© 2012 Pearson Education, Inc. Chapter 11 Liquids and Intermolecular Forces John D. Bookstaver St. Charles Community College Cottleville, MO Lecture Presentation

© 2012 Pearson Education, Inc. Helpful links: Bozeman Solids and liquids Intermolecular forces Dipole forces London dispersion forces Intermolecular potential energy Water: A polar molecule Helpful links: Crash course Chemistry Partial pressure and Vapor pressure Liquids

Intermolecular Forces © 2012 Pearson Education, Inc States of Matter The fundamental difference between states of matter is the distance between particles. Gases: molecules far apart, random straight line translational motion, etc. Liquids: molecules close together, slide past one another Translational and vibrational motion Solids: molecules tightly packed in regular geometrical pattern. Vibrational motion only

Intermolecular Forces © 2012 Pearson Education, Inc. The States of Matter The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: –The kinetic (speed) energy of the particles. –The strength of the attractions between the particles. The properties follow from the kinetic molecular theory: arrangement and motion of molecules

Intermolecular Forces © 2012 Pearson Education, Inc. Ex: At the same temperature (average kinetic energy) CO 2 is a gas, while H 2 O is a liquid. Water molecules have “strong” intermolecular forces between their molecules. Increasing the kinetic energy will allow water to vaporize. Likewise, decreasing the kinetic energy can allow CO 2 to solidify, as weak IM forces develop between their molecules.

Intermolecular Forces © 2012 Pearson Education, Inc Intermolecular Forces The attractions between molecules are not as strong as the intramolecular bonds that hold compounds together. These intermolecular attractions are, however, strong enough to control physical properties, such as boiling and melting points, vapor pressures, and viscosities. There are three types of intermolecular Van Der Waal’s forces: dipole attractions, dispersion forces and hydrogen bonds.

Intermolecular Forces © 2012 Pearson Education, Inc. Dipole–Dipole Interactions Molecules that have permanent dipoles are attracted to each other. –The positive end of one is attracted to the negative end of the other, and vice versa. –These forces are only important when the molecules are close to each other.

Intermolecular Forces © 2012 Pearson Education, Inc. Dipole–Dipole Interactions The more polar the molecule, the higher its boiling point. Recall that polarity is expressed in “Debyes”?

Intermolecular Forces © 2012 Pearson Education, Inc. London Dispersion Forces Random distribution of electrons creates molecules with temporary dipoles even in nonpolar molecules Notice that the electrons on the same side make this He atom temporarily polar

Intermolecular Forces © 2012 Pearson Education, Inc. London Dispersion Forces These electrons in molecule B repel the electrons in neighboring molecule A inducing a dipole in A.

Intermolecular Forces © 2012 Pearson Education, Inc. London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

Intermolecular Forces © 2012 Pearson Education, Inc. London Dispersion Forces Important: these forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability. The greater the number of electrons, the more polarizable the molecule I call these “big and sticky” forces because they get stronger as molecules get larger (more electrons)

Intermolecular Forces © 2012 Pearson Education, Inc. Factors Affecting London Forces Shape of the molecule affects the strength of dispersion forces: long, skinny molecules have stronger dispersion forces than short, fat ones due to increased surface contact

Intermolecular Forces © 2012 Pearson Education, Inc. Factors Affecting London Forces The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms have larger electron clouds that are easier to polarize. Did I mention, I call these “big and sticky” forces because they get stronger as molecules get larger (more electrons)?

Intermolecular Forces © 2012 Pearson Education, Inc. Since these are all nonpolar molecules, dispersion forces only can account for the strength of intermolecular attraction. How do the number of electrons (and polarizability) compare for these molecules? BP: CH 4 (smallest) < CCl 4 (bigger) < CBr 4 (largest)

Intermolecular Forces © 2012 Pearson Education, Inc. It depends: >If two molecules are of comparable size and shape, dipole–dipole interactions will likely be the dominating force. >If one molecule is much larger than another, dispersion forces will likely determine its physical properties. Which Has a Greater Effect? Dipole–Dipole Interactions or Dispersion Forces

Intermolecular Forces © 2012 Pearson Education, Inc. Can you explain this? The nonpolar series (SnH 4 to CH 4 ) follow the expected trend in dispersion forces. The polar series follow the trend until you get to the smallest molecules in each group. This is due to a special force called hydrogen “bonding”

