Chapter 14: Liquids and Solids (or States of Matter)

Goals:

  • Learn about Van der Waal’s forces (London-dispersion forces, dipole-dipole attraction, and hydrogen bonding), and understand the effects these intermolecular forces have on solids and liquids.
  • Explore water – including how hydrogen bonding accounts for many of its properties.
  • Understand what is occurring during phase changes, and how to use the heat of vaporization, heat of fusion, and heat capacities of a substance to perform calculations about the energy changes that occur with phase changes.
  • Understand the relationship between vaporization (both evaporation and boiling) and vapor pressure.
  • Relate the boiling point of substances (including water) to their vapor pressure.
  • Learn about the various types of solids, and how these types result from different types of bonding and intermolecular attractions.

As you’re no doubt aware, the three states of matter we work with in chemistry are solids, liquids, and gases. In the next few classes we’ll explore how these states work on a microscopic level.

Table: General properties of the three states of matter

property / solid / liquid / gas
density / high / high / Very low
compressibility / low / low / Very high
volume / constant / constant / fills container
shape / constant / Takes shape of container / Takes shape of container
structure / Neat, orderly / Somewhat random / disordered / completely random / disordered

We’ll discuss why each state has each of the properties above.

Gases (a quick review)

The Kinetic Molecular Theory of Gases:

  • Definition: The kinetic molecular theory (KMT) of gases says that the properties of gases are determined by the interactions of the gas atoms/molecules with each other.
  • Basic postulates of the kinetic molecular theory:
  • Gas particles are infinitely small:
  • In a gas, the gas molecules are separated from one another by a lot of empty space. As a result, the gas particles can be said to have negligible volume compared to the overall volume of the gas. This has the advantage of making the math much easier.
  • Gas particles are in constant, random motion:
  • In a gas, the particles constantly move very quickly all over the place, changing direction only when they hit something and bounce off. Energy is transferred between particles when they collide, but this energy is conserved (not lost).
  • The collisions of gas particles with the sides of the container they’re in is called pressure.
  • Gas particles don’t experience intermolecular forces:
  • Because the gas particles are so small, are so far apart, and are moving so quickly, the molecules don’t interact with each other much. To make our lives (and the math easier), we just ignore the very little interaction that does take place.
  • The kinetic energies of gas particles are proportional to their temperatures (in Kelvin).
  • This should make sense: All this says is that the hotter the particles in a gas are (the faster the average speed of the particles), the more energy the gas particles have.
  • Why Kelvin? Because if you used degrees Celsius, you’d get negative energy whenever the temperature dropped below the freezing point of water (0o C). As a result, we have to use a temperature scale where the zero point energy of a molecule corresponds to zero on the temperature scale.
  • This relationship is mathematically described by: (this formula should look familiar)

Properties of Gases:

  • Gases have low density: This makes sense: If there’s lots of space between gas molecules, then you’d expect low density.
  • Gases can be compressed and expanded:
  • You can compress a gas because all you’re doing is just squishing the particles together.
  • Gases expand because the particles move all over the place unless you stop them from doing so.
  • Gases diffuse: Diffusion is when two things put in the same container mix together.
  • For example, if you open a jar of pickles in a room, you can soon smell pickles all over the place – this is because the pickle smell diffuses throughout the room.
  • Gases effuse: Effusion is when a gas escapes through a hole in a container into a vacuum.
  • Because of ½ mv2, light molecules (those with a low molar mass) move more quickly than heavy ones (if they are at the same temperature). As a result, light molecules diffuse and effuse more quickly than heavy ones.
  • Graham’s Law of effusion:
  • Sample problem: If the scent molecules in pickle juice have a molar mass of 450 g/mol and the scent molecules in ammonia have a molar mass of 17 g/mol, how much faster will ammonia molecules move than pickle molecules?

