Chem PPT 5b & 16a - Real Gases Deviations from Ideality Ideal gases behave according to the gas laws Ideal Ideal gases obey the Kinetic-Molecular

Chem PPT 5b & 16a - Real Gases Deviations from Ideality...

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Unformatted text preview: Real Gases: Deviations from Ideality Ideal gases behave according to the gas laws. Ideal Ideal gases obey the Kinetic-Molecular theory. Ideal KineticMost of a gas sample is empty space Most Ideal gases are not attracted to each other because Ideal they are very far apart from each other. Chapter 5b & 16a Real Gases & Liquids 2 Real Gases: Deviations from Ideality Real Gases: Deviations from Ideality Real gases behave ideally at ordinary temperatures Real and pressures. At low temperatures and high pressures real gases do At not behave ideally. The reasons for the deviations from ideality are: The The molecules are very close to one another, thus The their volume is important. The molecular interactions also become important. The Real Gases: gases that are close to becoming liquids, Real do not behave according to the Kinetic-Molecular Kinetictheory. Real Gases are close to becoming liquids at: Real Extreme high pressures Extreme Low temperatures (very low) Low High pressures and low temperatures High 3 Real Gases: Deviations from Ideality 4 Real Gases: Deviations from Ideality Johannes van der Waals studied real gases and Johannes adjusted the Ideal Gas Equation to account for the behavior of gases under the conditions of extreme pressure and very low temperatures. Very high pressures compress gases. Very Very low temperatures cause the kinetic energy of Very gases to decrease. The molecules move so slowly that small attractive The forces between molecules become significant. The gases have less volume in the container in which to The move. 5 Very high pressures and very low temperatures Very cause gases to move slower and closer to each other in the container. 6 1 Real Gases: Deviations from Ideality van der Waals’ equation accounts for the behavior of van Waals’ real gases at low temperatures and high pressures. n 2a P + 2 ( V − nb ) = nRT V The van der Waals constants a and b take into account two things: 1. a accounts for intermolecular attraction 2. b accounts for volume of gas molecules At large volumes a and b are relatively small and van der Waal’s equation reduces to ideal gas law at high temperatures and low pressures. 7 For “a” (attractive forces): • the monoatomic molecules (noble gases) have smaller values of a. • the diatomic molecules and molecules with more than one kind of atom have larger values of a—polar molecules, such as ammonia, have even greater values of a. For “b” (volume taken by gas molecules): • Larger molecules take up more space, therefore; have a larger value of b. 8 Real Gases: Deviations from Ideality What are the intermolecular forces in gases that cause them to deviate from ideality? 1. For nonpolar gases the attractive forces are London Forces 2. For polar gases the attractive forces are dipoledipoledipole attractions or hydrogen bonds. 9 Real Gases: Deviations from Ideality An ideal gas has no intermolecular forces between An gas particles. The strong forces that can exist between molecules The are determined by intermolecular forces (IMF). The stronger the IMF forces, the gas behaves less The ideally. Polar molecules have the greatest IMF forces. Polar Even IMF forces between nonpolar molecules can Even become strong – this depends on the size of the molecule. 10 End of Chapter 4 Gases are the simplest state of matter. Gases Liquids and solids are more complex. Liquids They are the subject of Chapter 16. They 12 2 Kinetic-Molecular Description Kineticof Liquids and Solids CHAPTER 16a– Liquids and 16a– Solids LIQUIDS LIQUIDS Schematic representation of the three common states of Schematic matter. gas cool liquid heat cool solid heat 14 13 Kinetic-Molecular Description Kineticof Liquids and Solids Kinetic-Molecular Description Kineticof Liquids and Solids If we compare the strengths of interactions among particles If strengths and the degree of ordering of particles, we see that degree Gases < Liquids < Solids Miscible liquids are soluble in each other. Miscible Immiscible liquids are insoluble in each other. Immiscible Two examples of immiscible liquids: Two Water does not dissolve in oil. Water Water does not dissolve in cyclohexane. Water cyclohexane. Examples of miscible liquids: Examples Water dissolves in alcohol. Water Polar solvent and polar solute Polar Gasoline dissolves in motor oil. Gasoline Nonpolar solvent and nonpolar solute Nonpolar 15 16 Intermolecular Forces in Liquids Intermolecular Forces in Liquids For a liquid in an open container to boil, its vapor For pressure must reach atmospheric pressure. For some substances this occurs at temperatures For well below zero; for example, oxygen has a boiling point of –183°C. 183° Other substances do not boil until the Other temperature is much higher. An explanation for this variation involves a An consideration of the nature of the intermolecular forces that must be overcome in order for molecules to escape from the liquid state into the vapor state. Mercury, for example, has a boiling point of 357 °C, Mercury, which is 540°C higher than that of oxygen. 