Transcript Document

Non-ideal Gases
• Non-ideality naturally follows a consideration
of intermolecular forces since these, in part,
account for gas non-ideality. The next slide
reviews the kinetic theory assumptions related
to the ideal gas law. Let’s consider where these
assumptions might break down.
Kinetic Molecular Theory of Gases
• Particles are point masses in constant,
random, straight line motion.
• Particles are separated by great
distances.
• Collisions are rapid and elastic.
• No force between particles.
• Total energy remains constant.
Visualizing Molecular Motion
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General Chemistry: Chapter 6
Slide 2 of 41
Aside: rms velocities
• Ex. Calculate the root mean square velocity for
gaseous CH2F2 molecules at -15 oC. Would the
rms velocity of CO2 molecules be larger or
smaller at the same temperature?
Non-ideal Gases:
• Treating gaseous molecules as point masses
(zero molecular volume) is reasonable for
small molecules and dilute gases. As molecular
size increases the gas molecules less resemble
point masses. As molar volume decreases
(high P and low T) the volume occupied by
molecules is no longer negligible compared to
the overall gas volume.
Non-ideal Gases – cont’d:
• Intermolecular forces vary significantly from
molecule to molecule – ranging from the very
weak London dispersion forces (in He and
CO2) to the very strong hydrogen bonding
interactions in molecules such as H2O, NH3
and CH3OH. In some cases, hydrogen bonding
can cause molecules to dimerize in the gas
phase resulting in highly non-ideal behaviour.
•
FIGURE 12-9
CH3-COOH
An acetic acid dimer
General Chemistry: Chapter 12
Slide 6 of 61
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Non-ideality and the Physical World
• As mentioned (repeatedly?), intermolecular
forces are responsible for molecules forming
condensed phases – without which the
universe would be a boring place! What
intermolecular forces are important in the next
couple of slides?
A Beautiful Picture – Thanks to Gas Non-ideality?
“Gotta love that
hydrogen bonding!”
Class examples:
• 1. Which of the following gases would you
expect to behave more ideally? Why?
(a) Ne(g) or CO2(g)
(b) CH4(g) or CH2Cl2(g)
(c) CH3CH2F(g) or CH3CH2OH(g)
• Note that the molecules which behave least
ideally also have the highest melting points
and the highest boiling points.
Class examples:
• 2. Which of the following gas mixtures would
you expect to behave more ideally? Why?
(a) Ar(g) and CH4(g) or (b) CH3OH(g) and
CH3NH2(g). Indicate what types of interactions
are possible for both mixtures.
Manifestation of Nonideality
• Gas nonideality manifests itself as molar
volumes decrease due to external P increase or
a T decrease. Initially the ratio PV/nRT is
smaller than the ideal value of 1 due to
intermolecular forces. At very high pressures
(and/or low T’s) the volume occupied by the
gas molecules collectively is not negligible
compared to the “empty space” in the gas and
PV/nRT gets bigger than 1.
6-9 Nonideal (Real) Gases
Compressibility factor
PV/nRT = 1 for ideal gas.
Deviations for real gases.
PV/nRT > 1 - molecular
volume is significant.
PV/nRT < 1 – intermolecular
forces of attraction.
Figure 6-22
Intermolecular forces of attraction
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General Chemistry: Chapter 6
Slide 13 of 41
van der Waals Equation
n2a
P+
V2
V – nb
= nRT
The van der Waals equation reproduces the observed behavior of gases
with moderate accuracy. It is most
accurate for gases comprising approximately spherical molecules that have
small dipole moments.
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General Chemistry: Chapter 6
Slide 14 of 41
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General Chemistry: Chapter 1
Slide 15 of 41
Solutes, Solvents and Solutions:
• In Newfoundland and Labrador one can imagine
that the following process might have been
important at some time.
• NaCl(s) + Water →NaCl(aq) (Na+(aq) + Cl-(aq))
• For this process to occur the ionic bonds in the
NaCl(s) lattice must be broken, the solute
particles (Na+ and Cl- ions) must be separated,
solvent molecules (water) must be separated and,
finally, a solution is formed.
