Chapter 11 Particle Forces.

Chapter 11 Particle 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 3

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.

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 M+(g) + X-(g) ---> MX(s)
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)

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

Practice Determine which salt has the greater lattice energy.
MgO and NaF MgO and MgS

Lattice Energy Using Hess’s Law

Electron Affinity Cl(g) + e-(g) ---> Cl-(g)
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 U Na+(g) + e-(g) ---> Na(g) -HIE1
Na(g) ---> Na(s) -Hsub Cl-(g) ---> Cl(g) + e-(g) -HEA Cl(g) ---> 1/2Cl2(g) -1/2HBE Na(s) + 1/2Cl2(g) ---> NaCl(s) Hf Na+(g) + Cl-(g) ---> NaCl(s) ΔU U = Hf - 1/2HBE - 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 Process Attractive 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 Process Theromodynamics Enthalpy
Entropy (Perfect crystal, assumed to be zero) Gibbs free energy

The Solution Process Oil 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 Process Oil 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 Process Sodium 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 Forces H Cl H Cl H Cl Dipole-dipole (Polar molecules) δ+
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% δ+ δ– δ+ δ– δ– δ+ H Cl H Cl H Cl

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
Chemistry 140 Fall 2002 Hydrogen Bonding in Water Solid ice has lower density than liquid water. H-bonding holds the ice in a rigid but open structure. Maximum density of water at 3.98 C. 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 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

Hydrocarbon Alcohol 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 -0.5 CH3CH2CH2OH 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 MM IM Forces CH3OH 32.0 London and H-bonding CH3CH2CH2CH3 58.0 London, only CH3CH2OCH3 60.0 London and Dipole-dipole

Practice Rank the following compound in order of increasing boiling point. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3 MM IM Forces CH3OH 32.0 London and H-bonding CH3CH2CH2CH3 58.0 London, only CH3CH2OCH3 58.0 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
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 Forces Example Beading or wetting on a surface
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 water Beading or wetting on a surface Rounded surface of liquid mercury in a tube

Adhesive Forces Examples.
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 right tube shows cohesive forces, since mercury is nonpolar and attracts more strongly to itself, rather than to the glass (SiO2) The left test tube shows adhesive forces due to the attraction of water solvent to the polar glass (SiO2) Hydrogen bonding, right? 51

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. 52

Surface Tension 53

The Liquid State Example Blood up a capillary Meniscus
Adhesive Forces Intermolecular forces between unlike molecules Example 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
Slide 55

The Liquid State Viscosity (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)
deposition freezing Condensation H2O(s) H2O(l) H2O(g) Water Thermodynamic Properties ∆Hfus = 6.02 kj/mole ∆Hvap= 40.7 kj/mole melting evaporation sublimation

Change of State (Water)
Heat capacity of water = 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 Slide 59

Some Properties of Solids
Melting Point Freezing Point Super cooling ΔHfus(H2O) = kJ/mol Slide 60 of 35

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

Vapor Pressure Physical 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 Clausius-Clapeyron Equation Linear Which one is water?
(e) (d) (c) (b) (a) Clausius-Clapeyron Equation ( ) 1 T -ΔHvap Linear Ln P = + B R ∆S Which one is water? (Entropy of vaporization) No Units! B = y-intercept = R Slide 68

ln((P1/P2) = - ΔH/R(T2-T1)/T1T2
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 equation ln((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 fraction Xsolvent + 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

Liquid-Vapor Equilibrium
Chemistry 140 Fall 2002 Liquid-Vapor Equilibrium The vapor phase is richer in the more volatile component. July 2009 General Chemistry: Chapter 11 Slide 74 of 46 74

Two Volatile Liquids Ideal Solution Positive deviation
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
July 2009 General Chemistry: Chapter 11 Slide 76 of 46

Fractional Distillation
July 2009 General Chemistry: Chapter 11 Slide 77 of 46

Practice A solution contains mL of water and 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?

Practice A solution contains mL of water and 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

Practice A solution contains mL of water and 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? 1.00 g mole 100.0mL mL 18.0 g

Practice A solution contains mL of water and 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? 1.00 g mole 100.0mL = 5.56 mole mL 18.0 g 0.500 mole C2H6O 6.06 mole 5.56 XHOH = = 0.917 6.06

Practice A solution contains mL of water and 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? 1.00 g mole 100.0mL = 5.56 mole mL 18.0 g 0.500 mole C2H6O 6.06 mole PHOH = 0.917(23.8 torr) PHOH = 21.8 torr 5.56 XHOH = = 0.917 6.06

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 moles solute Molality (m) = Kg solvent

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

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

Practice Calculate the molality of a solution containing mol of glucose (C6H12O6) in 1.5 kg of water. 0.875 mole = 0.58 m 1.5 kg

Practice Seawater contains M Cl- at the surface at 25oC. If the density of sea water is g/mL, what is the molality of Cl- in sea water?

Practice Seawater contains M Cl- at the surface at 25oC. If the density of sea water is g/mL, what is the molality of Cl- in sea water? 103 mL solution 1.022 g = 1022 g solution mL

Practice Seawater contains M Cl- at the surface at 25oC. If the density of sea water is g/mL, what is the molality of Cl- in sea water? 103 mL solution 1.022 g = 1022 g solution mL 0.558 mole Cl- 45.45 g Cl- = g Cl- mole Cl-

Practice Seawater contains M Cl- at the surface at 25oC. If the density of sea water is g/mL, what is the molality of Cl- in sea water? 103 mL solution 1.022 g = 1022 g solution mL 0.558 mole Cl- 45.45 g Cl- = g Cl- mole Cl- 1022 g solution – g Cl- = g H2O

Practice Seawater contains M Cl- at the surface at 25oC. If the density of sea water is g/mL, what is the molality of Cl- in sea water? 103 mL solution 1.022 g = 1022 g solution mL 0.558 mole Cl- 45.45 g Cl- = g Cl- mole Cl- 1022 g solution – g Cl- = g H2O 0.558 mole Cl- 103 g = m 996.6 g H2O Kg

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 10-3 mole 103 g mmole = m 1.00 g CCl4 mg mmole Kg 2.34 °C m = 1.77°C 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 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 Factor Used 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 C/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 ( l•atm/(mol•K)), and T is in Kelvin 107

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 In osmosis, solvent passes through a semipermeable membrane
to balance the concentration of solutes in solution on both sides of the membrane. Figure 10.30

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
Click to launch animation PC | Mac 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.

Click to launch animation
ChemTour: Henry’s Law Click to launch animation PC | Mac 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
Click to launch animation PC | Mac Students use an interactive graph to explore the relationship between kinetic energy and temperature. Includes Practice Exercises.

ChemTour: Raoult’s Law
Click to launch animation PC | Mac 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
Click to launch animation PC | Mac 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
Click to launch animation PC | Mac 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
Click to launch animation PC | Mac 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
Click to launch animation PC | Mac 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? © 2008 W. W. Norton & Company Inc. All rights reserved. A) CH4(g) B) CH2Cl2(λ) C) CCl4(λ) Solubility of CH4, CH2Cl2, and CCl4

Solubility of CH4, CH2Cl2, and CCl4
Consider the following arguments for each answer and vote again: A gas is inherently easier to dissolve in a liquid than is another liquid, since its density is much lower. The polar molecule CH2Cl2 can form stabilizing dipole-dipole interactions with the water molecules, corresponding to a decrease in ΔH°soln. 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. Answer: B Solubility of CH4, CH2Cl2, and CCl4

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Chapter 11 Particle Forces.

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