Intermolecular Forces © 2012 Pearson Education, Inc. Hydrogen Bonding The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong. We call these interactions hydrogen bonds. Since the attraction is physical and not chemical in nature, we should call them hydrogen “ bonds ” “

Intermolecular Forces © 2012 Pearson Education, Inc. Hydrogen Bonding Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed. Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. We call hydrogen atom a “naked” proton since its one electron is pulled away. This makes the hydrogen an unshielded positive charge. “

Intermolecular Forces © 2012 Pearson Education, Inc. SAMPLE EXERCISE 11.1 Identifying Substances that Can Form Hydrogen Bonds In which of the following substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH 4 ), hydrazine (H 2 NNH 2 ), methyl fluoride (CH 3 F), or hydrogen sulfide (H 2 S)? Solution Analyze: We are given the chemical formulas of four substances and asked to predict whether they can participate in hydrogen bonding. All of these compounds contain H, but hydrogen bonding usually occurs only when the hydrogen is covalently bonded to N, O, or F. Plan: We can analyze each formula to see if it contains N, O, or F directly bonded to H. There also needs to be an unshared pair of electrons on an electronegative atom (usually N, O, or F) in a nearby molecule, which can be revealed by drawing the Lewis structure for the molecule.

Intermolecular Forces © 2012 Pearson Education, Inc. Check: While we can generally identify substances that participate in hydrogen bonding based on their containing N, O, or F covalently bonded to H, drawing the Lewis structure for the interaction, as shown above, provides a way to check the prediction. Solution Solve: The criteria listed above eliminate CH 4 and H 2 S, which do not contain H bonded to N, O, or F. They also eliminate CH 3 F, whose Lewis structure shows a central C atom surrounded by three H atoms and an F atom. (Carbon always forms four bonds, whereas hydrogen and fluorine form one each.) Because the molecule contains a C––F bond and not an H––F bond, it does not form hydrogen bonds. In H 2 NNH 2, however, we find N––H bonds. If we draw the Lewis structure for the molecule, we see that there is a nonbonding pair of electrons on each N atom. Therefore, hydrogen bonds can exist between the molecules as depicted below.

Intermolecular Forces © 2012 Pearson Education, Inc. Answer: HOOH PRACTICE EXERCISE In which of the following substances is significant hydrogen bonding possible: methylene chloride (CH 2 Cl 2 ) phosphine (PH 3 ) hydrogen peroxide (HOOH), or acetone (CH 3 COCH 3 )?

Intermolecular Forces © 2012 Pearson Education, Inc. Ion–Dipole Interactions Ion–dipole interactions (a fourth type of force) are important in solutions of ions. The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents. (eg. Salt dissolves in water)

Intermolecular Forces © 2012 Pearson Education, Inc. The molecular compound CH 3 OH will dissolve and will develop H-bonds between the OH group and neighboring water molecules. The Ionic Ca(NO 3 ) 2 will dissociate into Ca +2 ions and NO 3 - ions which will be attracted to the polar water molecules.

Intermolecular Forces © 2012 Pearson Education, Inc. Summarizing Intermolecular Forces

Intermolecular Forces © 2012 Pearson Education, Inc. SAMPLE EXERCISE 11.2 Predicting the Types and Relative Strengths of Intermolecular Forces List the substances BaCl 2, H 2, CO, HF, and Ne in order of increasing boiling points. Check: The actual normal boiling points are H 2 (20 K), Ne (27 K), CO (83 K), HF (293 K), and BaCl 2 (1813 K), in agreement with our predictions. Solution Analyze: We need to relate the properties of the listed substances to boiling point. Plan: The boiling point depends in part on the attractive forces in the liquid. We need to order these according to the relative strengths of the different kinds of forces. Solve: BaCl 2 is ionic vs. others are molecular. Ionic is strong, continuous chemical bond which will be much stronger than any intermolecular force, so BaCl 2 will have highest BP and would in fact prob be a solid. Next, notice HF has H bonded to F so it will likely exhibit hydrogen bonding, considered the strongest of the IM forces. If an extremely large molecule presents itself, the question would get more complex since dispersion forces can be significant in larger molecules. Since all molecules are relatively small (low molecular weight) we can exclude this as a problem. From the remaining list, CO is polar so will exhibit dipole – dipole attraction; compared with the other two which are both nonpolar. Of the remaining two Ne has a larger mass than H 2, so it will likely have slightly stronger dispersion forces and be next in line. H 2, the smallest nonpolar molecule should have the weakest IM forces.

