Liquids, Solids, and Intermolecular Forces
Mr. Matthew Totaro Legacy High School AP Chemistry Copyright 2011 Pearson Education, Inc.
Properties of the Three Phases of Matter of r t h la n g cu re le St o rm ns te o In t i ac ? ttr ow A Fl it e? ill bl si W es pr om C
State Shape Volume fixed fixed Solid Liquid indefinite fixed Gas indefinite indefinite
Density high high low
No No No Yes Yes Yes
very strong intermediate weak
• Fixed = keeps shape when placed in a container • Indefinite = takes the shape of the container 2
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Three Phases of Water Notice that the densities of ice and liquid water are much larger than the density of steam Notice that the densities and molar volumes of ice and liquid water are much closer to each other than to steam Notice that the densities of ice is larger than the density of liquid water. This is not the norm, but is vital to the development of life as we know it. 3
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Degrees of Freedom
• Particles may have one or several types of freedom of motion
and various degrees of each type
• Translational freedom is the ability to move from one position in space to another • Rotational freedom is the ability to reorient the particles direction in space • Vibrational freedom is the ability to oscillate about a particular point in space 4
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Solids
• The particles in a solid are packed close together and are fixed in position though they may vibrate
• The close packing of the particles
results in solids being incompressible • The inability of the particles to move around results in solids retaining their shape and volume when placed in a new container, and prevents the solid from flowing 5
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Solids • Some solids have their particles
arranged in an orderly geometric pattern – we call these crystalline solids salt and diamonds
• Other solids have particles that do not show a regular geometric pattern over a long range – we call these amorphous solids plastic and glass
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Liquids
• The particles in a liquid are closely
packed, but they have some ability to move around • The close packing results in liquids being incompressible • But the ability of the particles to move allows liquids to take the shape of their container and to flow – however, they don’t have enough freedom to escape or expand to fill the container 7
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Kinetic – Molecular Theory
• What state a material is in depends largely on two major factors
1. the amount of kinetic energy the particles possess 2. the strength of attraction between the particles
• These two factors are in competition with each other
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States and Degrees of Freedom • The molecules in a gas have complete freedom of motion
their kinetic energy overcomes the attractive forces between the molecules
• The molecules in a solid are locked in place, they cannot move around
though they do vibrate, they don’t have enough kinetic energy to overcome the attractive forces
• The molecules in a liquid have limited freedom – they can move around a little within the structure of the liquid
they have enough kinetic energy to overcome some of the attractive forces, but not enough to escape each other 9
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Kinetic Energy
• Increasing kinetic energy
increases the motion energy of the particles • The more motion energy the molecules have, the more freedom they can have • The average kinetic energy is directly proportional to the temperature KEavg = 1.5 kT
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Attractive Forces
• The particles are attracted to each other by
electrostatic forces • The strength of the attractive forces varies, some are small and some are large • The strength of the attractive forces depends on the kind(s) of particles • The stronger the attractive forces between the particles, the more they resist moving though no material completely lacks particle motion
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Kinetic–Molecular Theory of Solids
• When the attractive forces are strong enough so the kinetic energy cannot overcome it at all, the material will be a solid • In a solid, the particles are packed together without any translational or rotational motion the only freedom they have is vibrational motion
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Kinetic–Molecular Theory of Liquids
• When the attractive
forces are strong enough so the kinetic energy can only partially overcome them, the material will be a liquid • In a liquid, the particles are packed together with only very limited translational or rotational freedom 13
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Explaining the Properties of Liquids
• Liquids have higher densities than gases and are
incompressible because the particles are in contact • They have an indefinite shape because the limited translational freedom of the particles allows them to move around enough to get to the container walls • It also allows them to flow • But they have a definite volume because the limit on their freedom keeps the particles from escaping each other 14
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Phase Changes
• Because the attractive forces between the molecules are fixed, • •
