chapter 11 particle forces. states of matter solid- particles moving about a fixed point...

Chapter 11Particle Forces

States of Matter

Solid- Particles moving about a fixed point

Liquid-Particles moving about a moving point

Gas-Particles filling the volume of the container with complete random motions.

Particle Forces Affect• Solubility

• Vapor Pressures

• Freezing Points

• Boiling Points

Particle Forces• Intramolecular forces (Relative strength = 100)

Ionic bonding Covalent bonding

• Interparticle forces Ion-dipole forces Dipole-dipole (Polar molecules)

(relative Strength = 1) London Forces (Dispersion forces)( Nonpolar molecules)

(relative strength = 1) Hydrogen Bonding (Relative strength = 10)

Ion-Ion Interactions• Coulomb’s law states that the energy (E) of

the interaction between two ions is directly proportional to the product of the charges of the two ions (Q1 and Q2) and inversely proportional to the distance (d) between them.

E (Q1Q2)

d

Predicting Forces of Attraction• Coulombs Law indicates the increases in the

charges of ions will cause an increase in the force of attraction between a cation and an anion.

• Increases in the distance between ions will decrease the force of attraction between them.

Size of Ions

Lattice Energy• The lattice energy (U) of an ionic compound

is the energy released when one mole of the ionic compound forms from its free ions in the gas phase.

M+(g) + X-

(g) ---> MX(s)

d)Qk(Q = U 21

Comparing Lattice Energies

Lattice Energies of Common Ionic Compounds

Compound U(kJ/mol)

LiF -1047

LiCl -864

NaCl -790

KCl -720

KBr -691

MgCl2 -2540

MgO -3791

PracticeDetermine which salt has the greater lattice

energy.

A. MgO and NaF

B. MgO and MgS

Lattice Energy Using Hess’s Law

Electron Affinity

• Electron affinity is the energy change occurring when one mole of electrons combines with one mole of atoms or ion in the gas phase.

• Step 4 in diagram on the last slide.

Cl(g) + e-(g) ---> Cl-(g)ΔHEa = -349 kj/mole

Calculating UNa+(g) + e-(g) ---> Na(g) -HIE1

Na(g) ---> Na(s) -Hsub

Cl-(g) ---> Cl(g) + e-(g) -HEA

Cl(g) ---> 1/2Cl2(g) -1/2HBE

Na(s) + 1/2Cl2(g) ---> NaCl(s) Hf

Na+(g) + Cl-(g) ---> NaCl(s) ΔU

U = Hf - 1/2HBE - HEA - Hsub - HIE1

Lattice energy for NaCl.

ΔU

Interactions Involving Polar Molecules

• An ion-dipole interaction occurs between an ion and the partial charge of a molecule with a permanent dipole.

• The cluster of water molecules that surround an ion in aqueous medium is a sphere of hydration.

Illustrates of Ion-Dipole Interaction

The Solution Process

Bond Breaking Processes• Break solute particle forces (expanding

the solute), endothermic• Break solvent particle forces (expanding

the solvent), endothermic

The Solution ProcessAttractive Forces• Energy released when solute solvent are

attracted, exothermic• Energy is released due to new attractions

Ion dipole if the solute is ionic and the solvent polar.

London-Dipole for nonpolar solute and polar solvent

Dipole-dipole for polar solute and polar solvent

The Solution ProcessTheromodynamics

• Enthalpy • Entropy (Perfect crystal, assumed to be

zero)• Gibbs free energy

The Solution ProcessOil dissolving in water• London forces holding the oil molecules

together are large do to the large surface area of the oil

• The hydrogen bonds holding water molecules together are large

• The forces of attraction of between nonpolar oil and polar water are weak at best

• Thus the overall process is highly endothermic and not allowed thermo chemically

The Solution ProcessOil dissolving in water• Entropy should be greater than zero

• Free energy should be greater than zero, since the process is highly endothermic

• Thus the overall process is nonspontaneous

The Solution ProcessSodium chloride dissolving in water

• Large amount of energy is required to break the ionic lattice of the sodium chloride (expand solute)

• Large amount of energy is required to separate the water molecules to expand the solvent breaking hydrogen bonds

• Formation of the ion dipole forces releases a large amount of energy, strong forces (why?)

• The sum of the enthalpies is about +6 kJ (slightly endothermic), which is easily overcome by the entropy of the solution formation.

Water as a Solvent• Water most important solvent, important

to understand its solvent properties

• Most of the unusual solvent properties of water stem from it hydrogen bonding nature

• Consider the following ∆S of solution

KCl →75j/K-mole

LiF→-36j/K-mole

CaS→-138 j/K-mole

Water as a Solvent• We would expect ∆S>0 for all solutions,

right?

