Chapter_11-IntrmolForces_Liquids_Solids-all

Chapter_11-IntrmolForces_Liquids_Solids-all - Chapter 11....

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Unformatted text preview: Chapter 11. Intermolecular forces, liquids, solids forces, Intermolecular forces, liquids and solids Chapter 11. Intermolecular forces, liquids, solids forces, Gases vs liquids vs solids Condensed phases Balance between kinetic energies of particles and interparticle attraction defines if substance is gas, liquid or solid. Chapter 11. Intermolecular forces, liquids, solids forces, Intermolecular forces 431 kJ/mol 16 kJ/mol Intermolecular forces are generally weaker than covalent or ionic bonds. Chapter 11. Intermolecular forces, liquids, solids forces, Intermolecular forces define boiling and melting points The stronger the intermolecular forces is the higher boiling point and higher melting point are (because it takes more kinetic energy to overcome intermolecular attraction). Boiling point = the temperature at which liquid vapor forms inside the liquid. Melting point = the temperature at which solid turns liquid. Crystalline solids have a very distinct melting point (salts, metals). Amorphous solids gradually soften before they completely turn into liquid (paraffin). Chapter 11. Intermolecular forces, liquids, solids forces, Major types of intermolecular forces • Ion-dipole • Dipole-dipole forces (polar substances) • London dispertion forces (all substances) • Hydrogen bonds (substances with hydrogen covalently bonded to N, O or F – highly electronegative small atoms) Chapter 11. Intermolecular forces, liquids, solids forces, Ion-dipole forces Ion-dipole forces are important in the solution of electrolytes in polar liquids. Dipoles of solvent orient toward a cation (or anion) by their oppositely charged ends and are attracted to it. Chapter 11. Intermolecular forces, liquids, solids forces, Dipole-dipole forces • Exist in polar substances. • Act on a short distances only. Although there are both attraction and repulsion between dipoles, the dipoles tend to orient in such way that attraction prevail (if temperature allows). Chapter 11. Intermolecular forces, liquids, solids forces, Dipole moment, molecular size and dipole-dipole forces • The larger dipole moment the stronger d.-d. forces (compare higher boiling points above). • The larger molecule the father away dipoles are the weaker role d.d. forces play. Chapter 11. Intermolecular forces, liquids, solids forces, London dispersion forces Fritz London (1930): motion of electrons in molecules can create instantaneous dipole moment, which can induce a dipole in the neighbor molecule making a temporary dipoledipole force. We know that dispersion forces exist, because we can observe non-polar substances, even noble gases, been liquefied and even solidified. Chapter 11. Intermolecular forces, liquids, solids forces, Features of London dispersion forces. 1. • Magnitude of dispersion force depends upon polarizability of a molecule (the ease with which dipole moment can be induced). • Larger molecules (or atoms) (i.e. with larger molecular weight) have electrons father from atomic nuclei, i.e. weaker attracted, thus their polarizability is higher dispersion forces are greater. Chapter 11. Intermolecular forces, liquids, solids forces, Features of London dispersion forces. 2. • Dispersion forces are stronger when number of contacts between molecules is bigger. Long molecule more contacts higher boiling point H3C-CH2-CH2-CH2-CH3 Compact molecule less contacts lower disp.force lower boiling point CH3 | H3C-C-CH3 | CH3 Chapter 11. Intermolecular forces, liquids, solids forces, Features of London dispersion forces summary • Dispersion forces are noticeable only at very short distances between molecules (similar to dipole-dipole). • Dispersion forces operate between all molecules, whether they are polar or non-polar. • When molecular weights and shapes are comparable, then the dipole-dipole forces decide which molecules have greater intermolecular forces. • When molecular weights are widely different, then the dispersion forces contribute more into intermolecular interaction than dipole-dipole. Chapter 11. Intermolecular forces, liquids, solids forces, Hydrogen bonds (H-bonds) H-bonds are formed between a hydrogen atom attached to a small electronegative atom and a lone pair on another nearby small electronegative atom. Most usually the small electronegative atoms above are N, O or F. Examples of H-bonding: Chapter 11. Intermolecular forces, liquids, solids forces, Features of hydrogen bonds. 