changing the material’s state requires changing the amount of kinetic energy the particles have, or limiting their freedom Solids melt when heated because the particles gain enough kinetic energy to partially overcome the attactive forces Liquids boil when heated because the particles gain enough kinetic energy to completely overcome the attractive forces the stronger the attractive forces, the higher you will need to raise the temperature
• Gases can be condensed by decreasing the temperature and/or increasing the pressure
pressure can be increased by decreasing the gas volume reducing the volume reduces the amount of translational freedom the particles have
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Phase Changes
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Intermolecular Attractions • The strength of the attractions between the
particles of a substance determines its state • At room temperature, moderate to strong attractive forces result in materials being solids or liquids • The stronger the attractive forces are, the higher will be the boiling point of the liquid and the melting point of the solid other factors also influence the melting point
Tro: Chemistry: A Molecular Approach, 2/e
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Why Are Molecules Attracted to Each Other? • Intermolecular attractions are due to attractive forces between opposite charges
+ ion to − ion + end of polar molecule to − end of polar molecule H-bonding especially strong
even nonpolar molecules will have temporary charges
• Larger charge = stronger attraction • Longer distance = weaker attraction • However, these attractive forces are small relative to the bonding forces between atoms generally smaller charges generally over much larger distances 18
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Trends in the Strength of Intermolecular Attraction • The stronger the attractions between the atoms or molecules, the more energy it will take to separate them • Boiling a liquid requires we add enough energy to overcome all the attractions between the particles However, not breaking the covalent bonds
• The higher the normal boiling point of the liquid, the stronger the intermolecular attractive forces
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Kinds of Attractive Forces • Temporary polarity in the molecules due to
unequal electron distribution leads to attractions called dispersion forces • Permanent polarity in the molecules due to their structure leads to attractive forces called dipole– dipole attractions • An especially strong dipole–dipole attraction results when H is attached to an extremely electronegative atom. These are called hydrogen bonds. 20
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Dispersion Forces
• Fluctuations in the electron distribution in atoms and molecules result in a temporary dipole
region with excess electron density has partial (─) charge region with depleted electron density has partial (+) charge
• The attractive forces caused by these temporary dipoles are called dispersion forces aka London Forces
• All molecules and atoms will have them • As a temporary dipole is established in one molecule, it induces a dipole in all the surrounding molecules
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Dispersion Force
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Size of the Induced Dipole • The magnitude of the induced dipole •
•
depends on several factors Polarizability of the electrons
+ + + + + + + - - - --
volume of the electron cloud
larger molar mass = more electrons = larger electron cloud = increased polarizability = stronger attractions Shape of the molecule more surface-to-surface contact = larger induced dipole = stronger attraction
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molecules that are flat larger molecules have have more surface more electrons, leading interaction than to increased polarizability spherical ones
+ + + + + + + + + + ++ + + + + − −- − − −− − − −- − −− −
++
+
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Effect of Molecular Size on Size of Dispersion Force As molar mass Thethe Noble gases increases, the number are all nonpolar of electrons increases. atomic elements Therefore the strength of the dispersion forces increases. The stronger the attractive forces between the molecules, the higher the boiling point will be. 24
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Properties of Straight Chain Alkanes NonPolar Molecules
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Boiling Points of n-Alkanes
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Effect of Molecular Shape on Size of Dispersion Force the larger surface-tosurface contact between molecules in n-pentane results in stronger dispersion force attractions
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Alkane Boiling Points • Branched
chains have lower BPs than straight chains • The straight chain isomers have more surface-tosurface contact 30
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Practice – Choose the Substance in Each Pair with the Higher Boiling Point a) CH4
C4H10
b) C6H12
C6H12
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Practice – Choose the Substance in Each Pair with the Higher Boiling Point a) CH4
CH3CH2CH2CH3
b) CH3CH2CH=CHCH2CH3
Both molecules are nonpolar larger molar mass
cyclohexane Both molecules are nonpolar, but the flatter ring molecule has larger surface-tosurface contact
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Dipole–Dipole Attractions