• But two are negative, why?

• Obviously, something must be happening for the increased order.

• Ion-dipole forces are ordering the water molecules around the ions, thus causing more order in water i.e. less positions for water than in the pure liquid state

Water as a Solvent• Smaller ions, have stronger ion dipole forces,

thus pulling water closer, therefore less positions

• Also, ions with a charge greater than one will attract to water stronger than a one plus charge, thus more order due to less space between particles

Dipole-Dipole Interactions• Dipole-dipole interactions are

attractive forces between polar molecules.

• An example is the interaction between water molecules.

• The hydrogen bond is a special class of dipole-dipole interactions due to its strength.

Dipole-Dipole ForcesDipole-dipole (Polar molecules)

Alignment of polar molecules to two electrodes

charged + and δ–Forces compared to ionic/covalent are about 1 in strength

compared to a scale of 100, thus 1%

H Cl H Cl H Clδ–δ–δ– δ+δ+ δ+

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Dipole Dipole Interactions

Hydrogen Bonding• Hydrogen bonding a stronger intermolecular

force involving hydrogen and usually N, O, F, and sometimes Cl

–Stronger that dipole-dipole, about 10 out of 100, or 10

–Hydrogen needs to be directly bonded to the heteroatom

–Since hydrogen is small it can get close to the heteroatom

–Also, the second factor is the great polarity of the bond.

Hydrogen Bonding in HF(g)

Hydrogen Bonding in Water

around a molecule in the solid in the liquid

Boiling Points of Binary Hydrides

Interacting Nonpolar Molecules• Dispersion forces (London forces) are

intermolecular forces caused by the presence of temporary dipoles in molecules.

• A temporary dipole (or induced dipole) is a separation of charge produced in an atom or molecule by a momentary uneven distribution of electrons.

Illustrations

Strength of Dispersion Forces• The strength of dispersion forces depends

on the polarizability of the atoms or molecules involved.

• Poarizability is a term that describes the relative ease with which an electron cloud is distorted by an external charge.

• Larger atoms or molecules are generally more polarizable than small atoms or molecules.

London Forces (Dispersion)

• Induced dipoles (Instantaneous )• Strength is surface area dependent• More significant in larger molecules• All molecules show dispersion forces• Larger molecules are more polarizable

Instantaneous and Induced Dipoles

Molar Mass and Boiling Points of Common Species.

Halogen M(g/mol) Bp(K) Noble Gas M(g/mol) Bp(K)

He 2 4

F2 38 85 Ne 20 27

Cl2 71 239 Ar 40 87

Br2 160 332 Kr 84 120

I2 254 457 Xe 131 165

Rn 211 211

Hydrocarbon AlcoholMolecular Formula

Molar Mass

Bp (oC)

Molecular Formula

Molar Mass

Bp (oC)

CH4 16.04 -161.5

CH3CH3 30.07 -88 CH3OH 32.04 64.5

CH3CH2CH3 44.09 -42 CH3CH2OH 46.07 78.5

CH3CH(CH)CH3 58.12 -11.7 CH3CH(OH)CH3 60.09 82

CH3CH2CH2CH3 58.12 -0.5 CH3CH2CH2OH 60.09 97

The Effect of Shape on Forces

Practice Rank the following compound in order of increasing

boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3

Practice

Rank the following compound in order of increasing boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3

CH3OH

CH3CH2CH2CH3

CH3CH2OCH3

MM32.0

58.0

60.0

IM ForcesLondon and H-bonding

London, only

London and Dipole-dipole

Practice Rank the following compound in order of increasing

boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3

CH3OH

CH3CH2CH2CH3

CH3CH2OCH3

MM32.0

58.0

58.0

IM ForcesLondon and H-bonding

London, only

London and Dipole-dipole

The order is:

CH3CH2CH2CH3 < CH3CH2OCH3< CH3OH

Polarity and Solubility• If two or more liquids are miscible, they form

a homogeneous solution when mixed in any proportion.

• Ionic materials are more soluble in polar solvents then in nonpolar solvents.

• Nonpolar materials are soluble in nonpolar solvents.

• Like dissolves like

Solubility of Gases in Water• Henry’s Law states that the solubility of a

sparingly soluble chemically unreactive gas in a liquid is proportional to the partial pressure of the gas.

• Cgas = kHPgas where C is the concentration of the gas, kH is Henry’s Law constant for the gas.

