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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. Liquids and Solids. Liquids properties lie between those of solids and gases

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liquids and solids
Liquids and Solids
  • Gas
    • low density
    • high compressibility
    • completely fills its container
  • Solid
    • high density
    • only slightly compressible
    • rigid
    • maintains its shape
liquids and solids1
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 solids2
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
Intermolecular Forces
  • Bonds are formed between atoms to form molecules
    • intramolecular bonding (within the molecule)
intermolecular forces1
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 forces2
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
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 forces1
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
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 bonding1
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
London Dispersion Forces
  • aka Van der Waals forces
  • Nonpolar molecules must exert some kind of force or they would never solidify
london dispersion forces1
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 forces2
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 forces3
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 forces4
London Dispersion Forces
  • Element Freezing Point (oC)

Helium -269.7

Neon -248.6

Argon -189.4

Krypton -157.3

Xenon -111.9

london dispersion forces5
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
The Liquid State
  • Properties of liquids
    • low compressibility
    • lack of rigidity
    • high density (compared to gases)
the liquid state1
The Liquid State
  • Surface Tension
    • results in droplets when a liquid is poured onto a surface
    • depends on IMF’s
the liquid state2
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 state3
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 state4
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 state5
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 state6
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
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 solids1
Types of Solids
  • three common unit cells exist
    • simple cubic
    • body centered cubic
    • face centered cubic
types of solids2
Types of Solids
  • Amorphous Solids
    • noncrystalline
    • glass is an example
    • disorder abounds
types of solids3
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
types of solids4
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
types of solids5
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 solids6
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 solids7
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 solids8
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
Structure and Bonding inMetals
  • 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 metals1
Structure and Bonding inMetals
  • 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 metals2
Structure and Bonding inMetals
  • 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 metals3
Structure and Bonding inMetals
  • Hexagonal Closest Packed Structure (hcp)
    • aba arrangement
    • First Layer
      • each sphere is surrounded by six other spheres
structure and bonding in metals4
Structure and Bonding inMetals
  • 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 metals5
Structure and Bonding inMetals
  • 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 metals6
Structure and Bonding inMetals
  • 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 metals7
Structure and Bonding inMetals
  • 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 metals8
Structure and Bonding inMetals
  • 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
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 metals1
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 metals2
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 metals3
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 metals4
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 metals5
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 metals6
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 metals7
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 metals8
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
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 alloys1
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 alloys2
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 alloys3
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
  • 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 solids1
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 solids2
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 solids3
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 solids4
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 solids5
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 solids6
Covalent Network Solids
  • Carbon compounds usually contain long chains with C-C bonds
  • Silicon compounds usually contain chains with Si-O bonds
covalent network solids7
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 solids8
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 solids9
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 solids10
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
  • 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
  • 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)
    • 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 formedwith a missing valence electron
types of solids9
Types of Solids
  • Ionic Solids
    • between positive and negative ions
    • held by ionic bonds
      • electrostatic forces between oppositely charged ions