• Polar molecules have a permanent dipole because of bond polarity and shape dipole moment as well as the always present induced dipole
• The permanent dipole adds to the attractive forces between the molecules
raising the boiling and melting points relative to nonpolar molecules of similar size and shape
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Effect of Dipole–Dipole Attraction on Boiling and Melting Points
Tro: Chemistry: A Molecular Approach, 2/e
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Example: Determine if dipole–dipole attractions occur between CH2Cl2 molecules Given: CH2Cl2, EN C = 2.5, H = 2.1, Cl = 3.0 Find: If there are dipole–dipole attractions Conceptual Plan:
Formula
Lewis Structure
Bond Polarity
EN Difference
Relationships: Solution:
Molecule Polarity Shape
molecules that have dipole–dipole attractions must be polar
Cl—C
3.0−2.5areas = 0.5 4 bonding polarpairs = no lone tetrahedral C—H shape 2.5−2.1 = 0.4 nonpolar 36
polar bonds and tetrahedral shape = polar molecule polar molecule; therefore dipole– dipole attractions
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Practice – Choose the substance in each pair with the higher boiling point a)
CH2FCH2F
CH3CHF2
b) or
Tro: Chemistry: A Molecular Approach, 2/e
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Practice – Choose the substance in each pair with the higher boiling point a)
CH2FCH2F
CH3CHF2
more polar
b) or polar
nonpolar
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Hydrogen Bonding
• When a very electronegative atom is bonded to
hydrogen, it strongly pulls the bonding electrons toward it O─H, N─H, or F─H
• Because hydrogen has no other electrons, when its electron is pulled away, the nucleus becomes deshielded exposing the H proton
• The exposed proton acts as a very strong center of positive charge, attracting all the electron clouds from neighboring molecules 39
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H-Bonding
HF
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H-Bonding in Water
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H-Bonds • Hydrogen bonds are very strong intermolecular attractive forces
stronger than dipole–dipole or dispersion forces
• Substances that can hydrogen bond will have
higher boiling points and melting points than similar substances that cannot • But hydrogen bonds are not nearly as strong as chemical bonds 2 to 5% the strength of covalent bonds 42
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Effect of H-Bonding on Boiling Point
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For NH3 have unusually such as the HF,nonpolar H2O, andmolecules, hydrides strong dipole-dipole of Group 4, attractions, the intermolecular called hydrogen attractions bonds. Therefore are due to dispersion they have forces. higher boiling Therefore points they than increase you would down expect the from column, the causing general the trends. boiling point to increase. 44
Polar molecules, such as the hydrides of Groups 5–7, have both dispersion forces and dipole–dipole attractions. Therefore they have higher boiling points than the corresponding Group 4 molecules.
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Example: One of these compounds is a liquid at room temperature (the others are gases). Which one and why? MM = 30.03 Polar No H-Bonds
MM = 34.03 Polar No H-Bonds
MM = 34.02 Polar H-Bonds
Step 2. 3. Compare 1. Evaluate the Determine intermolecular thestrengths kinds of of intermolecular attractive the total intermolecular forces attractive forces attractive forces. The substance with the strongest will be Because the molar masses are similar, the size of the liquid. Fluoromethane: Formaldehyde: Hydrogen peroxide: the dispersion force attractions shouldplanar bebent similar 30.03, tetrahedral 34.03, trigonal dispersion forces: MM 34.02, Because only hydrogen peroxide has the additional very very polar uncancelled dipole–dipole: O–H C–F bonds uncancelled C=O bond uncancelled Because thepolar molecules arebond polar, the size of the strong H-bond all additional attractions, its intermolecular noattractions O–H, N–H,should orH-bond F–Hbetherefore no H-bond H-bonding: O–H, therefore dipole–dipole similarexpect attractions will be the strongest. We therefore hydrogen to be peroxide the liquid.also has additional Onlyperoxide the hydrogen hydrogen bond attractions Tro: Chemistry: A Molecular Approach, 2/e
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Practice – Choose the substance in each pair that is a liquid at room temperature (the other is a gas)
a) CH3OH
CH3CHF2
can H-bond
b) CH3-O-CH2CH3
CH3CH2CH2NH2 can H-bond
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Attractive Forces and Solubility
• Solubility depends, in part, on the attractive forces of the solute and solvent molecules
like dissolves like miscible liquids will always dissolve in each other
• Polar substances dissolve in polar solvents
hydrophilic groups = OH, CHO, C=O, COOH, NH2, Cl
• Nonpolar molecules dissolve in nonpolar solvents hydrophobic groups = C-H, C-C
• Many molecules have both hydrophilic and hydrophobic parts – solubility in water becomes a competition between the attraction of the polar groups for the water and the attraction of the nonpolar groups for their own kind 47
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Immiscible Liquids
Pentane, C5H12 is a nonpolar molecule. Water is a polar molecule. The attractive forces between the water molecules is much stronger than their attractions for the pentane molecules. The result is the liquids are immiscible.