Henry’s Law Constants for Gas

Henry’s Law Constants

Gas kH[mol/(L•atm)] kH[mol/(kg•mmHg)]

He 3.5 x 10-4 5.1 x 10-7

O2 1.3 x 10-3 1.9 x 10-6

N2 6.7 x 10-4 9.7 x 10-7

CO2 3.5 x 10-2 5.1 x 10-5

Terms• A hydrophobic (“water-fearing)

interaction repels water and diminishes water solubility. Polar vs. nonpolar

• A hydrophilic (“water-loving”) interaction attracts water and promotes water solubility. Polar vs. polar, best with hydrogen bonding involved.

Types of Forces• Cohesive Forces

Intermolecular forces between the same particles.• Adhesive Forces

Intermolecular forces between the different particles.

Cohesive ForcesExample Surface tension (resistance to increasing

the surface area)Def: To increase surface area

molecules must move from the middle. This requires energy j/m2

The stronger the IMF the stronger the surface tension

Needle or paper clip on top of waterBeading or wetting on a surfaceRounded surface of liquid mercury

in a tube

Adhesive ForcesExamples.

Capillary rise water forms a meniscus since the forces between the glass and water are stronger than between water and water. Both are hydrogen bonds

Cohesive and Adhesive Forces

The left test tube shows adhesive forces due to the attraction of water solvent to the polar glass (SiO2) Hydrogen bonding, right?

The right tube shows cohesive forces, since mercury is nonpolar and attracts more strongly to itself, rather than to the glass (SiO2)

Terms• Capillary rise is the rise of a liquid up a narrow

tube as a result of adhesive forces between the liquid and the tube and cohesive forces within the liquid.

• Viscosity is a measure of the resistance to flow of a fluid.

Surface Tension

The Liquid StateAdhesive Forces

Intermolecular forces between unlike moleculesExample

Capillary rise• Blood up a capillary• Meniscus• Capillary rise is when the adhesive forces

are stronger than the cohesive forces• Capillary rise when polar bonds are present

in the container walls like glass, SiO2

• Mercury is an example where the cohesive forces are stronger than the adhesive forces

Intermolecular Forces

The Liquid StateViscosity (resistance to flow)

• How fast liquids flow• Due in part to intermolecular forces, but also

entanglement• Newton’s/m2 called poise

Change of State (Water)

H2O(s) H2O(l) H2O(g)

Water Thermodynamic Properties∆Hfus = 6.02 kj/mole

∆Hvap= 40.7 kj/mole

melting evaporation

Condensationfreezing

sublimation

deposition

Change of State (Water)

Heat capacity of water = 4.184 j/g-°C• Water has a very large heat capacity since it

hydrogen bonds and a lot of energy is required to break these bonds

• Why water is used in radiators• Used to cool animals

Sublimation

ΔHsub = ΔHfus + ΔHvap

= -ΔHdeposition

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Some Properties of Solids

Freezing PointMelting Point

ΔHfus(H2O) = +6.01 kJ/mol

Super cooling

Super Cooling and Heating• Super cooled, when a liquid exists below its freezing

point.• Super cooling occurs when the rate of cooling is faster than it

takes for the molecules to rotate for correct alignment to form crystals.

• When the crystals rotate and form inter-particle forces, heat is released, thus raising the temperature up to the correct m.p.

• Super heated Called bumping Use boiling stones, cannot reuse the stones Hot vapor at bottom expands rapidly and bursts

Vapor Pressure• Vaporization or

evaporation is the transformation of molecules in the liquid phase to the gas phase.

• Vapor pressure is the force exerted at a given temperature by a vapor in equilibrium with its liquid phase.

Vapor Pressure

What evaporates faster pure distilled water in the beaker on the left, or seawater in the beaker on the right?? Both beakers are the same size and at the same temperature.

Intermolecular Forces

Vapor PressureYes, pure distilled water evaporates faster, since there are more water molecules on the surface to evaporate?

Vapor PressurePhysical properties that depend on the number of particles, and not on the particle nature are called colligative properties

An Aqueous Solution and Pure Water in a Closed Environment

Vapor Pressure(e) (d) (c) (b) (a)

( ) 1

T

Which one is water?

-ΔHvap+ B

RLn P = Linear

B = y-intercept = ∆SR

(Entropy of vaporization) No Units!