540° 17 18 3 Intermolecular Forces in Liquids Intermolecular Forces in Liquids The changes of state are caused by intermolecular The interactions. During a change of state, the molecules remain During intact. Intermolecular forces are the forces between Intermolecular molecules and formula units. Much weaker attractions that occur between Much molecule-molecule of a covalent compound moleculeand ion-ion in a formula unit. ionIntramolecular attractions are the forces Intramolecular between atoms in molecules and formula units Covalent & Ionic Covalent Gaseous water is composed of H2O molecules in a Gaseous vapor state with high kinetic energies. Liquid water is composed of H2O molecules that are Liquid held together by intermolecular forces with low kinetic energies. Solid water is composed of H2O molecules that are held Solid together by intermolecular forces with very low kinetic energies. 19 20 Intermolecular Forces in Liquids Intermolecular Forces in Liquids Fig. 5.9 Intermolecular forces are similar in one way to intramolecular forces (between atoms) involved in covalent bonding. They are electrostatic in origin. They A major difference between inter- and intramolecular major interforces is their magnitude; intermolecular forces are much weaker. Intermolecular forces Intramolecular forces Intermolecular forces can be several magnitudes stronger in Intermolecular strength than intramolecular forces However, intermolecular forces, despite their relative However, weakness, are sufficiently strong to influence the behavior of liquids, often in a very dramatic way. 22 21 Intermolecular Attractions and Phase Changes (Ion-Ion) (Ion- Intermolecular Forces in Liquids 1. Ion-ion interactions (interionic) Ion(interionic) The force of attraction between two oppositely The charged ions is governed by Coulomb’s law. Coulomb’ There are five important intermolecular attractions. There They will be presented from strongest attraction to They the weakest attraction. ( q )( q ) + Ep ∝ - r E p is the potential energy q + and q - are the ion charges. r is the distance between the ions. 23 24 4 Intermolecular Attractions and Phase Changes Intermolecular Attractions and Phase Changes Coulomb’s law determines: Coulomb’ 1. The melting and boiling points of ionic 1. compounds. Example: Arrange the following ionic compounds in the expected order of increasing melting and boiling points. NaF, CaO, CaF2 compounds with small highly charge ions have the compounds highest b.p. and m.p. b.p. m.p. 2. The solubility of ionic compounds. 25 Intermolecular Attractions and Phase Changes 26 Intermolecular Attractions and Phase Changes The strength of the intermolecular forces between The ion-ion interactions determine the solubility of ionic ioncompounds in water. The greater the force (coulomb attraction) between The ions in a salt crystal the more difficult it is to overcome during hydration. 2. Hydrogen bonding a special type of dipole-dipole dipoleintermolecular bonding. One way to determine the strength of One intermolecular forces is to look at the boiling points of different substances. As an example, we will look at the b.p. of As b.p. substances that contain hydrogen. (Called hydrides) 27 Boiling points of molecular hydrides vs. molar mass Hydrogen Bonds The three hydrides with the highest boiling points are: H2O HF Increased molar mass increases b.p. 28 NH3 The increased boiling points are due to the differences in electronegativity between the H and O, F, and N atoms. 29 Unusually strong dipole-dipole interactions are found dipoleamong hydrogen-containing molecules: hydrogenWhere hydrogen is covalently bonded to a highly Where electronegative element of small atomic size. The three small atoms bonded to hydrogen are The fluorine, oxygen, and nitrogen. fluorine, 30 5 Hydrogen Bonds Hydrogen Bonds Two factors account for the extra strength of these Two dipole-dipole interactions. dipole1. The highly electronegative element to which hydrogen is covalently bonded dominates the electron-sharing process electronto such a degree that the hydrogen atom is left with significant partial positive charge. ∆EN = 1.9 1.4 0.9 The hydrogen atom is essentially a “bare” nucleus, The bare” because it has no electrons besides the one attracted to the electronegative element—a unique property element— of hydrogen. 2. The small size of the hydrogen atom allows the "bare" nucleus to approach closely and be strongly attracted to an unshared pair of electrons on the electronegative atom of another molecule. 