5.15 (21.21/4.12)

Liquids: The Joy of Intermolecular Forces

Whereas solids are materials in which the particles are very tightly stuck to one another, liquids are materials in which the particles have a little more interaction.

As you already know from experience, liquids are materials that can easily change shape but also tend to have fairly high densities (nearly the same as solids). Why do the molecules hang together at all? The answer: Intermolecular forces.

Intermolecular forces -Interactions that hold the particles in a liquid together. There are three types of intermolecular forces (often called Van der Waals’ forces)that we need to consider. (All three are varying degrees of attraction between partial positive charges and partial negative):

  • Dipole-dipole forces: Interactions in which polar molecules stick to each other like little magnets when the partial positive side of one is attracted to the partial negative side of another.
  • The more polar the molecule, the stronger the attraction!
  • Hydrogen bonds: A very strong dipole-dipole force that occurs when the lone pair electrons on O, F, or N interacts strongly with a hydrogen atom bonded to O, F, or N.
  • Essentially, these bonds are so polar that the lone pair electrons (negative) on one molecule want to stick to the very positive hydrogen atoms on another molecule.
  • The more hydrogen bonding that a molecule can do, the stronger this force is.
  • Water has a MP of 00 C, while methanol has a MP of -980 C.
  • London Dispersion Forces: When nonpolar molecules are attracted to one another via temporarily induced dipoles.
  • Essentially, nonpolar molecules stick together magnetically – like polar molecules. How can this work?

  • The bigger the molecules, the stronger the force (because there are more electrons to become unbalanced and interact with each other).
  • Why are intermolecular forces important?
  • The stronger the intermolecular force, the higher the melting and boiling points.
  • Because melting and boiling both involve the movement of particles from their neighbors, anything that causes neighboring particles to stick together will raise them.
  • Intermolecular forces make liquids almost as dense as solids:
  • Because the intermolecular forces in liquids keep the molecules stuck to each other, liquids are nearly as dense (and in a few cases even denser) than solids.
  • Intermolecular forces allow liquids to flow – a property called “fluidity”.
  • Because the molecules in a liquid are attracted to each other but not permanently stuck in place, the molecules can move from one place to another – flowing!
  • Intermolecular forces cause differences in the viscosities of liquids:
  • Definition: Viscosity is the resistance to a liquid flowing. High viscosity = slow flowing.
  • The higher the intermolecular force, the more viscous the liquid. This is because molecules that are held tightly together want to move apart less than molecules that are loosely held.
  • Viscosity also goes down with increasing temperature (i.e. things flow more easily at high temperatures) – this is because the energies of the particles in the liquid are beginning to get to the point where they can overcome some of the attractive forces.

Intermolecular forces cause surface tension in a liquid.

  • Definition: Surface tension is the energy needed to increase the surface area of a liquid – the higher the surface tension, the harder it is for something to push through the surface of a liquid.
  • Stronger intermolecular forces cause the surface molecules to hold together more tightly, making the surface tension higher.
  • This is why some things that are heavier than water (i.e. water bugs, leaves, etc) don’t fall through.
  • Intermolecular forces cause capillary action.
  • Capillary action: The tendency of some liquids to rise when placed in a small tube – this explains why putting one edge of a paper towel will eventually cause the whole towel to get wet, as well as how water gets to the top of a tall tree.
  • This happens because water molecules want to grab the surface of the walls of the tube with their intermolecular forces more than they want to grab each other. This causes them to move up the sides of the tube (away from each other).
  • This causes the meniscus.
  • Eventually, the pressure of the water height overcomes the attraction of the water for the sides of the tube and the water stops rising.
  • Ranking of intermolecular (and intramolecular)forces (strongest to weakest):
  • Ionic interactions and covalent bonds (they are not intermolecular forces, they are intramolecular forces) – it takes a huge amount of energy to break these.
  • Example: The BP of NaCl is 14130 C.
  • Hydrogen bonding
  • Example: The BP of water is 1000 C.
  • Dipole-dipole forces
  • Example: The BP of H2S is -59.60 C.
  • London dispersion forces
  • Example: The BP of methane is -161.50 C.