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Polar Solvents Dichloromethane (methylene chloride) Water Ethanol (ethyl alcohol)
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Nonpolar Solvents
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Ion–Dipole Attraction
• In a mixture, ions from an ionic compound are
attracted to the dipole of polar molecules • The strength of the ion–dipole attraction is one of the main factors that determines the solubility of ionic compounds in water
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Practice – Choose the substance in each pair that is more soluble in water a) CH3OH
CH3CHF2
can H-bond with H2O
b) CH3CH2CH2CH2CH3
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CH3Cl
more polar
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Summary • Dispersion forces are the weakest of the
intermolecular attractions • Dispersion forces are present in all molecules and atoms • The magnitude of the dispersion forces increases with molar mass • Polar molecules also have dipole–dipole attractive forces 53
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Summary (cont’d)
• Hydrogen bonds are the strongest of the intermolecular attractive forces a pure substance can have
• Hydrogen bonds will be present when a molecule has H directly bonded to either O , N, or F atoms only example of H bonded to F is HF
• Ion–dipole attractions are present in mixtures of
ionic compounds with polar molecules. • Ion–dipole attractions are the strongest intermolecular attraction • Ion–dipole attractions are especially important in aqueous solutions of ionic compounds 54
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Tro: Chemistry: A Molecular Approach, 2/e
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Liquids
Properties & Structure Copyright 2011 Pearson Education, Inc.
Surface Tension
• Surface tension is a property of liquids that
results from the tendency of liquids to minimize their surface area • To minimize their surface area, liquids form drops that are spherical as long as there gravity
is no
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Surface Tension
• The layer of molecules on the surface behave differently than the interior
because the cohesive forces on the surface molecules have a net pull into the liquid interior
• The surface layer acts like an elastic skin
allowing you to “float” a paper clip even though steel is denser than water
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Surface Tension • Because they have fewer neighbors to attract them, the surface molecules are less stable than those in the interior have a higher potential energy
• The surface tension of a liquid is
the energy required to increase the surface area a given amount
surface tension of H2O = 72.8 mJ/m2 at room temperature
surface tension of C6H6 = 28 mJ/m2 59
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Factors Affecting Surface Tension
• The stronger the intermolecular attractive forces, the higher the surface tension will be • Raising the temperature of a liquid reduces its surface tension raising the temperature of the liquid increases the average kinetic energy of the molecules the increased molecular motion makes it easier to stretch the surface
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Viscosity
• Viscosity is the resistance of a liquid to flow 1 poise = 1 P = 1 g/cm∙s often given in centipoise, cP
H2O = 1 cP at room temperature
• Larger intermolecular attractions = larger viscosity
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Factors Affecting Viscosity
• The stronger the intermolecular attractive forces, the
higher the liquid’s viscosity will be • The more spherical the molecular shape, the lower the viscosity will be molecules roll more easily less surface-to-surface contact lowers attractions
• Raising the temperature of a liquid reduces its viscosity
raising the temperature of the liquid increases the average kinetic energy of the molecules the increased molecular motion makes it easier to overcome the intermolecular attractions and flow 63
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Insert Table 11.6
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Capillary Action • Capillary action is the ability of a liquid to flow up a thin tube against the influence of gravity
the narrower the tube, the higher the liquid rises
• Capillary action is the result of two forces working in conjunction, the cohesive and adhesive forces
cohesive forces hold the liquid molecules together adhesive forces attract the outer liquid molecules to the tube’s surface
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Capillary Action
• The adhesive forces pull the surface liquid up the
side of the tube, and the cohesive forces pull the interior liquid with it • The liquid rises up the tube until the force of gravity counteracts the capillary action forces • The narrower the tube diameter, the higher the liquid will rise up the tube
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Meniscus
• The curving of the liquid surface in a thin
tube is due to the competition between adhesive and cohesive forces • The meniscus of water is concave in a glass tube because its adhesion to the glass is stronger than its cohesion for itself • The meniscus of mercury is convex in a glass tube because its cohesion for itself is stronger than its adhesion for the glass metallic bonds are stronger than intermolecular attractions 67
