States of Matter - Yonsei Universitychem.yonsei.ac.kr/~mhmoon/pdf/GenChem_Brown/11_Liquids.pdf · 2013-08-27 · Intermolecular Forces 3 States of Matter • Because in the solid

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Chapter 11

Liquids and Intermolecular

Forces

Lecture Presentation

IntermolecularForces

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States of MatterThe fundamental difference between states of matter is the distance between particles.


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States of Matter

• Because in the solid and liquid states, particles are closer together, we refer to them as condensed phases.

• The state a substance is in at a particular temperature and pressure depends on two antagonistic entities:– The kinetic energy of the particles.– The strength of the attractions between the particles.


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Intermolecular Forces

• The attractions between molecules are not nearly as strong as the intramolecular attractions 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.

• These intermolecular forces as a group are referred to as van der Waals forces.


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van der Waals Forces

• Dipole–dipole interactions

• Hydrogen bonding

• London dispersion forces


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London Dispersion Forces


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While the electrons in the 1s orbital of heliumwould repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.


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At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.


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Another helium atom nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.


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London Dispersion ForcesLondon dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

• 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.


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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.


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Factors Affecting London Forces

• The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane) tend to have stronger dispersion forces than short, fat ones (like neopentane).

• This is due to the increasedsurface area in n-pentane.


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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.


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Dipole-Dipole Forces


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Dipole–Dipole Interactions

The more polar the molecule,

the higher its boiling point.


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Which Have a Greater Effect?Dipole–Dipole Interactions or Dispersion Forces

• 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.


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How Do We Explain This?

• The nonpolar series (SnH4 to CH4) follow the expected trend.

• The polar series follow the trend until you get to the smallest molecules in each group.


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Hydrogen Bonds


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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.

• Due to the high electronegativity of nitrogen, oxygen, and fluorine

• When hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.


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H-Bonds in DNA


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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.


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Summarizing Intermolecular Forces


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Sample Exercise 11.2 Predicting Types and Relative Strengths of Intermolecular Attractions

SolutionH2 < Ne < CO < HF < BaCl2

Check boiling points are H2 (20 K), Ne (27 K), CO (83 K), HF (293 K), and BaCl2(1813 K)—in agreement with our predictions.

List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling point.

(a) Identify the intermolecular attractions present in the following substances, and (b)select the substance with the highest boiling point: CH3CH3, CH3OH, and CH3CH2OH.

Answers: (a) CH3CH3 has only dispersion forces, whereas the other two substances have both dispersion forces and hydrogen bonds, (b) CH3CH2OH

Practice Exercise


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Intermolecular Forces Affect Many Physical Properties

The strength of the attractions between particles can greatly affect the properties of a substance or solution.


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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.


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Surface Tension

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.: Surface tension is the amount of energy required to increase the surface area of a liquid by a unit amount.

Ex) H2O(l), Hg(l)


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Cohesive and adhesive forces

• Cohesive forces are intermolecular forces that bind molecules to one another.

• Adhesive forces are intermolecular forces that bind molecules to a surface.

• Illustrate this by looking at the meniscus in a tube filled with liquid.

- The meniscus is the shape of the liquid surface.

- If adhesive forces are greater than cohesive forces, the liquid surface is attracted to its container more than the bulk molecules. Therefore, the meniscus is U-shaped (e.g., water in glass).

- If cohesive forces are greater than adhesive forces, the meniscus is curved downwards (e.g., Hg(l) in glass)


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Capillary action

• Capillary action is the rise of liquids up very narrow tubes.

• The liquid climbs until adhesive and cohesive forces are balanced by gravity.


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Phase Changes


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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.


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The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas.

The heat of sublimation is defined as the energy required to change a solid directly to a gas. Generally ∆Hfusion < ∆Hvap

• deposition: ∆Hdep < 0 (exothermic)

• condensation: ∆Hcon < 0 (exothermic)

• freezing: ∆Hfre < 0 (exothermic)


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• The heat added to the system at the melting and boiling points goes intopulling the molecules farther apart from each other.

• The temperature of the substance does not rise during a phase change.

• Supercooling: When a liquid is cooled below its freezing point and it still remains a liquid.


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Heating Curves


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Sample Exercise 11.3 Calculating H for Temperature and Phase Changes

Solution

Calculate the enthalpy change upon converting 1.00 mol of ice at 25 C to steam at 125 C under a constant pressure of 1 atm. The specific heats of ice, liquid water, and steam are 2.03 J/g-K, 4.18 J/g-K, and 1.84 J/g-K, respectively. For H2O, Hfus = 6.01 kJ/mol and Hvap = 40.67 kJ/mol.

AB: H = (1.00 mol)(18.0 g/mol)(2.03 J/g-K)(25 K) = 914 J = 0.91 kJ

CD: H = (1.00 mol)(18.0 g/mol)(4.18 J/g-K)(100 K) = 7520 J = 7.52 kJDE: H = (1.00 mol)(40.67 kJ/mol) = 40.7 kJEF: H = (1.00 mol)(18.0 g/mol)(1.84 J/g-K)(25 K) = 830 J = 0.83 kJ

H = 0.91 kJ + 6.01 kJ + 7.52 kJ + 40.7 kJ + 0.83 kJ = 56.0 kJ

BC: H = (1.00 mol)(6.01 kJ/mol) = 6.01 kJ

Practice ExerciseWhat is the enthalpy change during the process in which 100.0 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 11.3.)

