Chapter 10 Particle Forces. States of Matter Solid- Particles moving about a fixed point...

Chapter 10Particle 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) U

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

Lattice energy for NaCl.

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+

of 35

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 instantaneous 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

Molar Mass and Boiling Point

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

London vs Hydrogen Bonding

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

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.


How does polarity effect solubility?

The thermodynamic argument, is that the lower the potential energy, the more stable the system. If subtracting the potential energy of the solute from the potential energy of the original solute and solvent is negative (exothermic) then solution is thermodynamically favored.


How does polarity effect solubility?

Non polar solute and solvent: The forces holding these particles together are London Dispersion forces, the weakest of all of the inter-particle forces. The strength of these forces are relative to the surface area if solute and solvent are of similar size, then about the same amount of energy is required to separate solute and solvent particles from each other. And about the same amount of energy is released when solute and solvent are attracted to each other forming a solution. Thus we predict non polar solutes and solvents should dissolve

Polarity and SolubilityHow does polarity effect solubility?

Non polar solute and polar solvent: Considering solutes and solvents of similar surface area it should be noted that more energy is required to separate the polar solvent molecules from each other, since dipole-dipole interactions are stronger. The only interaction between a nonpolar solute and polar solvent would be London Dispersion forces, so the energy released is much less than required for separating the solvent and solute. Subtracting the potential energy of the products from reactants would give a positive (endothermic) result and the solution would be less stable than the dissolution.

Practice Rank the following compound in order

of increasing boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3

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

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.

• A hydrophilic (“water-loving”) interaction attracts water and promotes water solubility.

Affects of Intermolecular Forces

• Solubility

• Vapor Pressures

• Freezing Points

• Boiling Points

• Surface tension

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

Vapor Pressure

The normal boiling point of a liquid is the temperature at which its vapor pressure equals 1 atmosphere.

Vapor Pressure of SolutionsWhat evaporates faster, sugar water or pure water?

s w s w s w w w w w w w



w(g)w(g) w(g)



w(g)w(g) w(g)

Pure water evaporates faster, since there are more water particles on the surface, thus lowering the average kinetic energy. Evaporation of a solution is inversely proportional to concentration.

Vapor Pressure of Solutions

• Raoult’s Law

Psolution = Xsolvent (Psolvent)

• P - vapor pressure

• X - mole fraction

• Xsolute + Xsolvent = 1

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?

Surface Tension

Cohesive and Adhesive Forces Produce a Meniscus

Physical State and Phase Transformations

• 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

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

Terms

• Capillary action 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.

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.

Colligative PropertiesColligative property is a physical property that depends on the number of particles present, and not on the nature of the particle. Since evaporation is dependent on the number of solvent particles present on the surface that makes evaporation and vapor pressure colligative properties.

Colligative Properties

Is density a colligative property?

Colligative PropertiesIs density a colligative property?

While density depends on the number of particles in a given area, it is also effected by the weight of the substance, which is a nature thing, so no density is not a colligative property.

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

m= nsolute

kg solvent

Practice

Calculate the molality of a solution containing 0.875 mol of glucose (C6H12O6) in 1.5 kg of water.

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?

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).

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

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).

Osmosis

Figure 10.30

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

Osmosis at the Molecular Level

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

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 Law


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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 Diagrams


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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 Points


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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 Pressure


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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 10 Particle Forces. States of Matter Solid- Particles moving about a fixed point...

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