1. Anomalously high boiling point of water is due to H-bonds (O⋅ ⋅ H) Boiling points H-bonds are generally stronger than dipole-dipole or dispersion forces, but weaker than covalent bonds. Dispersion forces Chapter 11. Intermolecular forces, liquids, solids forces, Classic example of hydrogen bonds – the water ice Strong H-bonds (O⋅ ⋅ H) hold the ice crystal structure very rigid with plenty of empty space inside – that is why ice is less dense than water (where molecules move randomly). This is not typical for substances which cannot form H-bonds. RasMol demo of water and ice. Chapter 11. Intermolecular forces, liquids, solids forces, Example: Importance of H-bonds for life: DNA H-bonds (O⋅ ⋅ H and N⋅ ⋅ ⋅ H) between nucleobases is one of the forces that support DNA double helix structure and allow AT and GC recognition. RasMol demo of B-DNA. Chapter 11. Intermolecular forces, liquids, solids forces, Intermolecular forces - summary • Effect of all forces (dipole-dipole, dispersion and Hbonds) is additive, i.e. presence of any of them increase intermolecular attraction. Energy range for different force types: • Dipole-dipole and dispersion forces (AKA Van-derWaals forces): 2 to 10 kJ/mol; • H-bonds: 5-25 kJ/mol (sometimes up to 100 kJ/mol); • Covalent bonds: 200-1100 kJ/mol. Chapter 11. Intermolecular forces, liquids, solids forces, Viscosity of liquids Definition: Viscosity is the resistance of a liquid to flow. •More viscous liquids flow slower. •More viscous liquids resist stronger to object motion. Viscosity units: 1 kg/(m⋅ s) = 10 P SI unit CGS unit - Poise Real gases have viscosity, too. It is lower than same of liquids. Chapter 11. Intermolecular forces, liquids, solids forces, Viscosity of liquids Viscosity is also called internal friction. The reason for viscosity is intermolecular attraction – layers that move slower retard the faster layers. Viscosity typically grows with molecular weight and decreases with temperature. Chapter 11. Intermolecular forces, liquids, solids forces, Surface tension of liquids Definition: surface tension is the energy required to increase the surface of a liquid by a unit of area. SI unit: J/m2 Molecules at the surface of the liquid predominantly are attracted toward inside the liquid. Thus they have some energy excess. For molecules inside the liquid there is no predominant direction of attraction. Chapter 11. Intermolecular forces, liquids, solids forces, Capillary effect and contact angle • Cohesive forces – intermolecular forces in a liquid. • Adhesive forces – intermolecular forces between liquid and a solid. Cohesive forces < adhesive forces Concave meniscus Cohesive forces > adhesive forces Convex meniscus Chapter 11. Intermolecular forces, liquids, solids forces, Effects of surface tension of liquids To minimize the surface energy a liquid tends to minimize its surface released water droplet becomes spherical. Chapter 11. Intermolecular forces, liquids, solids forces, Phase changes It always takes energy to melt or to vaporize a substance: Heat of vaporization Heat of fusion Fusion = melting Chapter 11. Intermolecular forces, liquids, solids forces, Phase changes (typically) Heat of fusion < heat of vaporization: because intermolecular forces still act in liquid, whereas much less noticeable in gas phase. Chapter 11. Intermolecular forces, liquids, solids forces, Heating curve During phase change the temperature does not change Chapter 11. Intermolecular forces, liquids, solids forces, Calculations using heat curve Example: warming the water ice to steam: ∆ Htotal = Cs,ice⋅ m⋅ ∆ T1 {warming the ice} + m⋅ ∆ Hfus {melting the ice} + Cs,water⋅ m⋅ ∆ T2 {warming the water} + m⋅ ∆ Hvap {vaporizing the water} + Cs,vapor⋅ m⋅ ∆ T3 {warming the steam} Sample exercise: How much heat is needed to convert 1.00 mol of ice at -25°C into steam at 125°C. Accept Cs for ice, water and steam = 2.03, 4.18 and 1.84 J/(g⋅ K). ∆ Hfus = 6.01 kJ/mol, ∆ Hvap = 40.67 kJ/mol. Answer: 56 kJ. Chapter 11. Intermolecular forces, liquids, solids forces, Critical temperature and pressure Often, instead of cooling a gas, we can compress it in order to liquefy. Higher T needs larger P. Definition: above a certain Tcrit no pressure can convert gas into liquid. Chapter 11. Intermolecular forces, liquids, solids forces, Vapor pressure Only fastest molecules can escape from liquid into gas phase. Partial pressure of vapor. Chapter 11. Intermolecular forces, liquids, solids forces, Vapor pressure vs temperature Liquid boils when its vapor pressure equal or more than external (e.g. atmospheric pressure). Example: on high altitudes water boils at lower than 100°C. More volatile substance (higher vapor pressure) Less volatile substance (lower vapor pressure) Chapter 11. Intermolecular forces, liquids, solids forces, Clausius-Clapeyron equation ln(P) = -∆ Hvap/(RT) + C The equation gives us a method to determine the enthalpy of vaporization from the slope of experimental ln(P) vs 1/T curves we can. Chapter 11. Intermolecular forces, liquids, solids forces, Phase diagram Pressure and temperature determine if substance exists as gas, liquid or solid. In triple point all three phases can coexist in equilibrium (usually we need a sealed and heat isolated container to observe this) Chapter 11. Intermolecular forces, liquids, solids forces, Features of a phase diagram •Triple point: temperature and pressure at which all temperature three phases are in equilibrium. •Vapor-pressure curve: generally as pressure generally increases, temperature required to boil increases. increases, •Critical point: critical temperature and pressure for critical the gas. (min. T when gas cannot be liq. by P) the •Melting point curve: as pressure increases, the as solid phase is favored if the solid is more dense