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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

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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, Ill let you knowor when
  it shows up on the ACS or AP testthen Ill
  figure it out!)

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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

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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

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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

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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

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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

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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

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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)

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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

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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

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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

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Bonding 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

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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

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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

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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

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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

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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

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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 MOs.
• Electrons cannot be transferred easily to empty
  MOs

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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 LDFs
• graphite is slippery due to these weak LDFs
  between layers

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Covalent 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

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Covalent 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

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Covalent 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

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Covalent Network Solids

• Carbon compounds usually contain long chains with
  C-C bonds
• Silicon compounds usually contain chains with
  Si-O bonds

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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

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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

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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

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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

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Semiconductors

• 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

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Semiconductors

• Enhance conductivity of semiconductors by doping
  the crystal with other atoms

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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

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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

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Semiconductors

• 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

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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

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Types of Solids

• Ionic Solids
• between positive and negative ions
• held by ionic bonds
• electrostatic forces between oppositely charged
  ions

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