Chapter 11 Intermolecular Forces and Liquids & Solids
The Kinetic Molecular Theory of Liquids & Solids Gases vs. Liquids & Solids difference is distance between molecules Liquids Molecules close together; very little empty space so difficult to compress, denser than gases Molecules held together by attractive forces Have a definite volume, yet can flow, be poured, & assume shape of container Solids Molecules held rigidly in position, virtually no freedom of motion Molecules typically arranged in regular configurations in 3-D Almost incompressible & have definite shape & volume With few exceptions, density of solid > density of liquid form for same substance Phase = homogeneous part of the system in contact with other parts of the system but separated by a well-defined boundary (i.e. Ice water) For us, phase will be used when discussing changes of state involving one substance, as well as systems containing >1 phase of a substance Table 1.1
2 Phases Solid phase - ice Liquid phase - water
Intermolecular Forces Intermolecular Forces = attractive forces between molecules Intramolecular Forces hold atoms together in a molecule Intermolecular vs. Intramolecular 41 kj to vaporize 1 mole of water (inter) 930 kj to break all O-H bonds in 1 mole of water (intra) Generally, intermolecular forces are much weaker than intramolecular forces. Measure of intermolecular force boiling point melting point DH vap DH fus DH sub
Intermolecular Forces Dipole-Dipole Forces Attractive forces between polar molecules Orientation of Polar Molecules in a Solid Molecules that have a permanent dipole moment tend to align with opposite polarities in the solid phase for maximum attractive interaction.
Intermolecular Forces Ion-Dipole Forces Attractive forces between an ion and a polar molecule Ion-Dipole Interaction
Hydration The strength of this interaction depends on the charge & size of the ion and on the magnitude of the dipole moment & size of the molecule. The charges on a cation are generally more concentrated, because cations are usually smaller than anions. Therefore, a cation interacts more strongly with dipoles than does an anion having a charge of the same magnitude.
Intermolecular Forces Dispersion Forces Occur via INDUCED dipoles in which the separation of + & - charges in the atom or nonpolar molecule is due to the proximity of an ion or polar molecule. Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules ion-induced dipole interaction dipole-induced dipole interaction
Induced Dipoles Interacting With Each Other
Dispersion Forces Intermolecular Forces Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted. Polarizability increases with: greater number of electrons more diffuse electron cloud (electrons not held tightly by nucleus) Dispersion forces usually increase with molar mass. Melting point as # of electrons
Dispersion Forces Intermolecular Forces Instantaneous dipole = lasts for a fraction of a second; at any moment is created by specific positions of electrons Averaged over time, instantaneous dipoles cancel each other resulting in no net dipole moment Dispersion forces = attractive forces that arise as a result of temporary dipoles induced in atoms or molecules Also called London forces Usually increase with molar mass because molecules with larger molar mass tend to have more electrons & electrons are held less tightly by nucleus Increase in strength with the number of electrons Comparable to & sometimes > dipole-dipole forces between molecules Exist between ALL molecules
What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH 4 CH 4 is nonpolar: dispersion forces. S SO 2 SO 2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO 2 molecules.
Hydrogen Bond Intermolecular Forces The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H B or A H A A & B are N, O, or F
Hydrogen Bond Average strength of a hydrogen bond is large for a dipole-dipole interaction (up to 40 kj/mol) Strength is determined by coulombic interaction between lone-pair electrons of the electronegative atom & the hydrogen nucleus.
Why is the hydrogen bond considered a special dipole-dipole interaction? Decreasing molar mass Decreasing boiling point
Properties of Liquids Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. Strong intermolecular forces = High surface tension
Properties of Liquids Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules Capillary action: Adhesion Cohesion
Properties of Liquids Viscosity is a measure of a fluid s resistance to flow. The greater the viscosity, the more slowly the liquid flows. Viscosity usually decreases as temperature increases. Strong intermolecular forces = Glycerol is most viscous liquid listed. High viscosity It can form hydrogen bonds, like water, has 3 -OH groups that participate in H-bonding, & because of their shape, molecules of glycerol tend to become entangled rather than slipping past one another. CH 2 -OH CH 2 -OH CH 2 -OH
Structure & Properties of Water Maximum Density 4 0 C Density of Water Ice is less dense than water
Water is a unique substance Has high specific heat (a lot of heat is absorbed with only a slight change in temp. & can give off much heat with small decrease in temperature) Has many intermolecular hydrogen bonds Each oxygen can be bonded to 4 hydrogen atoms: two covalently within the water molecule & two hydrogen bonds to adjacent molecules via 2 lone pairs on O. Results in 3-dimensional structure Solid form is less dense than liquid form Two processes act in opposite directions the trapping of free water molecules in cavities of ice structure & thermal expansion
Crystal Structure A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions. An amorphous solid, such as glass, does not possess a welldefined arrangement and long-range molecular order. A unit cell is the basic repeating structural unit of a crystalline solid. lattice point At lattice points: Atoms Molecules Ions Unit Cell Unit cells in 3 dimensions
Coordination # is the # of atoms (or ions) surrounding an atom (or ion) in a crystal lattice. For x above, the coordination # = 6 (x has 6 immediate neighbors) The larger the coordination #, the closer the spheres are to each other
Coordination # = 6 8 12
Shared by 8 unit cells Shared by 2 unit cells