Clausius-Clapeyron Equation

Vapor Pressure

Clasius Clapeyron Equation•Assume data for two different temperatures and pressures to generate two separate equations•By subtracting the equations the y-intercept component is eliminated.

ln P1 = - ΔH/R(1/T1 ) + C- (ln P2 = - ΔH/R(1/T2) + C)ln((P1/P2) = - ΔH/R(1/T2 – 1/T1)

Another useful version of the two point equationln((P1/P2) = - ΔH/R(T2-T1)/T1T2

Vapor Pressure As a liquid evaporates in a closed container the

concentration of vapor increases, thus the rate of condensation increases

As the rate of condensation is increasing eventually it will equal the constant rate of evaporation, then we have vapor in equilibrium with the liquid

The pressure of the vapor at equilibrium is called the equilibrium vapor pressure

Raoult’s Law

Psolution = Xsolvent (Psolvent)

P - vapor pressure

X - mole fractionXsolvent + Xsolute = 1

For a Solution that Obeys Raoult's Law, a Plot fo Psoln Versus Xsolvent, Give a Straight Line

Vapor Pressure of Solvent and Solution

July 2009 General Chemistry: Chapter 11 Slide 74 of 46

Liquid-Vapor Equilibrium

Two Volatile LiquidsPositive deviationIdeal Solution Negative deviation

Positive deviation exists when experimental value is larger than calculated value, weaker solute solvent attraction; more evaporation.

Negative deviation exists when experimental value is smaller than calculated value; stronger solvent solute attraction; less evaporation


Fractional Distillation

Practice A solution contains 100.0 mL of water and 0.500 mol of ethanol. What is the mole fraction of water and the vapor pressure of the solution at 25oC, if the vapor of pressure of pure water is 23.8 torr?


100.0mL


100.0mLmL1.00 g

18.0 gmole


100.0mLmL1.00 g

18.0 gmole

0.500 mole C2H6O= 5.56 mole

6.06 mole

XHOH = 5.566.06 = 0.917


100.0mLmL1.00 g

18.0 gmole

0.500 mole C2H6O= 5.56 mole

6.06 mole

XHOH = 5.566.06 = 0.917

PHOH = 0.917(23.8 torr)PHOH = 21.8 torr

Boiling Point Vs. Pressure

Phase Diagrams• A phase diagram is a graphic representation

of the dependence of the stabilities of the physical states of a substance on temperature and pressure.

Phase Diagram for Water

• Triple Point

• Critical Point

• Critical Temperature

• Critical Pressure

• Supercritical Fluid

The Critical Point

Phase Diagram Terms• The triple point defines the temperature and

pressure where all three phases of a substance coexist.

• The critical point is that specific temperature and pressure at which the liquid and gas phases of a substance have the same density and are indistinguishable for each other.

• A supercritical fluid is a substance at conditions above its critical temperature and pressure.

Phase Diagram for CO2

Colligative Properties of Solutions

• Colligative properties of solutions depend on the concentration and not the identity of particles dissolved in the solvent.

• Sea water boils at a higher temperature than pure water.

Calculating Changes in Boiling Point

Tb = Kbm Tb is the increase in

Bp Kb is the boiling-point

elevation constant m is a new

concentration unit called molality

Molality (m) = moles solute

Kg solvent

Practice Calculate the molality of a solution containing

0.875 mol of glucose (C6H12O6) in 1.5 kg of water.



0.875 mole1.5 kg



0.875 mole1.5 kg

= 0.58 m

Practice Seawater contains 0.558 M Cl- at the surface

at 25oC. If the density of sea water is 1.022 g/mL, what is the molality of Cl- in sea water?



103 mL solutionmL

1.022 g = 1022 g solution



103 mL solutionmL


0.558 mole Cl- 45.45 g Cl-

mole Cl- = 23.36 g Cl-



103 mL solutionmL


0.558 mole Cl- 45.45 g Cl-


1022 g solution – 23.36 g Cl- = 996.6 g H2O



103 mL solutionmL


0.558 mole Cl- 45.45 g Cl-


1022 g solution – 23.36 g Cl- = 996.6 g H2O

0.558 mole Cl-

996.6 g H2O103 gKg

= 0.560 m

Practice Cinnamon owes its flavor and odor to

cinnamaldehyde (C9H8O). Determine the boiling-point elevation of a solution of 100 mg of cinnamaldehyde dissolved in 1.00 g of carbon tetrachloride (Kb = 2.34oC/m).

Practice Cinnamon owes its flavor and odor to

cinnamaldehyde (C9H8O). Determine the boiling-point elevation of a solution of 100 mg of cinnamaldehyde dissolved in 1.00 g of carbon tetrachloride (Kb = 2.34oC/m).

100 mg C9H8O

132.54 mgmmole

1.00 g CCl4 mmole

10-3 mole 103 g

Kg= 0.7545 m

2.34 °Cm

0.7545 m= 1.77°C

Freezing-point Depression

Tf = Kfm

Kf is the freezing-point depression constant and m is the molality.

