Reading Notes for Chapter 8


These are Dr. Bodwin's reading notes for Chapter 8 of "Introduction to Chemistry". I am using a local .pdf copy that was downloaded in August 2020.

Chapter Summary:

There are a number of different ways we can define and describe states of matter, from a very simple macroscopic description (defined or variable shape and volume) to extremely theoretical and mathematical models used by computational chemists. We need to be somewhere in the middle so that we can not only classify matter, but also understand how it behaves under various conditions.

Intermolecular Forces (IMFs):

Molecules interact with one another in a variety of ways.
The strength of these various interactions on the microscale has a huge impact on the macroscale properties of a substance.
Covalent Network Interactions (very strong) - Covalent network solids are very stable because covalent bonds themselve are typically very stable. Covalent bonds are what hold molecules together. In fact, covalent network substances can actually be thought of a a single molecule!
Ionic Interactions (very strong, more variable) - These substances are goverend by Coulomb's Law. Opposite charges attract each other, like charges repel each other. The strength of the attraction and repulsion depend upon the magnitude of the charge (+2 attracts/repels more strongly than +1) and the distance between the particles (closer together = stronger interactions).
Polar Interactions (moderate to very strong) - Molecules are held together by covalent bonds in which electrons are shared. That sharing is not always equal; one "side" of the bond usually attracts the electrons more strongly than the other "side". The side that attracts electrons more strongly will appear slightly more negatively charged than the other side. The slightly negative end of one molecule is attracted to the slightly positive end of another molecule just like in an ionic interaction, but (usually) not as strong. The positive and negative ends are called "poles" (just like the North Pole and South Pole of the Earth), to these are called "dipoles" because there are 2 poles. Dipole-dipole interactions can be quite strong or quite weak, depending upon how strongly the electrons are drawn to one side of the covalent bond.
Hydrogen Bonds - these are often described as a separate category, but hydrogen bonds are really just a specific type of dipole-dipole interaction. When hydrogen is bonded to oxygen, nitrogen, fluorine, sulfur, chlorine, bromine, or iodine, the hydrogen does not attract the electrons as strongly as the other element, so the resulting bond is polar. There are a LOT of O-H, N-H, S-H bonds in biology, so it makes sense that this would be an important type of bond to consider.
Dispersion Forces (weak, individually) - This one has a LOT of names. "Dispersion forces", "London forces", "London dispersion forces", and "induced dipole forces"... they all mean the same thing. What are they? When a molecule has a "squishy" electron cloud, at some moments, the electrons squish toward one end of the molecule. When they do this, that end of the molecule "looks" slightly negative and the other end "looks" slightly positive. That makes it seem like a dipole. Dipoles attract one another. The difference here is that the dispersion of the elctrons changes, so different parts of the molecule "look" positive or negative at different times. That means these "induced dipoles" are pretty weak, but they still make molecules stick together.
Induced dipole-Induced dipole interactions are individually weak, but for large molecules, there are a lot of them, so their cumulative strength can be fairly strong. If I have some 2-pound test fishing line, it's pretty weak, BUT if I use enough pieces of it, I can lift up a car!


States of Matter (Molecular View):

Why are solids solid? Why are gases gases? States of matter (macroscopic) are the result of the delicate balance between the intermolecular forces present between molecules and the kinetic energy of those molecules.
High kinetic energy, Weak IMFs - If a molecule has high kinetic energy (fast moving) and there are weak interactions between the molecules, then the moelcules will fly by each other so fast that the weak IMFs won't give them time to stick together. This is likely to be a gas.
Low kinetic energy, Strong IMFs - Strongly attracted molecules have kinetic energy (the molecules vibrate), but those vibrations aren't strong enough to overcome the "stickiness" of the strong IMFs. This is likely to be a solid.
Temperature is a way for use to measure (kinetic) energy: as the temperature decreases, the molecules move more slowly. Since IMFs aren't really affected by temperature, as a sample cools (decreases kinetic energy), eventually the balance between kinetic energy and IMFs shifts toward IMFs - gases condense, liquids freeze.



Condensed Phases:

Solids
Crystalline solids have atoms, ions, and/or molecules arranged in a regular pattern. Amorphous solids do not have a regular arrangement.
The kinetic energy in solids is (mostly) due to vibrational motion of the atoms, ions, and/or molecules; the individual pieces do not tend to move with respect to one another.

Liquids
Intermediate phase. Individual atoms, ions, and/or molecules have enough kinetic energy to move with respect to one another, but the IMFs are strong enough to prevent individual pieces from escaping to the gas phase.