1 atom/unit cell (8 x 1/8 = 1) 2 atoms/unit cell (8 x 1/8 + 1 = 2) 4 atoms/unit cell (8 x 1/8 + 6 x 1/2 = 4)
Closest Packing The most efficient arrangement of spheres, coordination # = 12 Hexagonal closepacked = ABA (ABABABABA ) Cubic close-packed = ABC = face-centered cubic (ABCABCABCABC)
When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 409 pm. Calculate the density of silver. d = m V V = a 3 = (409 pm) 3 = 6.83 x 10-23 cm 3 4 atoms/unit cell in a face-centered cubic cell m = 4 Ag atoms 107.9 g x mole Ag x 1 mole Ag 6.022 x 10 23 atoms = 7.17 x 10-22 g d = m V 7.17 x 10-22 g = = 10.5 g/cm 3 6.83 x 10-23 cm 3
X-ray Diffraction by Crystals X-ray diffraction refers to the scattering of X rays by the units of a crystalline solid The scattering, or diffraction, patterns produced are used to determine the arrangement of particles in the solid lattice. Refer to section 10.6 for discussion of interference of waves X-rays exhibit interference behavior; an X-ray diffraction pattern is the result of interference of waves associated with X rays The wavelength of X rays is comparable in magnitude to the distances between lattice points in a crystal (von Laue) Figures 11.23, 11.24 X-ray diffraction is most accurate method for determining bond lengths, bond angles in molecules in the solid state
Extra distance = BC + CD = 2d sinq = nl (Bragg Equation)
X rays of wavelength 0.154 nm are diffracted from a crystal at an angle of 14.17 0. Assuming that n = 1, what is the distance (in pm) between layers in the crystal? nl = 2d sin q n = 1 q = 14.17 0 l = 0.154 nm = 154 pm d = nl 2sinq = 1 x 154 pm 2 x sin(14.17) = 314.0 pm
Types of Crystals 1. Ionic Crystals Lattice points occupied by cations and anions Held together by electrostatic attraction Hard, brittle, high melting point Poor conductor of heat and electricity CsCl ZnS CaF 2
Cell edge length = 2(r Na+ + r Cl- ) Ionic Crystals: NaCl
Types of Crystals 2. Covalent Crystals Lattice points occupied by atoms Held together by covalent bonds Hard, high melting point Poor conductor of heat and electricity carbon atoms diamond graphite
Types of Crystals 3. Molecular Crystals Lattice points occupied by molecules Held together by intermolecular forces Soft, low melting point Poor conductor of heat and electricity SO 2, Ice, I 2, P 4, S 8 are examples of molecular crystals. S
Types of Crystals Metallic Crystals Lattice points occupied by metal atoms Held together by metallic bonds Soft to hard, low to high melting point Good conductors of heat and electricity nucleus & inner shell e - Cross Section of a Metallic Crystal mobile sea of e -
Crystal Structures of Metals
Types of Crystals
Chemistry In Action: High-Temperature Superconductors MgB 2
Amorphous Solids An amorphous solid does not possess a well-defined arrangement and long-range molecular order. A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing Crystalline quartz (SiO 2 ) Non-crystalline quartz glass
Phase Changes Phase changes = transformations from one phase to another & occur when energy (usually in the form of heat) is added or removed. Physical changes characterized by changes in molecular order Molecules in solid phase have greatest order; those in the gas phase have the greatest randomness or least order
Evaporation Condensation Least Order T 2 > T 1 Greatest Order
Liquid-Vapor Equilibrium Evaporation or Vaporization = the process in which a liquid is transformed into a gas Condensation = the change from the gas phase to the liquid phase, often because a molecule strikes the liquid surface and becomes trapped by intermolecular forces in the liquid. The rate of evaporation is constant at any given temperature The rate of condensation increases with increasing concentration of molecules in the vapor (gas) phase Dynamic Equilibrium = the rate of a forward process is exactly balanced by the rate of the reverse process; is reached when the rates of evaporation & condensation become equal.
Vapor Pressure The equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation Maximum vapor pressure of a liquid at a given temp. Constant at constant temp. Independent of amount of liquid as long as some liquid is present H 2 O (l) H 2 O (g) Dynamic Equilibrium Rate of condensation = Rate of evaporation
Before Evaporation At Equilibrium
Molar Heat of Vaporization & Boiling Point Molar heat of vaporization (DH vap ) is the energy required to vaporize 1 mole of a liquid at its boiling point. Clausius-Clapeyron Equation ln P = - DH vap RT P = (equilibrium) vapor pressure + C ln Vapor Pressure Versus Temperature P P DH R R = gas constant (8.314 J/K mol) T1 T ( T T 1 vap 2 1 2 2 ) T = temperature (K)
The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure. The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm.
Critical Temperature & Pressure The critical temperature (T c ) is the temperature above which the gas cannot be made to liquefy, no matter how great the applied pressure. (highest temp. at which a substance can exist as a liquid). The critical pressure (P c ) is the minimum pressure that must be applied to bring about liquefaction at the critical temperature.
The Critical Phenomenon of SF 6 T < T c T > T c T ~ T c T < T c
Melting Freezing Liquid-Solid Equilibrium H 2 O (s) H 2 O (l) The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium
Molar heat of fusion (DH fus ) is the energy required to melt 1 mole of a solid substance at its freezing point.
Heating Curve Supercooling is a situation in which a liquid can be temporarily cooled to below its freezing point
Sublimation Deposition Solid-Vapor Equilibrium H 2 O (s) H 2 O (g) Molar heat of sublimation (DH sub ) is the energy required to sublime 1 mole of a solid. DH sub = DH fus + DH vap ( Hess s Law)
Phase Diagrams A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas. Phase Diagram of Water Each region represents a pure phase The line separating any 2 regions indicates conditions under which these 2 phases can exist in equilibrium The triple point is the point at which all 3 curves meet. It is the ONLY condition under which all 3 phases can be in equilibrium with one another
Effect of Increase in Pressure on the Melting Point of Ice and the Boiling Point of Water
Phase Diagram of Carbon Dioxide Positive slope At 1 atm CO 2 (s) CO 2 (g)