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The Molecular Dance
• Molecules in the liquid are constantly in motion
vibrational, and limited rotational and translational
• The average kinetic energy is proportional to the temperature • However, some molecules have more kinetic energy than the average, and others have less
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Vaporization
• If these high energy molecules are at the surface, they may have enough energy to overcome the attractive forces therefore – the larger the surface area, the faster the rate of evaporation
• This will allow them to escape the liquid and become a vapor
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Distribution of Thermal Energy
• Only a small fraction of the molecules in a liquid have enough energy to escape • But, as the temperature increases, the fraction of the molecules with “escape energy” increases • The higher the temperature, the faster the rate of evaporation
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Condensation • Some molecules of the vapor will lose energy
through molecular collisions • The result will be that some of the molecules will get captured back into the liquid when they collide with it • Also some may stick and gather together to form droplets of liquid particularly on surrounding surfaces
• We call this process condensation 71
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Evaporation vs. Condensation • Vaporization and condensation are opposite processes • In an open container, the vapor molecules generally spread
out faster than they can condense • The net result is that the rate of vaporization is greater than the rate of condensation, and there is a net loss of liquid • However, in a closed container, the vapor is not allowed to spread out indefinitely • The net result in a closed container is that at some time the rates of vaporization and condensation will be equal 72
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Effect of Intermolecular Attraction on Evaporation and Condensation
• The weaker the attractive forces between molecules,
the less energy they will need to vaporize • Also, weaker attractive forces means that more energy will need to be removed from the vapor molecules before they can condense • The net result will be more molecules in the vapor phase, and a liquid that evaporates faster – the weaker the attractive forces, the faster the rate of evaporation • Liquids that evaporate easily are said to be volatile e.g., gasoline, fingernail polish remover liquids that do not evaporate easily are called nonvolatile e.g., motor oil
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Energetics of Vaporization • When the high energy molecules are lost from the liquid, it lowers the average kinetic energy • If energy is not drawn back into the liquid, its temperature will decrease – therefore, vaporization is an endothermic process and condensation is an exothermic process
• Vaporization requires input of energy to
overcome the attractions between molecules
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Heat of Vaporization
• The amount of heat energy required to vaporize one mole of the liquid is called the heat of vaporization, ∆Hvap sometimes called the enthalpy of vaporization
• Always endothermic, therefore ∆Hvap is + • Somewhat temperature dependent •
∆Hcondensation = −∆Hvaporization
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Example: Calculate the mass of water that can be vaporized with 155 kJ of heat at 100 °C Given: Find: Conceptual Plan:
155 kJ g H2O kJ
mol H2O
g H2O
Relationships: 1 mol H2O = 40.7 kJ, 1 mol = 18.02 g Solution:
Check: because the given amount of heat is almost 4x the ∆Hvap, the amount of water makes sense 76
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Practice – Calculate the amount of heat needed to vaporize 90.0 g of C3H7OH at its boiling point (∆Hvap = 39.9 kJ/mol)
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Practice – Calculate the amount of heat needed to vaporize 90.0 g of C3H7OH at its boiling point Given: Find: Conceptual Plan:
90.0 g kJ g
mol
kJ
Relationships: 1 mol C3H7OH = 39.9 kJ, 1 mol = 60.09 g Solution:
Check:
because the given amount of C3H7OH is more than 1 mole the amount of heat makes sense 78
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Dynamic Equilibrium
• In a closed container, once the rates of vaporization
and condensation are equal, the total amount of vapor and liquid will not change • Evaporation and condensation are still occurring, but because they are opposite processes, there is no net gain or loss of either vapor or liquid • When two opposite processes reach the same rate so that there is no gain or loss of material, we call it a dynamic equilibrium this does not mean there are equal amounts of vapor and liquid – it means that they are changing by equal amounts 79
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Dynamic Equilibrium
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Vapor Pressure
• The pressure exerted by the vapor when it is in
dynamic equilibrium with its liquid is called the vapor pressure
remember using Dalton’s Law of Partial Pressures to account for the pressure of the water vapor when collecting gases by water displacement?