31 Hydrogen Bonds 32 Hydrogen Bonds Dipole-dipole interactions of the type we are now Dipoleconsidering are given a special name; hydrogen bonds (H-bonds). (HA hydrogen bond is an extra strong dipole-dipole hydrogen dipoleinteraction involving a hydrogen atom covalently bonded to a small, very hydrogen electronegative atom (F, O, or N). (F, and an unshared pair of electrons on another small, very and electronegative atom (F, O, or N). (F, The two molecules that participate in a hydrogen bond need not be identical. Hydrogen bond formation is possible whenever two Hydrogen molecules, the same or different, have the following characteristics. 1. One molecule has a hydrogen atom attached by a covalent bond to an atom of nitrogen, oxygen, or fluorine. nitrogen, fluorine. 2. The other molecule has a nitrogen, oxygen, or fluorine atom present that possesses one or more nonbonding electron pairs. 33 Hydrogen Bonds HF + HF HF + H2O H2O + HF 34 Hydrogen Bonding NH3 + NH3 H2O a very polar molecule. The oxygen atom is more electronegative than the The bonded hydrogen atoms. The oxygen nucleus has very little shielding and is The able, therefore, to pull the electrons away from the hydrogen in the bonds. The very small hydrogen atoms with their partial The positive charges are able to get very close to a lone pair of electrons on an oxygen in a neighboring water molecule. NH3 + H2O H2O + NH3 35 36 6 Hydrogen Bonding Intermolecular Attractions and Phase Changes δ+H •• δ− O: δ+H δ+H •• δ− O: δ+H The lone pair on oxygen and the partial charge on hydrogen are strongly attracted. The H-bonding is the strongest when the H atom is on a straight line between two oxygen atoms. 37 38 H-bonding is strong enough to maintain the intermolecular interactions even as vapors. Hydrogen Bonds H-bonding between H2O molecules. Liquid hydrogen fluoride contains chains of H-F molecules that zig-zag. The vapor has short fragments of the chains. CH3COOH (acetic acid) vapor has dimers (pairs of molecules) The H-bonding occurs between the oxygen and hydrogen in the molecule. 39 40 Hydrogen Bonds Hydrogen Bonds 1. 2. Indicate whether hydrogen bonding is possible between two molecules of each of the following substances. 41 Indicate whether hydrogen bonding is possible between two molecules of each of the following substances. 42 7 Consequences of Hydrogen Bonding Consequences of Hydrogen Bonding The vapor pressures of liquids that have significant The hydrogen bonding are much lower than those of similar liquids wherein little or no hydrogen bonding occurs. Vapor pressure is the pressure exerted by a vapor above a Vapor liquid when the liquid and vapor are in equilibrium. This is because the presence of hydrogen bonds makes it This more difficult for molecules to escape from the condensed state; additional energy's needed to overcome the hydrogen bonds. For this reason, boiling points are much higher for For liquids in which hydrogen bonding occurs. Water is one of the few substances known in which the solid Water form (ice) is less dense than the liquid form. This results from water expanding as it freezes. This This "expansion phenomenon" is directly related to hydrogen bonding. In the solid state, hydrogen bonds, which must have In specific orientations, that is, have directional nature, cause the water molecules to be farther apart than they are in the liquid state. In the liquid state hydrogen bonds are continually being In formed and broken; in the solid state hydrogen bonds are "permanent'”. "permanent'” 43 44 Intermolecular Attractions and Phase Changes Consequences of Hydrogen Bonding Because ice is less dense than liquid water, lakes Because freeze from top to bottom and not vice versa. If this "density inversion" for water did not exist If (no hydrogen bonding), surface ice would sink to the bottom of a lake as it was formed, and water bodies in colder climates would freeze solid in winter. 3. Ion-Dipole Forces IonA species (such as a salt) that ionizes in water: species The ions will have water molecules attached The The attachment of water molecules to solutes particles The (particularly ions) is called hydration. hydration 45 Hydration occurs because of the polar character of H2O molecules. δ _ Intermolecular Attractions and Phase Changes Attracted to cations Water molecules cluster around a cation. Water cation. Oxygen atoms point inward to the positive charge on the Oxygen cation Hydrogen atoms point outward away from the positive Hydrogen charge. Attracted to anions δ + 46 δ Water molecules cluster around a anion. Water + Attracted to anions 47 Hydrogen atoms point inward to the negative charge on the Hydrogen anion Oxygen atoms point outward away from the negative Oxygen charge. 