Intermolecular Forces © 2012 Pearson Education, Inc. PRACTICE EXERCISE (a) Identify the intermolecular forces present in the following substances, and (b) select the substance with the highest boiling point: CH 3 CH 3, CH 3 OH, and CH 3 CH 2 OH. Answers: (a) CH 3 CH 3 has only dispersion forces, whereas the other two substances have both dispersion forces and hydrogen bonds; (b) CH 3 CH 2 OH

Intermolecular Forces © 2012 Pearson Education, Inc Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.

Intermolecular Forces © 2012 Pearson Education, Inc. Viscosity Resistance of a liquid to flow is called viscosity. It is related to the ease with which molecules can move past each other. Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

Intermolecular Forces © 2012 Pearson Education, Inc. Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

Intermolecular Forces © 2012 Pearson Education, Inc Phase Changes

Intermolecular Forces © 2012 Pearson Education, Inc. Energy Changes Associated with Changes of State The heat of fusion is the energy required to change a solid at its melting point to a liquid.

Intermolecular Forces © 2012 Pearson Education, Inc. Energy Changes Associated with Changes of State The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas.

Intermolecular Forces © 2012 Pearson Education, Inc. Energy Changes Associated with Changes of State The heat of sublimation is defined as the energy required to change a solid directly to a gas.

Intermolecular Forces © 2012 Pearson Education, Inc. Notice, sublimation requires both melting (fusion) and boiling (vaporizing)

Intermolecular Forces © 2012 Pearson Education, Inc. Energy Revisited: Kinetic energy – energy of motion molecular KE indicated by temperature changes. Potential energy – energy of position or attraction between objects. Indicated by changes in state. Ex: Baseball has potential energy because of gravity’s pull. Further from earth = higher PE. Water boils: molecules move apart = higher PE. Chemical potential energy: atoms and molecules are attracted to each other: Chemical bonds and IM forces PE of phases Gashighest (far apart) | | Liquid to | | solidlowest(close together)

Intermolecular Forces © 2012 Pearson Education, Inc. Heating curves – graph changes in Kinetic energy (temperature) as a sample is heated Ex: For water starting as a solid (ice) below its melting point, to a gas (vapor) above it boiling point Temp 0 C time melting boiling solid Liquid gas

Intermolecular Forces © 2012 Pearson Education, Inc. Temp Phase changes are Changes in potential energy Temperature changes are Changes in kinetic energy Cooling curve: Energy is Released (KineticEnergy)

Intermolecular Forces © 2012 Pearson Education, Inc. Energy Changes Associated with Changes of State Diagonal parts: The heat added to the system mostly goes into kinetic energy (speeding up the motion of molecules. The temperature of the substance rises when kinetic energy is gained.

Intermolecular Forces © 2012 Pearson Education, Inc. Energy Changes Associated with Changes of State Horizontal parts: Temperature stays the same. The heat added to the system at the melting and boiling points goes into potential energy, pulling the molecules apart from each other.

Intermolecular Forces © 2012 Pearson Education, Inc. SAMPLE EXERCISE 11.3 Calculating  H for Temperature and Phase Changes Calculate the enthalpy change upon converting 1.00 mol of ice at –25°C to water vapor (steam) at 125°C under a constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.09 J/g-K, 4.18 J/g-K and 1.84 J/g-K, respectively. For H 2 O,  H fus = 6.01 kJ/mol and  H vap = kJ/mol.

Intermolecular Forces © 2012 Pearson Education, Inc. Solution Analyze: Our goal is to calculate the total heat required to convert 1 mol of ice at –25°C to steam at 125°C. Plan: We can calculate the enthalpy change for each segment and then sum them to get the total enthalpy change (Hess’s law, Section 5.6). Solve: For segment AB in Figure 11.19, we are adding enough heat to ice to increase its temperature by 25°C. A temperature change of 25°C is the same as a temperature change of 25 K, so we can use the specific heat of ice to calculate the enthalpy change during this process: Remember Q=mCdt?

