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39 Cards in this Set
- Front
- Back
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4 types of intermolecular forces
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dispersion, dipole-dipole, hydrogen bonding, ion-dipole
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dspersion force
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present in all molecules and atoms, also called london force, result from fluctuations in electron distribution within an atom or molecule
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 larger molar mass = more electrons = larger electron cloud = increased polarizability = stronger attractions more surface-to-surface contact = larger induced dipole = stronger attraction |
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stronger bonds = boiling point, melting point?
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stronger bonds are caused by larger differences in electronegativity, the more polar molecule = the stronger the bond = higher boiling and melting points
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dipole-dipole attraction
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all polar atoms have 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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solubility
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Solubility depends on the attractive forces of solute and solvent molecules
Like dissolves Like miscible liquids will always dissolve in each other polar substance 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 becomes competition between parts |
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4 types of intermolecular forces
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dispersion, dipole-dipole, hydrogen bonding, ion-dipole
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dspersion force
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present in all molecules and atoms, also called london force, result from fluctuations in electron distribution within an atom or molecule
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 larger molar mass = more electrons = larger electron cloud = increased polarizability = stronger attractions more surface-to-surface contact = larger induced dipole = stronger attraction |
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stronger bonds = boiling point, melting point?
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stronger bonds are caused by larger differences in electronegativity, the more polar molecule = the stronger the bond = higher boiling and melting points
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dipole-dipole attraction
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all polar atoms have 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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solubility
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Solubility depends on the attractive forces of solute and solvent molecules
Like dissolves Like miscible liquids will always dissolve in each other polar substance 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 becomes competition between parts |
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hydrogen bonding
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When a very electronegative atom is bonded to hydrogen, it strongly pulls the bonding electrons toward it
O-H, N-H, or F-H Since hydrogen has no other electrons, when it loses the electrons, 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 |
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Ion-Dipole Attraction
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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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Summary
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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 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 |
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Surface Tension
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surface tension is a property of liquids that results from the tendency of liquids to minimize their surface area
in order to minimize their surface area, liquids form drops that are spherical as long as there is no gravity 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 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 at room temp, surface tension of H2O = 72.8 mJ/m2 |
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Factors Affecting Surface Tension
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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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surface tension vs. temperature
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as temperature rises, surface tension decreases
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Viscosity
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viscosity is the resistance of a liquid to flow
1 poise = 1 P = 1 g/cm∙s often given in centipoise, cP larger intermolecular attractions = larger viscosity higher temperature = lower viscosity |
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Vaporization
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molecules in the liquid are constantly in motion
the average kinetic energy is proportional to the temperature however, some molecules have more kinetic energy than the average if these 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
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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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Effect of Intermolecular Attraction on Evaporation and Condensation
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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
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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
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the amount of heat energy required to vaporize one mole of the liquid is called the Heat of Vaporization, DHvap
sometimes called the enthalpy of vaporization always endothermic, therefore DHvap is + somewhat temperature dependent DHcondensation = -DHvaporization |
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Example 11.3 – Calculate the mass of water that can be vaporized with 155 kJ of heat at 100°C
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given: 155kJ
find: hH2O 155kJ x (1mol/40.7kJ) x (18.02g/1mol) = gH2O |
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Vapor Pressure
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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 |
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Vapor Pressure vs. Temperature
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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
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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 required to have the vapor pressure = external pressure is the 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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Heating Curve of a Liquid
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as you heat a liquid, its temperature increases linearly until it reaches the boiling point
q = mass x Cs x DT 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 |
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The Critical Point
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the temperature required to produce a supercritical fluid is called the critical temperature
the pressure at the critical temperature is called the critical pressure at the critical temperature or higher temperatures, the gas cannot be condensed to a liquid, no matter how high the pressure gets |
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Sublimation and Deposition
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solid to gas = sublimation
gas to solid = deposition |
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Melting = Fusion
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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
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as you heat a solid, its temperature increases linearly until it reaches the melting point
q = mass x Cs x DT 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
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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
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the amount of heat energy required to melt one mole of the solid is called the Heat of Fusion, DHfus
sometimes called the enthalpy of fusion always endothermic, therefore DHfus is + somewhat temperature dependent DHcrystallization = -DHfusion generally much less than DHvap DHsublimation = DHfusion + DHvaporization |
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Heating Curve of Water
Segment 1: |
Segment 1: heating 1.00 mole of ice at -25.0°C up to the melting point, 0.0°C
q = (mass ice 18g) x (Cs ice 2.09 J/g dC) x (delta T) = J |
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Heating Curve of Water
Segment 2: |
melting 1.00 mole of ice at the melting point, 0.0°C
q = n∙DHfus q = (mol ice) x (6.02 kJ/mol) = kJ |
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Heating Curve of Water
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 DT q = (mass H2O 18g) x (Cs H2O 4.18 J/g dC) x (delta T) = J |
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Heating Curve of Water
Segment 4: |
boiling 1.00 mole of water at the boiling point, 100.0°C
q = n∙DHvap q = (mol H2O) x (40.7 kJ/mol) = kJ |
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Heating Curve of Water
Segment 5: |
heating 1.00 mole of steam at 100.0°C up to 125.0°C
q = mass x Cs x DT q = (mass H2O 18g) x (Cs H2O 2.01 J/g dC) x (delta T) = J |
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Phase Diagrams
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describe the different states and state changes that occur at various temperature - pressure conditions
areas represent states lines represent state changes liquid/gas line is vapor pressure curve both states exist simultaneously critical point is the furthest point on the vapor pressure curve triple point is the temperature/pressure condition where all three states exist simultaneously for most substances, freezing point increases as pressure increases top left = solid top right = liquid bottom = gas |
