Part I: Kinetic Molecular Theory, Diffusion and Effusion, Real Gases

CHM 102

Friday, June 29th

Summer II 2007

Kinetic Molecular Theory

• The ideal gas equations describe HOW a gas will behave as it does.

• The Kinetic Molecular Theory describes WHY a gas will behave as it does.

• So the physical model described by Kinetic Molecular Theory is summed up in a series of five postulates.

Kinetic Molecular Theory in 5 Steps

1. Gases are made of A LOT of molecules that are in continuous random linear motion.

2. The combined volume of all the molecules of gas is negligible compared to the volume of the container (gases are assumed to have no volume).

3. Gas molecules don’t attract or repel one another.

KMT, the last two postulates…4. When gas molecules collide, energy can

be transferred, but the average KE of all the molecules doesn’t change over time. (as long as the temperature remains constant).

5. From #4, we can pull that Avg. KE is proportional to absolute temperature. At any given T, the molecules of all gases (identity doesn’t matter) have the same average KE.

Let’s apply KMT!

• Let’s take a 1L container of gas, and double its volume at constant T.

• What happens to its average KE?

• What happens to the pressure?

• According to the ideal gas law, P is inversely proportional to V. So we’re raising V, so P should drop. That’s the how, what’s the why? (Hint: KMT)

Let’s apply KMT again!

• Let’s take a 1L container of gas, and double temperature at constant volume.

• What happens to its average KE?

• What happens to the pressure?

• According to the ideal gas law, P is directly proportional to T. So we’re raising T, so P should raise. That’s the how, what’s the why? (Hint: KMT)

Let’s apply KMT yet again!• Let’s look at two separate gases now: O2

and CO2.• When each of these gases is held

separately in two different containers at exactly 298.15 K, which has the higher average KE?

• Which gas molecule, O2 or CO2 moves at a faster velocity, on average?

[ Hint: KE = (1/2)(mass)(velocity)2 ]

Diffusion: d-limonene• Diffusion is what helps you smell an onion

when you slice into it. It helps you catch a whiff of a fresh cut orange. It can also make a date awkward after the extra spicy fajitas.

• Diffusion is the spread of of substance through a space (or another substance).

• D-limonene is a chemical that provides some of the aroma of oranges.

Effusion, on the other hand…• Notice that balloons sag and drop after a

few days of floating around? This is effusion in action.

• Effusion is the escape, or movement, of a gas through a tiny hole, such as a pore in a latex balloon.

• There is a law that can give the relative rates of effusion for two different gases based on their molecular weights.

Graham’s Law of Effusion

• This law can compare the relative rates (r1 and r2) of effusion (escape through a tiny hole) of two gases, when their molecular weights (M1 and M2 are known).

• WARNING: Very common mistake time! Make sure r1 is in the top when you calculate but M1 is in the bottom!

1

2

2

1

M

M

r

r

Practice with Effusion

• Hydrogen has 2 naturally occuring isotopes, 1H, and 2H. Chlorine also has 2 naturally occuring isotopes, 35Cl and 37Cl. How many different combinations can we make to make H-Cl? Put these in the correct order of rate of effusion from slowest to fastest!

More Practice with Effusion

• You’ve got an unknown homonuclear diatomic gas. O2 effuses 2.816 times faster than the unknown gas at the same temperature. Identify the unknown gas and its molecular weight.

Real gases• Real gases do not behave as perfect “ideal”

gases do.• KMT gives us clues as to why these do not.• Remember that in KMT, we assumed gases

have essentially NO volume. In reality, however small, gases do take up space.

• Also in KMT, we assumed that gases neither attracted nor repelled one another. In reality, they can, and more effectively under certain conditions.

When keepin’ it real goes wrong

• Statement: Gases behave most ideally or least ideally under conditions of high pressure and low volume.

• Is this true or false? How could we correct this? (if it needs correcting)

• Let’s look at a model of this here.

Part II: Molecular Comparison of Liquids and Solids, Intermolecular Forces (IMF)

CHM 102

Friday, June 29th

Summer II 2007

States of Matter: Gases

• Gases assume the shape and volume of their container.