Intermolecular Forces © 2012 Pearson Education, Inc. For segment BC in Figure 11.19, in which we convert ice to water at 0°C, we can use the molar enthalpy of fusion directly: Q=H fus m for ice

Intermolecular Forces © 2012 Pearson Education, Inc. The enthalpy changes for segments CD, DE, and EF can be calculated in similar fashion: Q=mCdt for liquid water Q=H vap m Q=mCdt for water vapor

Intermolecular Forces © 2012 Pearson Education, Inc. The total enthalpy change is the sum of the changes of the individual steps: Check: The components of the total energy change are reasonable in comparison with the lengths of the horizontal segments of the lines in Figure Notice that the largest component is the heat of vaporization.

Intermolecular Forces © 2012 Pearson Education, Inc. PRACTICE EXERCISE What is the enthalpy change during the process in which g of water at 50.0°C is cooled to ice at –30.0°C? (Use the specific heats and enthalpies for phase changes given in Sample Exercise 11.4.) Answer: –20.9 kJ – 33.4 kJ – 6.27 kJ = –60.6 kJ

Intermolecular Forces © 2012 Pearson Education, Inc Vapor Pressure At any temperature some molecules in a liquid have enough energy to break free. As the temperature rises, the fraction of molecules that have enough energy to break free increases.

Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure As more molecules escape the liquid, the pressure they exert increases.

Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate. Equilibrium vapor pressure

Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure Vapor bubbles start to form when their vapor pressure is high enough to overcome the outside air pressure. The boiling point of a liquid then, is the temperature at which its vapor pressure equals atmospheric pressure.

Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure A vapor pressure curve shows changes as temperature increases. The normal boiling point is the temperature at which its vapor pressure reaches standard pressure: 760 torr. For water that is defined as C

Intermolecular Forces © 2012 Pearson Education, Inc. Vapor Pressure Liquids like ethanol and ethers with weaker intermolecular forces will produce higher vapor pressures Notice that those higher vapor pressures cause boiling to occur at lower temperatures

Intermolecular Forces © 2012 Pearson Education, Inc. Normal boiling point means at sea level pressure: 1 atm, 760 torr, kpa

Intermolecular Forces © 2012 Pearson Education, Inc.

Intermolecular Forces © 2012 Pearson Education, Inc.

Intermolecular Forces © 2012 Pearson Education, Inc. Reference table H Which compound has the strongest IM forces? At what temperature does ethanol boil if the outside pressure is reduced to 50 kpa? What is the vapor pressure of propanone if water is boiling at 60 0 C?

Intermolecular Forces © 2012 Pearson Education, Inc Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures, and the places where equilibria exist between phases. Phase diagrams are not heavily tested on the AP but you should recognize them conceptually

Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagrams Let’s draw one for a “typical” liquid Notice at standard pressure the substance has a typical melting and boiling point temperature 1 atm -xx mpbp pressure Temperature Phase diagrams show where solids change to liquids change to gases at different temperatures and pressures

Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagrams 1 atm -xx Increasing pressure raises the boiling point but changes the melting point little solidliquidgas Recall that liquids boil at successively higher temps when pressure is increased? Temperature

Intermolecular Forces © 2012 Pearson Education, Inc. 1 atm -xx Phase Diagrams Decreasing pressure has the same expected effect Notice at some low pressure mp and bp meet at the triple point At the triple point all three phases can exist in equilibrium solidliquidgas

Intermolecular Forces © 2012 Pearson Education, Inc. 1 atm -xx solidliquidgas Phase Diagrams Notice at extreme low pressure the substance would sublime from solid directly to gas

Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagrams xx solidliquidgas Notice at extreme high pressure and high temp no definite line separates the liquid and gas so there is no boiling point. This is the “critical point” At this high pressure gases become like liquids without a transition point.

Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagrams This phase diagram shows the phase changes occuring at different interfaces.

Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagram of Water The slope of the solid– liquid line is negative. –This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid. Freezing of water is requires lower temperatures, when pressure is higher since ice must expand as it forms

Intermolecular Forces © 2012 Pearson Education, Inc. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures. notice at standard pressure only the solid and gas exist

Intermolecular Forces © 2012 Pearson Education, Inc.

Intermolecular Forces © 2012 Pearson Education, Inc.

Intermolecular Forces © 2012 Pearson Education, Inc.