Answer: 20.9 kJ 33.4 kJ 6.09 kJ = 60.4 kJ


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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.


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Vapor PressureAs more molecules escape the liquid, the pressure they exert increases.

The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.


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Vapor Pressure

• The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.

• The normal boiling point is the temperature at which its vapor pressure is 760 torr.


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Vapor PressureThe natural log of the vapor pressure of a liquid is inversely proportional to its temperature.

This relationship is quantified in the Clausius–Clapeyron equation:

ln P = −Hvap/RT + C,

where C is a constant.


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Volatility, Vapor Pressure, and Temperature

• If equilibrium is never established, the vapor continues to form.

- Eventually, the liquid evaporates to dryness.

• Liquids that evaporate easily are said to be volatile.

- The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.


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Phase Diagrams: the state of a substance at various pressures and temperatures, and the places where equilibria exist between phases.


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Phase Diagramstriple point (T), critical point (C), TC: boiling point

curve, melting point curve, sublimation point curve

Below the triple point the substance cannot exist in the liquid state.


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Critical Temperature and Pressure

• Gases may be liquefied by increasing the pressure at a suitable temperature.

• Critical temperature is the highest temperature at which a substance can exist as a liquid.

• Critical pressure is the pressure required for liquefaction at this critical temperature.

-The greater the intermolecular forces, the easier it is to liquefy a substance.

- Thus, the higher the critical temperature.

• Supercritical fluid.

- Supercritical fluid extraction (SF CO2)


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Phase Diagram of Water• Note the high critical

temperature and critical pressure.– These are due to the

strong van der Waals forces between water molecules.

The slope of the solid–liquid line is negative.This means that as the pressure is increased at a temperature, melting point decreases, thus water goes from a solid to a liquid.


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Phase Diagram of CO2

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2

sublimes at normal pressures.


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Sample Exercise 11.5 Interpreting a Phase Diagram

SolutionSolve (a) The critical point is approximately 80 C and 50 atm.(b) The triple point is approximately 180 C and 0.1 atm.(c) point 2 : gas(d) Solid methane melts(e) gas to supercritical fluid ( … exceed the critical pressure) (~50 atm).

Use the phase diagram for methane, CH4, shown in Figure 11.30 to answer the following questions. (a) What are the approximate temperature and pressure of the critical point? (b) What are the approximate temperature and pressure of the triple point? (c) Is methane a solid, liquid, or gas at 1 atm and 0 C? (d) If solid methane at 1 atm is heated while the pressure is held constant, will it melt or sublime? (e) If methane at 1 atm and 0 C is compressed until a phase change occurs, in which state is the methane when the compression is complete?

Practice ExerciseUse the phase diagram of methane to answer the following questions. (a) What is the normal boiling point of methane? (b) Over what pressure range does solid methane sublime? (c) Liquid methane does not exist above what temperature?Answer: (a) 162 C; (b) It sublimes whenever the pressure is less than 0.1 atm; (c) The highest temperature at which a liquid can exist is defined by the critical temperature. So we do not expect to find liquid methane when the temperature is higher than 80 C.


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Liquid Crystals• Some substances do

not go directly from the solid state to the liquid state.

• In this intermediate state, liquid crystals have some traits of solids and some of liquids.


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Liquid Crystals

Unlike liquids, molecules in liquid crystals have some degree of order.


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Types of Liquid-Crystalline Phases

In nematic liquid crystals, molecules are only ordered in one dimension, along the long axis.


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In smectic liquid crystals, molecules are ordered in two dimensions, along the long axis and in layers.


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In cholesteryl liquid crystals, nematic-like crystals are layered at angles to each other.


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Liquid Crystals

These crystals can exhibit color changes with changes in temperature.

LCD →


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Sample Exercise 11.6 Properties of Liquid Crystals

SolutionMolecule (ii) possesses the characteristic long axis and the kinds of structural features often seen in liquid crystals: The molecule has a rodlike shape, the double bonds and benzene rings provide rigidity, and the polar COOCH3 group creates a dipole moment.

Which of these substances is most likely to exhibit liquid crystalline behavior?

Answer: Because rotation can occur about carbon–carbon single bonds, molecules whosebackbone consists predominantly of C C single bonds are too flexible; the molecules tend tocoil in random ways and, thus, are not rodlike.

Practice ExerciseSuggest a reason why decane

does not exhibit liquid crystalline behavior.

CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3


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Problems

• 10, 20, 36, 48, 60, 78, 84, 90

Documents

States of Matter - Yonsei Universitychem.yonsei.ac.kr/~mhmoon/pdf/GenChem_Brown/11_Liquids.pdf · 2013-08-27 · Intermolecular Forces 3 States of Matter • Because in the solid