• The weaker the attractive forces between the
molecules, the more molecules will be in the vapor • Therefore, the weaker the attractive forces, the higher the vapor pressure
the higher the vapor pressure, the more volatile the liquid 81
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Dynamic Equilibrium • A system in dynamic equilibrium can respond to changes in the conditions • When conditions change, the system shifts its position to relieve or reduce the effects of the change
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Changing the Container’s Volume Disturbs the Equilibrium
Initially, the rate of vaporization and condensation are equal and the system is in dynamic equilibrium
When the volume is increased, the rate of vaporization becomes faster than the rate of condensation 83
When the volume is decreased, the rate of vaporization becomes slower than the rate of condensation
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Vapor Pressure vs. Temperature • Increasing the temperature increases the number of molecules able to escape the liquid • The net result is that as the temperature increases, the vapor pressure increases • Small changes in temperature can make big changes in vapor pressure the rate of growth depends on strength of the intermolecular forces
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Boiling Point • When the temperature of a liquid reaches a point where its vapor pressure is the same as the external pressure, vapor bubbles can form anywhere in the liquid not just on the surface
• This phenomenon is what is called boiling and the temperature at which the vapor pressure = external pressure is the boiling point
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Boiling Point • The normal boiling point is the temperature at which the vapor pressure of the liquid = 1 atm • The lower the external pressure, the lower the boiling point of the liquid
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Vapor Pressure Curves
normal BP 100 °C
760 mmHg
BP Ethanol at 500 mmHg 68.1°C 87
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Practice – Which of the following has the highest normal boiling point? a) b) c) d) e)
water TiCl4 ether ethanol acetone
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Practice – Which of the following is the most volatile? a) b) c) d) e)
water TiCl4 ether ethanol acetone
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Practice – Which of the following has the strongest Intermolecular attractions? a) b) c) d) e)
water TiCl4 ether ethanol acetone
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Heating Curve of a Liquid
• As you heat a liquid, its
temperature increases linearly until it reaches the boiling point q = mass x Cs x ∆T
• Once the temperature
reaches the boiling point, all the added heat goes into boiling the liquid – the temperature stays constant • Once all the liquid has been turned into gas, the temperature can again start to rise 91
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Sublimation and Deposition
• Molecules in the solid have thermal energy that
allows them to vibrate • Surface molecules with sufficient energy may break free from the surface and become a gas – this process is called sublimation • The capturing of vapor molecules into a solid is called deposition • The solid and vapor phases exist in dynamic equilibrium in a closed container at temperatures below the melting point therefore, molecular solids have a vapor pressure
solid
sublimation deposition 92
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Sublimation
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Melting = Fusion • As a solid is heated, its temperature rises and the molecules vibrate more vigorously • Once the temperature reaches the melting point, the molecules have sufficient energy to overcome some of the attractions that hold them in position and the solid melts (or fuses) • The opposite of melting is freezing
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Heating Curve of a Solid • As you heat a solid, its
temperature increases linearly until it reaches the melting point q = mass x Cs x ∆T
• Once the temperature reaches •
the melting point, all the added heat goes into melting the solid – the temperature stays constant Once all the solid has been turned into liquid, the temperature can again start to rise ice/water will always have a temperature of 0 °C at 1 atm
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Energetics of Melting