• You can compress them (i.e. scuba tanks)

• They flow very easily (i.e. wind)

• Diffusion of gas is very quick (if not, Beano and Glade would be out of business)

States of Matter: Liquids

• Liquids do assume the shape of the container, but will not expand to fill its container as gases do. (drink half your coke and it does not fill back up).

• Liquids are not compressible

• Liquids flow readily (i.e. rivers, etc.)

• Diffusion within a liquid occurs slowly

DEMO: FOOD DYE

States of Matter: Solids

• Solids retain their shape and volume, regardless of the container.

• They essentially can’t be compressed.

• They do not flow. (your chair is a solid, and is quite apparently not flowing).

• Diffusion within a solid occurs extremely slowly.

States of Matter and avg KE

• The average KE of gases is much higher than the attractive forces holding them together.

• The average KE of liquids is strong enough to keep liquids moving, but the attractive forces are strong enough to hold the liquid together.

• In solids, the attractive forces are sufficiently strong to hold it tightly arranged, with little or no flow.

States of Matter and Order

• Solids are less disordered than liquids which are less disordered than gases (which are in complete disarray).

• This ties in directly to space between particles and how tightly the molecules are held together. Gases are mostly empty space, while liquids have much less empty space, with solids having the least empty space.

Intermolecular Forces• Intermolecular forces are forces between

molecules.

• Covalent bonds and ionic bonds tend to be much stronger than IMF’s.

• For example, to bring HCl(l) to HCl(g) (vaporization), we’re breaking IMF interactions, and it takes 16 kJ/mol. To break the polar covalent bond between H and Cl , it takes 431 kJ/mol.

Phase Changes and IMF’s

• What we’re looking at in phase changes is not molecules undergoing chemical changes (i.e. covalent and ionic bonds breaking and forming).

• We’re looking at physical changes. Phase changes involve overcoming IMF interactions by supplying or removing enough energy from a system.

• The stronger the IMF, the higher the BP, the lower the FP.

The 3 General Intermolecular Forcesbetween Neutral Molecules

1. Dipole-Dipole Interactions- This IMF forms between polar molecules.

2. London Dispersion Forces- This IMF is the primary IMF between nonpolar molecules, and the weakest of the three.

3. Hydrogen Bonding Interactions- This IMF is the strongest interaction of the three, and will be discussed later.

Dipole-Dipole Forces

• Opposites attract (MC Skat Cat and Paula Abdul…anyone remember this?).

• Neutral polar molecules, (i.e. H-Cl) will form dipole-dipole interactions as shown below, with the positive pole on one molecule attracting the negative pole on the other.

London Dispersion Forces

• LDF is a force present in all molecules to some extent, but it is very weak compared to the other IMF’s discussed.

• LDF is the only IMF for nonpolar molecules.

• London Dispersion Force is sometimes called an “induced” dipole. It is a dipole caused by the shifting of the electron cloud to form a momentary dipole that can interact with another momentary dipole.

LDF: Leaving Damn Fast!

• London dispersion forces are constantly leaving and reforming due to the trainsient nature of the IMF. As the electron cloud around an atom/molecule shifts to produce an instantaneous dipole, it may induce a dipole in another atom/molecule.

• The strength of LDF deals directly with polarizability, or “squishiness” of the electron cloud.

A trend for LDF

• Stronger LDF is associated with higher degrees of polarizability.

• Higher polarizability is associated with larger, higher MW molecules.

• It follows that: As molecular weight increases for nonpolar molecules, so do the London Dispersion forces holding them together!

Practice: Comparing IMF Strength

• Compare the H-Cl, the Br-Br, and the CH4 molecules. Order them from weakest to strongest IMF holding them together. Why did we pick the order we did?

• First, look at polarity. H-Cl is a polar molecule, and the remaining two are nonpolar. So we know that H-Cl has dipole-dipole IMF, and the other two have LDF. So H-Cl is the strongest. Between the remaining two, Br2 is the largest MW, so it has the strongest LDF. The order is: CH4<Br2<H-Cl.