Condensed phases are (relatively) dense and (usually) not compressible - they might pour or deform, but the total volume of a condensed phase is (usually) nearly constant under a wide range of temperature & pressure conditions.


Why are gases special?:

Gases have sufficient kinetic energy that they are a little easier to study than condensed phases. There are fewer subtleties and we can often use simple models to understand the behaviour of gas samples.
Many behaviours of gases can be explained using the Kinetic Molecular Theory of Gases.
Kinetic Molecular Theory of Gases:
A gas sample that obeys KMT is said to behave "ideally" or is referred to as an "ideal gas". A gas sample is ideal under a specific range of temperature and pressure conditions. "Ideal gas" does not refer to a specific chemical entity that happens to be a gas... For example, helium behaves ideally under a wide range of temperature and pressure conditions, but helium is not always and ideal gas.

Gas particles are independent and non-interacting.
The properties of gases depend upon the number of particles, not their identity, so mixtures of gases behave just like pure samples.

Pressure: One of the most important properties of gases is pressure. We don't often talk about pressure in condensed phases because it's not really a property worth measuring in condensed phases. In gases, pressure is highly variable and is one of the easier properties to measure. Pressure is caused by gas particles colliding with the container. More collisions = higher pressure. Harder collisions = higher pressure.

Volume: Unlike condensed phases, gases are not dense and are mostly empty space. Because of this, the volume of a gas sample is highly variable when temperature & pressure conditions change.

Gas Laws:

Scientific Laws are statements of observations that are reliably true and consistent. Because temperature, pressure, and volume are (relatively) easy things to measure in a gas sample, there are a lot of reliably observations that can be made about an ideally-behaving gas. Most of these "gas laws" relate the volume of a sample to other properties.
"Simple" Gas Laws - These gas laws involve only 2 properties. All other properties of the sample are held constant.
Avogadro's Law - The volume of a gas sample is proportional to the number of moles of gas particles present in the sample. More moles = more volume.
Boyle's Law - The volume of a gas sample is inversely proportional to the pressure of the sample. Higher pressure = lower volume.
Charles' Law - The volume of a gas is proportional to the (absolute) temperature of the sample. Higher temperature = higher volume.

Absolute temperature - Celsius and Fahrenheit are both relative temperature scales - they assign "zero" to some arbitrary value and define other temperatures relative to that "zero". If we want to use temperature as a direct measurement of energy, we need to use an absolute temperature scale. In an absolute temperature scale, "zero" is defined as the temperature at which there is no energy. When using an absolute temeprature scale, chenges in temperature are directly related to changes in energy; if the absolute temperature doubles, that means the energy of the system has doubled. This also means that there are no negative absolute temperatures. If you haven't already checked out these Temperature Conversion slides, take a look at them.

The Ideal Gas Law - All of the other gas laws can be found in the Ideal Gas Law. If you want to save brain space, memorize this one!

R
The constant "R" comes up in a couple different places, so make sure you're using the right version of it for the problem you're trying to solve. For typical gas law problems, you are probably going to want to use 0.08206 L.atm/mol.K.


Ideal vs. Non-ideal Gases:

An ideal gas is one that obeys Kinetic Molecular Theory. When do gases deviate from ideal behaviour? Let's look at where some parts of KMT begin to break down...
Gases are mostly empty space - This is why gas samples are compressible. When gases are compressed, the space between particles decreases. Since that space is (relatively speaking) so large, it can shrink quite a bit before it becomes "not large", but if a sample of gas is subjected to high enough pressure, the particles are forced close together. Ultimately, this causes the gas to condense to a liquid (or deposit directly to a solid).
Gas particles move randomly - This one usually doesn't cause a lot of problems, but if a gas is forced to flow (like through a pipe or tube) then the particles are not moving quite as randomly, so there can be some non-ideal behaviour.
Attraction & repulsion are negligible - This one is all about the relative size of the kinetic energy and the IMFs of the specific gas particles. The strength of the IMFs is (relatively) constant for a given type of gas particle, but if the kinetic energy decreases enough, the strength of IMFs can end up being of a similar magnitude to the kinetic energy. At that point, the gas particles can "stick together" and start to condense to a liquid (or deposit directly to a solid). Kinetic energy is decreased by lowering the temperature.
So gases behave most ideally when the temperature is (relatively) high and the pressure is (relatively) low, and become less ideal as the temperature decreases and the pressure increases.








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