than the liquid. than •Normal melting point: melting point at 1 atm. Chapter 11. Intermolecular forces, liquids, solids forces, Example: Phase diagrams of H2O and CO2 Chapter 11. Intermolecular forces, liquids, solids forces, Solids Crystalline Amorphous Quartz, SiO2 Well-defined Well-defined 3D arrangemen arrangemen t of atoms, of ions or molecules. Melting point. point. Fused silica, SiO2 No No orderly structure. No specific melting point. point. Chapter 11. Intermolecular forces, liquids, solids forces, Unit cell and crystal lattice Unit cell – a smallest building block of a solid crystal. It is a parallelepiped with certain dimensions and angles. and Coordination number of a particle = the number of other particles immediately surrounding it. surrounding Crystal is formed by arranging the unit cell repeatedly on the crystal lattice. crystal There are seven There seven basic types of unit cells. unit Chapter 11. Intermolecular forces, liquids, solids forces, Cubic cells – the most common type Primitive cubic. 1/8⋅ 8=1 atom per 8=1 cell. Coord.num=6. cell. Body-centered Body-centered cubic. 2 atoms per cell. Coord.num.= 8. Coord.num.= Face-centered cubic. Face-centered 4 atoms per cell. Coord.num.=12. Coord.num.=12. Chapter 11. Intermolecular forces, liquids, solids forces, NaCl crystal There are two equal There ways to define the lattice – on Na+ or on Cl- ions. on The crystal structure is facecentered cubic. 4Na+ and 4Cl- ions per cell (the cation-to-anion ratio must be the same as for whole crystal). crystal). Relative sizes Relative of ions in the lattice: lattice: Chapter 11. Intermolecular forces, liquids, solids forces, Sample exercise: calculating density from unit cell type and size: One of iron crystal forms has bodycentered cubic cell with a size 2.8664 Å. centered 24 (1 amu = 1.66⋅ 10--24 g) (1 g) Density of Fe = ? Chapter 11. Intermolecular forces, liquids, solids forces, Close packing of spheres Important for metals – all atoms are of the Important same size. Atoms are close each to other. same 74% of space 74% is occupied. is Coordination Coordination number = 12 in both struct. in Hexagonal close Hexagonal packing packing (ABAB layer sequence) Cubic close packing (ABCA layer sequence) (face-centered unit cell) Chapter 11. Intermolecular forces, liquids, solids forces, Close packing of spheres ABAB ABCA Chapter 11. Intermolecular forces, liquids, solids forces, Packing of very unequal spheres Example: some ionic Example: compounds with large anions and small cations. small Bigger ions assume a Bigger close-packed arrangement and smaller ions occupy spaces between them. between Chapter 11. Intermolecular forces, liquids, solids forces, Lattice planes Lattice planes – imaginary Lattice planes containing lattice points. Many lattice planes can be selected, but the most important contain crystal particles most closely spaced. spaced. Chapter 11. Intermolecular forces, liquids, solids forces, X-ray diffraction of crystals Waves diffraction is Waves most efficient when wavelength is close to size of a diffraction pattern. diffraction This is true for Xrays and crystals (220 Å) 20 Å) Diffraction pattern tells us about Diffraction crystal structure. crystal For X-rays For scattered in crystal the lattice planes play a role of diffraction grid. diffraction Chapter 11. Intermolecular forces, liquids, solids forces, Bonding in solids. Molecular solids Particles are held together by Particles intermolecular forces which are weak. weak. Molecular solids normally are soft and have low melting points (<200°C). have Shape of Shape molecule is important for good packing, for stronger solid. solid. H-bonds are H-bonds important for stronger solid. stronger Chapter 11. Intermolecular forces, liquids, solids forces, Covalent network solids Diamond. C-C bonds, Diamond. sp3-hybridization, sp -hybridization, tetrahedral geometry. tetrahedral Graphite. C-C bonds, Graphite. sp2-hybridization, trigonal planar geometry geometry Strong solids, high melting points (thousands Strong °C). °C) Ex.: SiO2, SiC, BN,… Particles Particles are held by covalent bonds. bonds. Chapter 11. Intermolecular forces, liquids, solids forces, NaCl Ionic solids Particles are held by Particles electrostatic attraction. electrostatic The larger the charge the The stronger the solid: stronger Tmelt(NaCl) = 801°C Tmelt(MgO) = 2852°C Chapter 11. Intermolecular forces, liquids, solids forces, Typical ionic crystal structures CsCl ZnS – Zinc Blende ZnS (mineral) (mineral) CaF2- fluorite fluorite (mineral) (mineral) Chapter 11. Intermolecular forces, liquids, solids forces, Metallic solids Valence electrons are Valence shared by all atoms of a metallic object. metallic The more conduction electron the metal have the weaker the bonds are. Typically, the more Typically, valence electrons the metal has, the stronger the bonds are. the Na, one val.electron: Tmelt=97.5°C Cr, six val.electrons: Tmelt=1890°C ...
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This note was uploaded on 01/02/2012 for the course CHEM 114 taught by Professor Sergeiaksyonov during the Fall '11 term at ASU.

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