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Liquids and Solids

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Title: Liquids and Solids


1
Liquids and Solids
  • Gas
  • low density
  • high compressibility
  • completely fills its container
  • Solid
  • high density
  • only slightly compressible
  • rigid
  • maintains its shape

2
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

3
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

4
Intermolecular Forces
  • Bonds are formed between atoms to form molecules
  • intramolecular bonding (within the molecule)

5
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

6
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

7
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

8
Dipole-Dipole Forces
  • Dipole-dipole forces
  • about 1 as strong as covalent or ionic bonds
  • become weaker with distance
  • unimportant in the gas phase

9
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

10
Hydrogen Bonding
  • H-bonding has a very important effect on physical
    properties
  • For example, boiling points are greater when
    H-bonding is present

11
London Dispersion Forces
  • aka Van der Waals forces
  • Nonpolar molecules must exert some kind of force
    or they would never solidify

12
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)

13
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)

14
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

15
London Dispersion Forces
  • Element Freezing Point (oC)
  • Helium -269.7
  • Neon -248.6
  • Argon -189.4
  • Krypton -157.3
  • Xenon -111.9

16
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

17
The Liquid State
  • Properties of liquids
  • low compressibility
  • lack of rigidity
  • high density (compared to gases)

18
The Liquid State
  • Surface Tension
  • results in droplets when a liquid is poured onto
    a surface
  • depends on IMFs

19
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

20
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

21
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.

22
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

23
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

24
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

25
Types of Solids
  • three common unit cells exist
  • simple cubic
  • body centered cubic
  • face centered cubic

26
Types of Solids
  • Amorphous Solids
  • noncrystalline
  • glass is an example
  • disorder abounds

27
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

28
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

29
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)

30
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

31
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

32
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

33
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

34
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

35
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

36
Structure and Bonding in Metals
  • Hexagonal Closest Packed Structure (hcp)
  • aba arrangement
  • First Layer
  • each sphere is surrounded by six other spheres

37
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

38
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

39
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!)

40
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

41
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

42
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

43
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

44
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

45
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

46
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)

47
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

48
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

49
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

50
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

51
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

52
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

53
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

54
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

55
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

56
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

57
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

58
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

59
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

60
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

61
Covalent Network Solids
  • Carbon compounds usually contain long chains with
    C-C bonds
  • Silicon compounds usually contain chains with
    Si-O bonds

62
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

63
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

64
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

65
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

66
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

67
Semiconductors
  • Enhance conductivity of semiconductors by doping
    the crystal with other atoms

68
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

69
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

70
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

71
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

72
Types of Solids
  • Ionic Solids
  • between positive and negative ions
  • held by ionic bonds
  • electrostatic forces between oppositely charged
    ions

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