• When the high energy molecules are lost from the solid, it lowers the average kinetic energy • If energy is not drawn back into the solid its temperature will decrease – therefore, melting is an endothermic process and freezing is an exothermic process
• Melting requires input of energy to overcome the attractions between molecules
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Heat of Fusion
• The amount of heat energy required to melt one mole of the solid is called the Heat of Fusion, ∆Hfus
sometimes called the enthalpy of fusion
• Always endothermic, therefore ∆Hfus is + • Somewhat temperature dependent • • •
∆Hcrystallization = −∆Hfusion
Generally much less than ∆Hvap
∆Hsublimation = ∆Hfusion + ∆Hvaporization
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Heats of Fusion and Vaporization
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Heating Curve of Water
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Segment 1 • Heating 1.00 mole of ice at −25.0 °C up to the •
melting point, 0.0 °C q = mass x Cs x ∆T mass of 1.00 mole of ice = 18.0 g Cs = 2.09 J/mol∙°C
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Segment 2 • Melting 1.00 mole of ice at the melting point, •
0.0 °C q = n∙∆Hfus n = 1.00 mole of ice ∆Hfus = 6.02 kJ/mol
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Segment 3 • Heating 1.00 mole of water at 0.0 °C up to the •
boiling point, 100.0 °C q = mass x Cs x ∆T mass of 1.00 mole of water = 18.0 g Cs = 2.09 J/mol∙°C
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Segment 4 • Boiling 1.00 mole of water at the boiling point, •
100.0 °C q = n∙∆Hvap n = 1.00 mole of ice ∆Hfus = 40.7 kJ/mol
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Segment 5 • Heating 1.00 mole of steam at 100.0 °C up to •
125.0 °C q = mass x Cs x ∆T mass of 1.00 mole of water = 18.0 g Cs = 2.01 J/mol∙°C
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Practice – How much heat, in kJ, is needed to raise the temperature of a 12.0 g benzene sample from −10.0 °C to 25.0 °C?
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Practice – How much heat is needed to raise the temperature of a 12.0 g benzene sample from −10.0 °C to 25.0 °C? kJ,°C, T2 = 5.5 °C), Given: 12.0 g benzene, seg 1 =(T0.2325 1 = −10.0 1.51 kJ, seg 3 = (T0.3978 kJ T2 = 25.0 °C) seg 2 = melting, 1 = 5.5 °C, kJ
Find: Conceptual Seg 231 Plan: Relationships:
g
mol J
kJ
∆Hfus 9.8 kJ/mol, 1 mol = 78.11 g, 1 kJ = 1000 J, q = m∙Cs∙∆T Cs,sol = 1.25 J/g°C, Cs,liq = 1.70 J/g°C
Solution:
Tro: Chemistry: A Molecular Approach, 2/e
106
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Water – An Extraordinary Substance
• Water is a liquid at room temperature
most molecular substances with similar molar masses are gases at room temperature e.g. NH3, CH4
due to H-bonding between molecules
• Water is an excellent solvent – dissolving many ionic and polar molecular substances
because of its large dipole moment even many small nonpolar molecules have some solubility in water e.g. O2, CO2
• Water has a very high specific heat for a molecular substance moderating effect on coastal climates
• Water expands when it freezes at a pressure of 1 atm
about 9% making ice less dense than liquid water 107
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Solids Properties & Structure
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Crystal Lattice
• When allowed to cool slowly, the particles in a
liquid will arrange themselves to give the maximum attractive forces therefore minimize the energy
• The result will generally be a crystalline solid • The arrangement of the particles in a crystalline
solid is called the crystal lattice • The smallest unit that shows the pattern of arrangement for all the particles is called the unit cell 109
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Unit Cells
• Unit cells are 3-dimensional
usually containing 2 or 3 layers of particles
• Unit cells are repeated over and over to give the macroscopic • • •
crystal structure of the solid Starting anywhere within the crystal results in the same unit cell Each particle in the unit cell is called a lattice point Lattice planes are planes connecting equivalent points in unit cells throughout the lattice
110
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7 Unit Cells c c
c b
b
b a Cubic a=b=c all 90°
a Tetragonal a=c
a Orthorhombic a≠b≠c all 90°
c
c
b a Hexagonal a=c
c
a b Monoclinic a≠b≠c 2 faces 90° c b
b
a Rhombohedral a=b=c no 90° 111
a Triclinic a≠b≠c no 90° Copyright 2011 Pearson Education, Inc.