Practice

The freezing point of a solution prepared by dissolving 1.50 X 102 mg of caffeine in 10.0 g of camphor is 3.07 Celsius degree lower than that of pure camphor (Kf = 39.7oC/m). What is the molar mass of caffeine?

The van’t Hoff Factor

• Tb = iKbm & Tf = iKfm

• van’t Hoff factor, i is the number of ions in one formula unit

The van’t Hoff FactorUsed for ionic compounds, why not osmolarity?• The value of i assumes that all of the salt

dissolves and dissociates in to its component ions

• This is not always true, for example 0.10m NaCl I is 1.87

Ion pairing often occurs in solutions Ion pairing most important in concentrated

solutions Ion pairing important in highly charged solutions

Values of van’t Hoff Factors

Practice CaCl2 is widely used to melt frozen precipitation on

sidewalks after a winter storm. Could CaCl2 melt ice at -20oC? Assume that the solubility of CaCl2 at this temperature is 70.0 g/100.0 g of H2O and that the van’t Hoff factor for a saturated solution of CaCl2 is 2.5 (Kf for water is 1.86 0C/m).

Osmotic Pressure• Osmotic pressure () is the pressure that has

to be applied across a semipermeable membrane to stop the flow of solvent form the the compartment containing pure solvent or a less concentrated solution towards a more concentrated solution.

= iMRT where i is the van’t Hoff factor, M is molarity of solute, R is the idea gas constant (0.00821 l•atm/(mol•K)), and T is in Kelvin

Osmosis at the Molecular Level

Osmotic pressure• Equation from the ideal gas law (pv = nRT) = MRT• Semi permeable membrane• Isotonic same concentration• Cells placed in lower concentration hypotonic, cell

will swell called hemolosis• If concentration on the outside of the cells is

greater then the solution is called hypertonic and the cells shrink called crenation

Osmosis

Figure 10.30

In osmosis, solvent passes through a semipermeable membraneto balance the concentration of solutes in solution on both sidesof the membrane.

ChemTour: Lattice Energy

Click to launch animation

PC | Mac

Students learn to apply Coulomb’s law to calculate the exact lattice energies of ionic solids. Includes Practice Exercises.

ChemTour: Intermolecular Forces


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This ChemTour explores the different types of intermolecular forces and explains how these affect the boiling point, melting point, solubility, and miscibility of a substance. Includes Practice Exercises.

ChemTour: Henry’s Law


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Students learn to apply Henry’s law and calculate the concentration of a gas in solution under varying conditions of temperature and pressure. Includes interactive practice exercises.

ChemTour: Molecular Motion


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Students use an interactive graph to explore the relationship between kinetic energy and temperature. Includes Practice Exercises.

ChemTour: Raoult’s LawClick to launch animation

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Students explore the connection between the vapor pressure of a solution and its concentration as a gas above the solution. Includes Practice Exercises.

ChemTour: Phase DiagramsClick to launch animation

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Students use an interactive phase diagram and animated heating curve to explore how changes in temperature and pressure affect the physical state of a substance.

ChemTour: Capillary Action


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In this ChemTour, students learn that certain liquids will be drawn up a surface if the adhesive forces between the liquid on the surface of the tube exceed the cohesive forces between the liquid molecules.

ChemTour: Boiling and Freezing


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Students learn about colligative properties by exploring the relationship between solute concentration and the temperature at which a solution will undergo phase changes. Interactive exercises invite students to practice calculating the boiling and freezing points of different solutions.

ChemTour: Osmotic PressureClick to launch animation

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Students discover how a solute can build up pressure behind a semipermeable membrane. This tutorial also discusses the osmotic pressure equation and the van’t Hoff factor.

Solubility of CH4, CH2Cl2, and CCl4

Which of the following three compounds is most soluble in water?

A) CH4(g) B) CH2Cl2(λ) C) CCl4(λ)

Solubility of CH4, CH2Cl2, and CCl4

Consider the following arguments for each answer and vote again:

A. A gas is inherently easier to dissolve in a liquid than is another liquid, since its density is much lower.

B. The polar molecule CH2Cl2 can form stabilizing dipole-dipole interactions with the water molecules, corresponding to a decrease in ΔH°soln.

C. The nonpolar molecule CCl4 has the largest molecular mass, and so is most likely to partially disperse into the water, corresponding to an increase in ΔS°soln.

The End

chapter 11 particle forces. states of matter solid- particles moving about a fixed point...

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