Intermolecular Forces

Intermolecular Forces Mr. Nelson AP Chem

Intermolecular Forces • Forces that exist between molecules • Also called Van der Waals forces • Compare to intramolecular forces, which are the forces that hold atoms together • Holds individual molecules together, but does not determine macro properties (aka boiling point)

Intermolecular Forces • Largely determines physical properties of solids and liquids • Intermolecular forces are weaker than intramolecular forces (chemical bonding) • Boiling points and melting points are an indication of the strength of intermolecular forces

Ion-Dipole forces • Interaction between an ion and a partial charge in a polar molecule • Not technically a van der Waals force • Positive ions are attracted to negative end of a dipole • Force depends on dipole moment of polar molecule and charge of the ion • Example: NaCl in Water

Dipole-dipole force • Attractive forces between polar molecules • Weaker than ion-dipole forces • Positive end of one molecule is near negative end of another • In liquids, molecules of equal size and mass have increasing intermolecular attractions with increasing polarity

London Dispersion Forces • London proposed that the motion of electrons in an atom or molecule can create an instantaneous dipole moment • Example: If both electrons on an He atom are on the same side of a nucleus, you have an instantaneous dipole moment

London Dispersion Forces • A temporary dipole on one atom can induce a similar dipole on an adjacent atom (called an induced dipole)

London Dispersion Forces • Because of induced dipoles, molecules attract one another • Dispersion forces exist in all types of molecules (polar or non-polar, charged or not) • In many cases, dispersion forces can be stronger than dipole-dipole • Example: CH3F (B.P. -78.9 °C) vs. CCl4 (B.P. 76.5 °C)

London Dispersion Forces • Dispersion forces increase with molar mass of the molecule • More electrons allows for greater chances of dispersion • When comparing molecules of similar weights and shapes, dipole-dipole forces tend to be the decisive factor • When comparing differing weights, dispersion forces is the decisive factor

Intermolecular Forces Recap • Example • Of Br2, Ne, HCl, HBr, and N2, which has: • The largest intermolecular dispersion forces? • The largest dipole-dipole attractive forces?

Hydrogen Bonding • A special type of dipole-dipole attraction that exists between hydrogen atoms in a polar bond and small electronegative ions • Can only form between hydrogen and N, O, or F

Hydrogen Bonding • Water has a lower molar mass and a higher boiling point • Dispersion forces do not account for this

Hydrogen Bonding • Explains density of ice • Solid is less dense than liquid (not common)

Review of IM forces

Phase Changes • Changes in state of matter • In general, each state of matter can change into either of the two other states • Phase changes require energy • When becoming a more disordered state, requires energy to overcome intermolecular forces that hold them together

Phase Changes

Phase Changes • Melting, vaporization, and sublimation are all endothermic processes • Freezing, condensation, and deposition are all exothermic processes

Phase Changes • Fusion • Melting of a solid • Molar heat of fusion or enthalpy of fusion (∆Hfus) • ∆Hfusion= -∆Hsolidification • Vaporization • Vaporizing of a liquid • Molar heat of vaporization or enthalpy of vaporization (∆Hvap) • ∆Hvaporization= -∆Hcondensation

Heating Curves • Graph of temperature of the system versus heat added to the system • Some segments of the graph are heating as single phase others are converting one phase to another • Heat change when temp. inc. is given by q=mC∆T. During phase change, it is q=n∆H(n is moles of substance)

Heating Curve of Water

Phase Change Problem • Freon-11 (CCl3F) has a normal boiling point of 23.8 °C. The specific heat of CCl3F (l) is 0.87 J/g K and CCl3F (g) is 0.59 J/g K. The heat of vaporization is 24.75 kJ/mole. Calculate the heat required to convert 10.0 g of Freon-11 from liquid at -50.0 °C to gas at 50.0 °C.

Vapor Pressure • In a closed system, liquid will initially evaporate and then condense back into its liquid form • The pressure exerted by this liquid/gas equilibrium is called vaporpressure • When evaporation occurs in an open system, no equilibrium can be established

Vapor Pressure • Substances with high vapor pressure evaporate more easily than those with a low vapor pressure • Volatile liquids (aka liquids that evaporate easily) have high vapor pressures • Water can vaporize at room temp. because molecules in liquid move at different speeds and some can overcome IM forces (think space shuttles/escape velocity)

Vapor Pressure • As the number of gas molecules increases in a closed system, the probability that gas molecules will strike the liquid increases. • This allows the molecules to become “trapped” by the IM forces of the liquid • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/vaporv3.swf

Temperature & Vapor Pressure • Vapor pressure increases nonlinearly with temperature • More molecules have enough KE to escape into gas phase

Vapor Pressure & Boiling Point • Boiling point is when the vapor pressure equals the external pressure acting on the surface • A “normal” boiling point is measured at 1 atm • A more volatile liquid has a lower boiling point and a higher vapor pressure (and vice versa)

Phase Diagrams • A graph that displays the conditions under which an equilibrium exists between different states of matter • Allows prediction of a substance’s state of matter at a given temperature and pressure

Phase Diagrams • Critical Temperature (Tc) is the highest temperature a substance can exist as a liquid • Meaning above Tc molecular motion breaks IM attraction (Tc measures strength of IM forces) • Critical Pressure (Pc) is pressure required to bring about liquefaction at the critical temperature

Phase Diagrams • Triple Point is the place at which all three phases exist at equilibrium

Intermolecular Forces