Liquids and Solids by vivi07


									                       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) DHofus = 6.02 kJ/mol
    • H2O(l) --> 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
          Liquids and Solids
• Densities of the three states of
  – 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
    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
  – e.g., H2O(s) --> H2O (l) …the molecules
    are still unchanged during the phase
      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
  – become weaker with distance
  – unimportant in the gas phase
          Hydrogen Bonding
• A particularly strong dipole-dipole
• 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
  London Dispersion Forces
• London dispersion forces (LDF)
  – due to an instantaneous dipole moment
    • created when electrons move about the
    • a temporary nonsymmetrical electron
      distribution can develop (I.e., all the
      electrons will shift to one side of the
  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
  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”
  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
  – 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
        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
• Adhesive forces - forces between the
  liquid molecules and the polar (glass)
  – 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
       – unit cell - smallest repeating unit of the
        Types of Solids
• three common unit cells exist
  – simple cubic
  – body centered cubic
  – face centered cubic
          Types of Solids
• Amorphous Solids
  – noncrystalline
  – glass is an example
  – disorder abounds
         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
  – obtain a diffraction pattern
  – Bragg equation: nl = 2d sinq
            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
             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)
          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
  – Molecular Solids have discrete, relatively
    small molecules at the lattice points
  – Ionic solids have ions at the lattice
         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
 Structure and Bonding in Metals
• Properties of Metals
  – high thermal conductivity
  – high electrical conductivity
  – malleability (metals can be pounded
  – 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
  – 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
    • Hexagonal Closest Packed Structure
    • Cubic Closest Packed Structure
 Structure and Bonding in Metals

• Hexagonal Closest Packed Structure
  – aba arrangement
  – First Layer
    • each sphere is surrounded by six other
 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
  (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
 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
  – strong
  – nondirectional
• It must be difficult to separate
  atoms, but easy to move them (as long
  as the atoms stay in contact with each
 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
    • these electrons can conduct heat and
    • meanwhile, the metal ions can move around
 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
  – 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
    • 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
         – 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 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
      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
     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
  – 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
• Silicon is a semiconductor
  – gap between filled and empty MO’s is
    smaller than the gap found in diamond (a
  – 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
• Enhance conductivity of
  semiconductors by doping the crystal
  with other atoms
• 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
• 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
• p-type semiconductor - dope Si with
  atoms with less valence e-’s (e.g. with
   – 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
• 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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