Formation of Ionic Solutions
FIGURE 13-6
•An ionic crystal dissolving in water
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General Chemistry: Chapter 13
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Enthalpies involved in the solution
process:
NaCl(s) → Na+(g) + Cl-(g)
ΔHlattice > 0
Na+(g) + xs H2O(l) → Na+(aq)
ΔHhydration < 0
Cl-(g) + xs H2O(l) → Cl-(aq)
ΔHhydration < 0
ΔHsoln > 0 (~ 5 kcal/mol) but ΔGsolution < 0
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General Chemistry: Chapter 13
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Water of Hydration:
• When ions move through water they are
commonly surrounded by several water
molecules. This is true, in particular, of the
smaller and higher charged metal ions. The
combination of transition metal ions and water
molecules gives us beautiful colours (blue for
hydrated Cu2+). Water molecules commonly
remain when ionic salts are crystallized – eg.
CuSO4∙5H2O.
Solution Process – Energetics:
• The energetics of the solution process are
discussed on the next slide. The overall
solution process/change can be exothermic or
endothermic. It’s easy to rationalize solutions
forming where the intermolecular forces in the
solute molecules (before mixing with solvent)
are similar in type and magnitude to the
intermolecular forces between solvent
molecules.
Intermolecular Forces and the Solution
Process
FIGURE 13-2
•Enthalpy diagram for solution formation
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General Chemistry: Chapter 13
Slide 21 of 46
Intermolecular Forces in Mixtures
Magnitude of ΔHa, ΔHb, and ΔHc
depend on intermolecular forces.
Ideal solution
Forces are similar between all
combinations of components.
ΔHsoln = 0
FIGURE 13-3
•Intermolecular forces in a solution
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General Chemistry: Chapter 13
Slide 22 of 46
Hydrocarbons – CXHY:
• Most hydrocarbon molecules are nonpolar or
very weakly polar. Dispersion forces and very
weak dipole-dipole forces hold hydrocarbons
together in condensed phases. It’s not
surprising that different hydrocarbon
molecules easily mix to form nearly “ideal”
solutions (ΔHSolution approximately zero).
FIGURE 13-4
Two components of a nearly ideal solution
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General Chemistry: Chapter 13
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“Like Dissolves Like”:
• It seems reasonable that highly polar
molecules – such as CH2Cl2 and CH2Br2 –
might happily mix together (form a solution).
This is true.
• One might expect that methanol (CH3OH) and
methyl amine (CH3NH2) might also form a
solution since both molecules can hydrogen
bond. This is also true.
More Complex Examples:
• Water also “happily” forms solutions with
CH3OH and CH3CH2OH but, more complex
alcohols, such as CH3(CH2)6OH are only
sparingly soluble in water. Why?
Solubilities of Alcohols and Alkanes
Alcohol
Alcohol
Solubility (g/L)
Alkane
(Hydrocarbon)
Alkane Solubility
(g/L)
CH3OH
Miscible
CH3CH3
CH3CH2OH
Miscible
CH3CH2CH3
CH3CH2CH2OH
Miscible
CH3CH2CH2CH3
0.061
CH3(CH2)3OH
73
CH3(CH2)3CH3
0.040
CH3(CH2)4OH
22
CH3(CH2)4CH3
0.0095
CH3(CH2)5OH
5.9
CH3(CH2)6OH
Boiling Points of Alcohols and Alkanes
Alcohol
Alcohol Boiling
Point
Alkane
(Hydrocarbon)
Alkane Boiling
Point
Boiling Point
Difference
CH3OH
66 oC
CH3CH3
-88 oC
154 oC
CH3CH2OH
78 oC
CH3CH2CH3
-43 oC
121 oC
CH3CH2CH2OH
97 oC
CH3CH2CH2CH3
-1 oC
98 oC
CH3(CH2)3OH
118 oC
CH3(CH2)3CH3
36 oC
82oC
CH3(CH2)4OH
137 oC
CH3(CH2)4CH3
69 oC
68oC
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13 - 29
Gas Law Review
• Ex. A sample of O2 gas initially present in a
2.00 L container at a pressure of 2.40 atm and
22 oC is moved to a 5.00 L container at 15 OC.
A sample of Ne gas initially held at a pressure
of 4.00 atm and -16 OC in a 4.00 L container is
added to the oxygen gas in the 5.00 L
container (still at 15 OC). What is the total
pressure in the 5.00 L container?