Classifying Crystalline Solids • Crystalline solids are classified by the kinds of
particles found • Some of the categories are sub-classified by the kinds of attractive forces holding the particles together
112
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Classifying Crystalline Solids
• Molecular solids are solids whose composite
particles are molecules • Ionic solids are solids whose composite particles are ions • Atomic solids are solids whose composite particles are atoms
nonbonding atomic solids are held together by dispersion forces metallic atomic solids are held together by metallic bonds network covalent atomic solids are held together by covalent bonds 113
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114
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Molecular Solids • The lattice sites are occupied by molecules CO2, H2O, C12H22O11
• The molecules are held together by intermolecular attractive forces
dispersion forces, dipole–dipole attractions, and Hbonds
• Because the attractive forces are weak, they tend to have low melting points generally < 300 °C
115
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Ionic Solids
• Lattice sites occupied by ions • Held together by attractions between oppositely charged ions nondirectional therefore every cation attracts all anions around it, and vice-versa
116
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Nonbonding Atomic Solids • Noble gases in solid form • Solid held together by weak dispersion forces very low melting
• Tend to arrange atoms in closest-packed structure
either hexagonal cp or cubic cp maximizes attractive forces and minimizes energy 117
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Metallic Atomic Solids • Solid held together by metallic bonds
strength varies with sizes and charges of cations coulombic attractions
• Melting point varies • Mostly closest-packed arrangements of the lattice points cations
118
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Metallic Structure
119
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Metallic Bonding
• Metal atoms release their valence electrons • Metal cation “islands” fixed in a “sea” of mobile electrons
+
+ e-
+
+ e-
+
+
+ ee-
+ +
+ ee-
+
+ e-
+
e+
+
+ ee-
Tro: Chemistry: A Molecular Approach, 2/e
+ +
+ ee-
+ +
+ ee-
120
+ +
+ ee-
+ +
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Network Covalent Solids
• Atoms attached to their nearest neighbors by
covalent bonds • Because of the directionality of the covalent bonds, these do not tend to form closest-packed arrangements in the crystal • Because of the strength of the covalent bonds, these have very high melting points generally > 1000 °C
• Dimensionality of the network affects other physical properties
121
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The Diamond Structure: a 3-Dimensional Network
• The carbon atoms in a diamond each have four covalent bonds to surrounding atoms sp3 tetrahedral geometry
• This effectively makes each crystal one giant molecule held together by covalent bonds
you can follow a path of covalent bonds from any atom to every other atom
122
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Properties of Diamond
• Very high melting point, ~3800 °C
need to overcome some covalent bonds
• Very rigid
due to the directionality of the covalent bonds
• Very hard
due to the strong covalent bonds holding the atoms in position used as abrasives
• Electrical insulator • Thermal conductor best known
• Chemically very nonreactive 123
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The Graphite Structure: a 2-Dimensional Network
• In graphite, the carbon atoms in a sheet are covalently bonded together
forming six-member flat rings fused together similar to benzene bond length = 142 pm
sp2
each C has three sigma and one pi bond
trigonal-planar geometry each sheet a giant molecule
• The sheets are then stacked and held together by dispersion forces
sheets are 341 pm apart 124
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Properties of Graphite • Hexagonal crystals • High melting point, ~3800 °C
need to overcome some covalent bonding
• Slippery feel
because there are only dispersion forces holding the sheets together, they can slide past each other glide planes
lubricants
• Electrical conductor parallel to sheets
• Thermal insulator • Chemically very nonreactive 125
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Practice – Pick the solid in each pair with the highest melting point a)
KCl ionic
SCl2 molecular
b)
C(s, graphite) cov. network S8 molecular
c)
Kr atomic
K
d)
SrCl2 ionic
SiO2 (s, quartz) cov. network
126
metallic
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