48 8 Hydration Intermolecular Attractions and Phase Changes δ+ δ+ δ− δ− + δ+ δ+ δ− − δ+ δ− The interaction of the partial charges on water with The the full charges on ions is a very good example of ion-dipole interactions. ionIon-dipole interactions are characteristic of ions Iondissolved in polar solvents δ+ δ+ δ+ δ+ δ+ Water (H2O) Water Liquid ammonia (NH3) Liquid δ+ Anion surrounded by water molecules Cation surrounded by water molecules 49 Intermolecular Attractions and Phase Changes Intermolecular Attractions and Phase Changes Polar molecules need to be very close to an ion Polar (almost in contact) before the interaction is significant. The ions in an aqueous solution may retain some of The the hydration water molecules when certain salts are crystallized from an aqueous solution. Na2CO3•10 H2O and CuSO4•5 H2O The size and charge of the cations determine to what The extent hydration will occur. The interaction between an ion and a dipole depends more The strongly on the distance than the interaction. Ep = − zµ r2 potential energy Potential energy is lowered by interaction 50 z is the charge of the ion µ is the electric dipole moment of the polar molecule 51 Intermolecular Attractions and Phase Changes 52 Intermolecular Attractions and Phase Changes Ba2+ and K+ both have similar radii. Ba Small ions attract water molecules more strongly. Small Small cations are usually experience more hydration Small Li+ and Na+ usually form hydrated salts Li+ K+, Rb+, and Cs+ usually do not form hydrated K+, Rb+, salts. Potassium salts are anhydrous (water free) Potassium Barium salts are often hydrated Barium The difference is the charge on barium The Ion-dipole interactions are strong for small, highly Ionsmall, charged ions. That is the reason that small, highly charged cations are That often hydrated in compounds. 53 54 9 Dipole-Dipole Interactions Dipole- Dipole-Dipole Interactions Dipole- 4. Dipole-dipole interactions are electrostatic Dipole- When polar molecules approach each other, they tend When to line up so that the relatively positive end of one molecule is directed toward the relatively negative end of the other molecule. attractions between polar molecules. Polar molecules (which are often called dipoles), Polar it should he recalled, are electrically electrically asymmetrical. Also known as permanent dipole-dipole Also permanent dipoleinteractions. As a result, there is an electrostatic attraction between the As molecules. The greater the polarity of the molecules, the greater the The strength of the dipole-dipole interaction. dipole- 55 Dipole-Dipole Interactions Dipole- Dipole-Dipole Interactions Dipole- Dipole-dipole interactions that occur in the solid phase are called Dipolestationary dipoles. µ1 and µ2 are dipole moments of interacting molecules where µ1 = µ2 The greater the polarity of the molecules, the stronger the The interactions. Doubling the distance between molecules reduces the strength Doubling of their interaction by a factor of 8 (23 = 8) Ep = − 56 Dipole-dipole interactions that occur in a liquid phase Dipoleare called rotating dipoles. Doubling the distance between molecules reduces the strength Doubling of their interaction by a factor of 64 (26 = 64) 64) Ep = − µ1µ 2 µ1µ 2 r6 r3 57 Dipole-Dipole Interactions 58 Dipole-Dipole Interactions Dipole-dipole interactions – Consider BrF a polar molecule. The figure shows the many dipole-dipole interactions possible for a random arrangement of polar CIF molecules. A represents attractive forces R represents the repulsive forces 59 60 10 Intermolecular Attractions and Phase Changes (London Dispersion) 5. London Forces (Dispersion Forces) The last type of intermolecular force, and the weakest, is the London (dispersion) force, named London after the German physicist Fritz London (1900(19001954), who first postulated its existence. They are also known as: They Dispersion forces (very common term) Dispersion van der Waal forces van London Forces are very weak. London They are the weakest of the intermolecular forces. They This is the only attractive force in nonpolar This molecules. 61 62 London Forces (Dispersion Forces) London Forces (Dispersion Forces) London forces are instantaneous dipole-dipole London instantaneous dipoleinteractions that exist between all atoms and molecules, nonpolar as well as polar. molecules, The origin of London forces is more difficult to The visualize than that of dipole-dipole interactions dipole- London forces result from momentary (temporary) London uneven electron distributions in molecules. Most of the time the electrons can be visualized as being Most distributed in a molecule in a definite pattern determined by their energies and the electronegativity of the atoms. However, there is a small statistical chance (probability) However, that the electrons will deviate from their normal pattern. 63 δ– 64 δ+ London Dispersion Forces o In a group of Ar atoms the temporary dipole in one atom induces other atomic dipoles. o For example, in the case of a nonpolar The induced polar molecule and diatomic molecule (A), more electron all the molecules with induced density may temporarily be located on one polarity are then attracted to side of the molecule than on the other. each other. o This condition causes the molecule to This happens many, many times become polar for an instant. per second throughout the o The negative side of this instantaneously liquid, resulting in a net polar molecule will tend to repel electrons attractive force. of adjoining molecules (B) and cause these molecules to also become polar (an induced polarity). 65 66 11 London Forces (Dispersion Force...
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