Questions that may be asked:

  • What intermolecular force is molecule X likely to experience?
  • Draw the Lewis structure of the molecule.
  • If the molecule is nonpolar, it’s Van der Waals forces.
  • If the molecule has H bonded to O, N, or F, it’s hydrogen bonding.
  • If the molecule is polar but H isn’t bonded to O, N, F, it’s dipole-dipole forces.
  • [do some examples: CH4, CH4O, HCN]
  • Rank the following by increasing melting/boiling point.
  • To solve, determine the type of intermolecular force that each molecule is undergoing.
  • You may assume that molecules with stronger intermolecular forces will have higher MP and BP than those with weaker ones.
  • Example: Rank CH2O, CO2, and CH2O2

(ranking intermolecular forces worksheet)

Phase Changes:

Solid  Liquid transformations:

  • Melting is when a solid becomes a liquid. The reverse of this process is called freezing.
  • Melting and freezing are the reverse of one another and happen at the same temperature.
  • Why things melt:
  • Solids melt when the amount of energy that’s available (because we’ve heated them) is greater than the amount of energy that’s holding them together.
  • Covalent compounds melt at low temperatures because the amount of energy that holds the particles together through intermolecular forces is very small.
  • Ionic compounds melt at high temperatures because the lattice energy that holds the ions together is very high.
  • Why things freeze:
  • Liquids freeze when enough energy has been taken away from the liquid that the particles are no longer able to stay separate – the intermolecular forces (or lattice energy, in the case of ionic compounds) now force them to combine into a solid.

Liquid  Gas transformations:Vaporization and Condensation

  • In a liquid, the particles move into the gas phase when they get enough energy to break free of the intermolecular forces that hold them together.

  • Vaporization – liquid to gas (can be evaporation or boiling)
  • Evaporation is the process in which only a very few of the molecules in a liquid have enough energy to break free. It occurs below the boiling point, and only at the surface.
  • The pressure of the molecules that have become a gas is called the vapor pressure of the liquid.
  • The higher the temperature of the liquid, the higher the vapor pressure (because more of the molecules have gotten enough energy to become a gas)
  • Boiling: When the vapor pressure of the liquid becomes equal to the atmospheric pressure, the molecules in a liquid have gotten enough energy to break free of the intermolecular forces that hold them together.
  • Boiling point: The temperature at which this happens (varies with pressure).
  • Condensation: When enough energy has been removed from a gas that intermolecular forces again hold them together as a liquid.
  • Condensation and vaporization are the reverse process of one another and happen at the same temperature.

Solid  Gas transformations: sublimation and deposition

  • Sublimation is when things go directly from the solid phase to the gas phase. This happens with dry ice (when the white gas comes off of the block).
  • This is why ice cubes in the freezer get smaller over time.
  • This is what causes “freezer burn.”
  • This is how things are freeze-dried (though at pressures below 1 atm)
  • Deposition is when things go from the gas phase directly to the solid phase.
  • This is why frost sometimes builds up on the sides of a freezer.
  • Snow and frost are deposition.

Phase diagrams: How we figure out what’s going on

  • Phase diagrams: Diagram that shows you what phase changes occur at different temperatures and pressure.
  • Obviously, phase changes take place when the temperature changes – that’s how we normally boil water and melt ice.
  • However, pressure is also important – remember how things boil if the vapor pressure = the atmospheric pressure? Well, if we decrease the pressure inside a container, things boil at lower temperatures.
  • Pressure changes have a similar effect on other phase changes.

