Unit 9 – Intermolecular Forces

1. Unit 9 – Intermolecular Forces Liquids, Solids, and Phase Changes Chapters 11 and 12

2. Intermolecular Forces • Interaction of molecules with other molecules • Weaker forces than ionic or covalent bonds • Less energy is required to overcome the force of an intermolecular attraction than a chemical bond. • Examples of “breaking” intermolecular forces: • Melting Point • Boiling Point • The stronger the IMF, the higher the melting point/boiling point.

3. Types of Intermolecular Forces • Ion-Dipole Forces • Dipole-Dipole Forces • London Dispersion Forces • Hydrogen Bonding • All are considered electrostatic in nature (attractions between + and – species)

4. Ion-Dipole Forces • Exist between an ion and the partial charge on the end of a polar molecule • Remember that polar molecules are dipoles and have a dipole moment + - + - +

5. Dipole-Dipole Forces • The positive end of one polar molecule is attracted to the negative end on another polar molecule. • Effective only when polar molecules are close together and in correct orientation • Weaker than ion-dipole forces • Dipole-dipole interactions involve both attractive and repulsive forces • Molecules of equal mass and size exhibit stronger intermolecular attractions with increasing polarity

6. Hydrogen Bonding Attraction between a hydrogen atom in a polar bond and an unshaired electron pair on a nearby small electronegative ion or atom (generally F, O, or N) A strong intermolecular attraction that accounts for high m.p. andb.p., specific heat, and heat of vaporization Sometimes considered to be a dipole-dipole attraction Strong because protons are more massive in size Reason ice is less dense than water

7. H2O HF H2Te H2Se NH3 SbH3 H2S HI AsH3 HCl HBr PH3 SnH4 GeH4 SiH4 CH4 100 Boiling Points 0ºC -100 200

8. Water d+ d- d+

9. London Dispersion Forces Attraction based on electron location of one molecule to electron location on a different molecule Non - polar molecules also exert forces on each other – otherwise, there would be no solids or liquids. Instantaneous dipole moment can be induced then attraction of opposite dipoles More polarizable molecules (“squishy clouds”) have greater L-D forces This force increases with larger molecule size (thus larger molecular weights)

10. London Dispersion Forces • Weak, short-lived. • Lasts longer at low temperature. • Explains the trend for bigger molecules, higher melting and boiling points. • Much, much weaker than other forces. • Also called Van der Waal’s forces.

11. Example

12. Vibrations of the nucleus within the electron charge cloud creates a temporary dipole, shown here as delta + and delta - charges. This causes sympathetic vibrations of the nucleus in neighboring atoms, so that the delta + and delta - charges become arranged opposite one another, plus near minus. There is now attraction between the opposite delta charges on neighboring atoms, creating Van der Waal's forces. Eventually all of the atoms or molecules vibrate in unison, making the partial charges enough to hold the structure together. When enough energy is given to a substance held together by only Van der Waal's forces, they are overcome and the compound melts.

13. Represented above are five identical balloons, each filled to the same volume at 25°C and 1.0 atm pressure with the pure gas indicated. For each question, be sure to explain your reasoning. (a) Which balloon contains the greatest mass of gas? (b) Compare the average kinetic energies of the gas molecules in the balloons. (c) Which balloon contains the gas that would be expected to deviate most from the behavior of an ideal gas? (d) Twelve hours after being filled, all the balloons have decreased in size. Predict which balloon will be the smallest.

14. Comparing Intermolecular Forces Are polar molecules AND ions both present? Interacting molecules/ions Are Polar Molecules Involved? ON YES Are ions involved? YES Are H atoms bonded to O, N, or F? ON YES ON London Dispersion Ar, I2 Ion-dipole Forces KBrin H2O Ionic Bond NaCl YES ON Hydrogen Bonding H2O, NH3, HF Dipole-Dipole H2S, CH3Cl

15. Viscosity – the resistance of a liquid to flow. The greater the viscosity, the slower it flows Can be measured as a rate of flow through a metal tube or the rate at which an object falls through the liquid. Viscosity increases with an increase in molecular weight Large IMFs = more viscous liquid Large molecules can get tangled up. Cyclohexane has a lower viscosity than hexane because it is a circle- more compact. Properties of Liquids – Many of the properties are due to internal attraction of atoms.

16. Energy required to increase the surface area of a liquid by a unit amount Inward forces that must be overcome Atoms at surface are closer together Forces that bind similar molecules (such as H-bonding) are called cohesive forces Forces that bind a substance to a surface are called adhesive forces Surface Tension

17. Surface tension • Molecules in the middle are attracted in all directions. • Molecules at the the top are only pulled inside. • Minimizes surface area.

18. Water in a glass adheres to the glass because the adhesive force is larger than the cohesive forces of the H2O molecules Mercury is the opposite Example:

19. Liquids spontaneously rise in a narrow tube Occurs when a small tube is placed in H2O and the water rises into the tube Glass is polar. It attracts water molecules Capillary Action

20. Beading • If a polar substance is placed on a non-polar surface. • There are cohesive, • But no adhesive forces. • And Visa Versa

21. Every phase change is accompanied by a change in the energy of the system ∆Hfus (heat of fusion) – molecules of a solid free up, having more ability to move. ∆Hvap(heat of vaporization) – molecules of a liquid break free ∆Hvap is almost always larger in comparison to ∆Hfus ∆Hsub(heat of sublimation) – molecules transition from a solid to a gas ∆Hsub =∆Hfus + ∆Hvap Phase Changes

22. A graphical way to summarize the conditions under which equilibria exist between different states of matter Predicts the phase of a a substance at a given temperature and pressure Phase Diagrams Boiling Point Variance Change in melting point All three phases are in equilibrium at this temperature and pressure

23. Critical Temperature: The highest temperature at which a distinct liquid phase can form Critical Pressure: pressure required to bring about liquefaction at this critical temperature Nonpolar, low molar mass substances have lower values for critical temperature and pressure Interpreting a Phase Diagram

24. The pressure exerted by a vapor in the space above a liquid Weaker intermolecular forces = higher vapor pressure Dynamic Equilibrium: rates of changing from gas to liquid and liquid to gas is the same Volatile Liquids: substances with higher vapor pressure that allow liquids to evaporate to dryness (no dynamic equilibrium exists) Vapor Pressure

25. Liquids boil when the vapor pressure equals the external pressure acting on the surface of the liquid. Bubbles of vapor form within the interior of the liquid Lower pressure = lower boiling point Vapor pressure and temperature are related through The Clausius-ClapeyronEquation where C is a constant Boiling

26. Crystalline Solid – atoms, ions, or molecules are ordered in well-defined arrangements Have flat surfaces, or faces Definite angles Regular shapes – “structured” Amorphous Solid – solid whose particles have no orderly structure Lack of well-defined faces and shapes Mixtures of molecules that do not stack together well Generally composed of large, complicated molecules Structures of Solids

27. Crystalline solids – melts at a specific temperature Amorphous solids – soften over a wide range of temperatures.

28. 1. Identify the most important types of inter-particle (between “molecules”) forces present in the solids of each of the following substances a) BaSO4 b) H2S c) Xe d) CHCl3 e) NH3 Warm-Up

29. 2. Predict which substance from the following pairs would have the greater intermolecular forces. (you WILL need to draw the Lewis structures to answer correctly) a) CO2or OCS b) PF3or PF5 c) SO3or SO2 d) HF or HBr e) CH3OH or H2CO f) CH3CH2CH2NH2 or H2NCH2CH2NH2 Warm-Up, cont.