CHAPTER 10

The forces with which molecules attract each other. Intermolecular forces are weaker than ionic or covalent bonds. Intermolecular forces are responsible for the physical state of a compound (solid, liquid or gas). TYPES OF INTERMOLECULAR FORCES Dipole Interactions (between polar molecules) London Dispersion Forces (between all molecules but mainly force between nonpolar molecules and noble gases) Hydrogen Bonds (between molecules where hydrogen is bonded to nitrogen, oxygen and fluorine) INTERMOLECULAR FORCES

Dipole-dipole forces exist between neutral polar molecules. Polar molecules need to be close together. Weaker than ion-dipole forces. There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity. DIPOLE-DIPOLE FORCES

Weakest of all intermolecular forces. It is possible for two adjacent neutral molecules to affect each other. The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom).For an instant, the electron clouds become distorted.In that instant a dipole is formed (called an instantaneous dipole). Polarizability is the ease with which an electron cloud can be deformed. The larger the molecule (the greater the number of electrons) the more polarizable. London dispersion forces increase as molecular weight increases. London dispersion forces exist between all molecules. London dispersion forces depend on the shape of the molecule. The greater the surface area available for contact, the greater the dispersion forces. LONDON DISPERSION FORCES

H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N). –Electrons in the H-X (X = electronegative element) lie much closer to X than H. –H has only one electron, so in the H-X bond, the  + H presents an almost bare proton to the  - X. –Therefore, H-bonds are strong. Special case of dipole-dipole forces. By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. Intermolecular forces are abnormally strong. HYDROGEN BONDING

Hydrogen bonds are responsible for: –Ice Floating Solids are usually more closely packed than liquids; Therefore, solids are more dense than liquids. Ice is ordered with an open structure to optimize H-bonding. Therefore, ice is less dense than water. In water the H-O bond length is 1.0 Å. The O…H hydrogen bond length is 1.8 Å. Ice has waters arranged in an open, regular hexagon. Each  + H points towards a lone pair on O.

EFFECTS OF INTERMOLECULAR FORCES ON PHYSICAL PROPERTIES

Viscosity Viscosity is the resistance of a liquid to flow. A liquid flows by sliding molecules over each other. The stronger the intermolecular forces, the higher the viscosity. Surface Tension Bulk molecules (those in the liquid) are equally attracted to their neighbors. VISCOSITY & SURFACE TENSION

Surface molecules are only attracted inwards towards the bulk molecules. –Therefore, surface molecules are packed more closely than bulk molecules. Surface tension is the amount of energy required to increase the surface area of a liquid. Cohesive forces bind molecules to each other. Adhesive forces bind molecules to a surface.

The stronger the intermolecular forces, the higher the boiling point. For compounds with approximately the same molecular weight: BOILING POINT

HYDROGEN BONDED TO GROUP 16 ELEMENTS: NOTICE THAT H2O HAS A GREATER BOILING POINT THAN THE OTHER EVEN THOUGH IT HAS THE LOWEST MOLECULAR MASS DUE TO THE HYDROGEN BONDING. THE OTHER HYDROGEN COMPOUNDS EXPERIENCE DIPOLE-DIPOLE BONDING AND THE BOILING POINT INCREASES WITH INCREASING MOLECULAR MASS. HYDROGEN BONDED TO GROUP 14 ELEMENTS: NOTICE THAT THESE MOLECULES WITH HYDROGEN ARE EXPERIENCE LONDON DISPERSION FORCES AND ALL BOILING POINTS INCREASE WITH INCREASE MOLECULAR MASS BOILING POINT TRENDS

A. Evaporation and Vapor Pressure Vaporization or evaporation Endothermic

MOLECULES AT THE SURFACE HAVE LESS INTERMOLECULAR FORCES ON THEM THAN THE MOLECULES BELOW THEM. THUS THEY CAN ESCAPE THE LIQUID PHASE EASIER EVAPORATION

Explaining Vapor Pressure on the Molecular Level Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid. These molecules move into the gas phase. As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid. After some time the pressure of the gas will be constant at the vapor pressure.

A. Evaporation and Vapor Pressure Amount of liquid first decreases then becomes constant Condensation - process by which vapor molecules convert to a liquid When no further change is visible the opposing processes balance each other - equilibrium Vapor Pressure

A. Evaporation and Vapor Pressure Vapor pressure - pressure of the vapor present at equilibrium with its liquid Vapor Pressure –Vapor pressures vary widely - relates to intermolecular forces

BOILING OCCURS WHEN THE VAPOR PRESSURE OF THE LIQUID EQUALS THE ATMOSPHERIC PRESSURE. LESS PRESSURE MEANS THAT THE MOLECULES CAN ESCAPE EASIER AND THUS HAS A LOWER BOILING POINT Boiling Point and Atmospheric Pressure

Vapor Pressure and Boiling Point Liquids boil when the external pressure equals the vapor pressure. Temperature of boiling point increases as pressure increases. Two ways to get a liquid to boil: increase temperature or decrease pressure. –Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required. Normal boiling point is the boiling point at 760 mmHg (1 atm).

The melting point is the temperature at which a solid is converted to its liquid phase. In melting, energy is needed to overcome the attractive forces in the more ordered crystalline solid. The stronger the intermolecular forces, the higher the melting point. Because ionic compounds are held together by extremely strong interactions, they have very high melting points. With covalent molecules, the melting point depends upon the identity of the intermolecular force. For compounds of approximately the same molecular weight: MELTING POINT

C. Energy Requirements for the Changes of State Molar heat of fusion – energy required to melt 1 mol of a substance Molar heat of vaporization – energy required to change 1 mol of a liquid to its vapor

Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization: –it takes more energy to completely separate molecules, than partially separate them.

CALCULATING HEAT OF VAPORIZATION & VAPOR PRESSURE OF WATER Ln(P vap ) =[ (-ΔH vap /R)(1/T)] Textbook : pg 486

TO FIND THE BOILING POINT AT A DIFFERENT PRESSURE Ln (P1/P2) = (ΔH/R)[ (1/T2) –(1/T1)] or Ln (P1/P2) = (ΔH/R)[ (1/T2) –(1/T1)] Ln (P2/P1) = (ΔH/R)[ (1/T1) –(1/T2)] ΔH is ΔH vap

HEATING AND COOLING CURVES

Energy Changes Accompanying Phase Changes All phase changes are possible under the right conditions. The sequence heat solid  melt  heat liquid  boil  heat gas is endothermic. The sequence cool gas  condense  cool liquid  freeze  cool solid is exothermic.

Plot of temperature change versus heat added is a heating curve. During a phase change, adding heat causes no temperature change. –These points are used to calculate  H fus and  H vap. Supercooling: When a liquid is cooled below its melting point and it still remains a liquid. Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.

PHASE DIAGRAMS

Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases. Given a temperature and pressure, phase diagrams tell us which phase will exist. Any temperature and pressure combination not on a curve represents a single phase. PHASE DIAGRAMS

Features of a phase diagram: –Triple point: temperature and pressure at which all three phases are in equilibrium. –Vapor-pressure curve: generally as pressure increases, temperature increases. –Critical point: critical temperature and pressure for the gas. –Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid. –Normal melting point: melting point at 1 atm.

Critical Temperature and Pressure Gases liquefied by increasing pressure at some temperature. Critical temperature: the minimum temperature for liquefaction of a gas using pressure. Critical pressure: pressure required for liquefaction.

Water: –The melting point curve slopes to the left because ice is less dense than water. –Triple point occurs at  C and 4.58 mmHg. –Normal melting (freezing) point is 0  C. –Normal boiling point is 100  C. –Critical point is 374  C and 218 atm. Carbon Dioxide: –Triple point occurs at  C and 5.11 atm. –Normal sublimation point is  C. (At 1 atm CO 2 sublimes it does not melt.) –Critical point occurs at 31.1  C and 73 atm.