Important features of phase diagrams:

  • Lines: Along the lines that separate the phases, both phases are equal to stably coexist. That’s why you can put a glass of ice water in the refrigerator and find both the ice and liquid water there after a few days.
  • Normal freezing point: The temperature at which a substance freezes/melts at a pressure of 1 atm.
  • Normal boiling point: The temperature at which a substance boils/condenses at a pressure of 1 atm.
  • Triple point: The conditions of pressure and temperature at which all three phases of matter can stably coexist. For water this is 0.006 atm and 0.010 C, which makes it impossible to observe without special equipment.
  • Critical point: The temperature above which water cannot exist as a liquid. Above this temperature water exists as something between a liquid and a gas.

Some terms and ideas you need for studying thermodynamics:

As with everything, thermodynamics has specialized terms to describe the things that go on in the real world. Before moving on to seeing how energy behaves, we need to first understand the terms that are used to describe it.

Energy: The ability to do work or to produce heat. There are two types of energy:

  • Kinetic energy: The energy something has when it moves. (i.e. moving objects, moving particles, vibrating molecules, etc.)
  • Temperature is a measure of the particles in an object. We know this from the KMT, which says that the amount of energy is proportional to the temperature (in K). The more the particles in an object move around, the higher the temperature.
  • Potential energy: Stored energy that’s waiting for its chance to get moving. (i.e. objects that are waiting to fall off of a shelf, energy stored in chemical bonds, etc.).
  • Chemical potential energy: The energy that’s stored in chemical bonds.
  • Heat (q): The movement of energy from one thing to another through the motion of molecules (thermal energy).
  • Heat spontaneously moves from hot things to cold. This is why a hot pan can burn you and you can’t burn a hot pan – the energy goes only from the pan to you because it’s hotter.
  • Heat and temperature are NOT the same thing: Heat is the transfer of energy, temperature is a measure of the kinetic energy of the object once the energy has finished transferring.

Quantifying energy:

  • The traditional unit of energy is the calorie (cal), which is the amount of energy you need to add to 1 gram of water to heat it by 10 C.
  • Food is measured in units of 1000 calories called kilocalories (kcal), which is more commonly known as the Calorie (Cal).
  • The metric unit of energy is the joule (J). There are 4.184 J/cal.
  • Because a joule isn’t very much energy, we usually measure energy in units of 1000 joules called kilojoules (kJ).

(In class practice worksheet about energy, followed by more notes)

Enthalpy: ∆H

Now that we’ve learned what heat is, it’s handy to think about how much heat a system can potentially give off to other systems. This term is called “enthalpy.”

Enthalpy (H): The amount of heat that a system can potentially give to other systems.

  • Unfortunately, it’s impossible to know what the overall enthalpy of a system is – after all, how much heat one system can give to another depends on a large number of factors, including the nature of the system and the nature of whatever system it wants to give energy to.
  • Instead of talking about how much enthalpy something has, we instead talk about how much the enthalpy of a system changes when heat is taken away from it or added to it.
  • This term is given the symbol ∆H (“delta H”), where ∆ represents the change in enthalpy that occurs during a process.
  • Enthalpy changes having to do with phase changes:
  • Heat of vaporization (∆Hvap) is the amount of heat required to convert 1 gram(or 1 mole) of a liquid at its boiling point into vapor without an increase in temperature.Watch the units! (may be called the Molar heat ofvaporization if it is the heat required for 1 mole.)
  • Heat of fusion (∆Hfus)is the amount of heat required to convert 1 gram (or 1 mole) of a solid at its melting point into a liquid without an increase in temperature. Watch the units! (may be called the Molar heat offusion if it is the heat required for 1 mole.)
  • Heat capacity is the heat required to raise 1 gram(or 1 mole) of substance by one degree Celsius (or one Kelvin). (Molar heat capacity is per one mole; while Specific heat is per 1 gram.) Heat capacity is often shown as C; and often with a subscript p indicating that the pressure is held constant.
  • The heat capacity of a substance varies depending on which phase it is in. The specific heat of liquid water is 1 calorie/gram °C = 4.184 joule/gram °C Watch the units!

Enthalpy changes caused by heating/cooling: