The structure of solids: There are two main types of structure that solids have:

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1 States of Matter Introduction: As you re no doubt aware, the three states of matter we work with in chemistry are solids, liquids, and gasestable: General properties of the three states of matter property solid liquid gas density highest high low compressibility low low high volume constant constant fills container shape constant fits container fills container structure ordered disordered disordered Solids The single characteristic that defines the properties of solids is this: In a solid, the particles don t move around very much. This is because, for various reasons, the particles in a solid are locked into place. As a result, solids are generally quite dense and hard. The structure of solids: There are two main types of structure that solids have: Crystalline solids: Solids win which the molecules/atoms are stacked together in orderly and repeating 3-D structures. o Each of the repeating units that a solid is made up of is called a unit cell these unit cells, when repeated over and over again, give the structure of the entire crystal. There are many different types of unit cells (described on p. 401). The three types I want you to commit to memory are the cubic structures: Simple cubic Body-centered cubic (BCC) Face-centered cubic (FCC)

2 o The five types of crystalline solid: Atomic solids: Noble gas crystals in which the atoms are held together by weak London dispersion forces. Examples: Noble gases. Properties: Very low MP/BP, soft. Molecular solids: Solids in which covalent molecules are locked into crystals by various intermolecular forces (London dispersion forces, dipole-dipole forces, or hydrogen bonds). Examples: water, ammonia, carbon dioxide, etc. Properties: Low/medium MP/BP, soft. Covalent network solids: Solids in which atoms bond together covalently to form crystals these are basically great big molecules. Examples: diamond, quartz, etc. Properties: Often indistinguishable from ionic compounds, except that they almost never conduct electricity in any form. Ionic solids: Solids in which ions are held together via electrostatic attractions. Examples: NaCl, Li 2 CO 3, etc. Properties: Check yer notes. Metallic solids: Solids in which delocalized electrons hold the metal atoms together. Examples: Any metals. Properties: Usually high MP/BP, hard, malleable/ductile, shiny.

3 Amorphous solids: Solids that have no regular structure (i.e. they don t form crystals). o This frequently occurs when liquids solidify extremely quickly - because it takes time for the particles in a crystal to orient themselves properly, a very rapid cooling may cause this process not to work. The resulting solids are typically hard and very brittle. Examples: Glasses. o It sometimes also happens (especially in organic compounds) that very long molecules tangle up with each other to form very soft and pliable solids. Examples: Rubber, plastics Lab: Identifying which type of solid is present from easily observed properties. Solids HWK sheet: Talks about the properties of solids as well as how and why they differ from those of liquids.

4 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). The big question: 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 that we need to consider. Dipole-dipole forces: Interactions in which polar molecules stick to each other like little magnets. o 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. o Essentially, these bonds are so polar that the lone pair electrons on one molecule want to stick to the very positive hydrogen atoms on another molecule.

5 o The more hydrogen bonding that a molecule can do, the stronger this force is. Water has a MP of 0 0 C, while methanol has a MP of C. London Dispersion Forces (also called Van der Waals forces): When nonpolar molecules are attracted to one another via temporarily induced dipoles. o Essentially, nonpolar molecules stick together magnetically like polar molecules. How can this work? o The bigger the molecules, the stronger the force (because, after all, there are more electrons to become unbalanced and interact with each other).

6 Neon has a MP of , while F 2 has a MP of C and methane has a MP of C. Why are intermolecular forces important? The stronger the intermolecular force, the higher the melting and boiling points. o 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: o 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. o 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: o Definition: Viscosity is the ease with which a liquid can flow. High viscosity = slow flowing. o 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. o 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.

7 Intermolecular forces cause surface tension in a liquid. o 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. o Stronger intermolecular forces cause the surface molecules to hold together more tightly, making the surface tension higher. o 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. o 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. o 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). o This causes the meniscus. o 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 forces (strongest to weakest): o 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 C. o Hydrogen bonding Example: The BP of water is C. o Dipole-dipole forces Example: The BP of H 2 S is C. o London dispersion forces Example: The BP of methane is C.

8 Gases Introduction: In the early days of chemistry, it was easy for scientists to study what happens in solids and liquids. However, this isn t the case with gases. Because gases are usually invisible, frequently have no smell, and are very hard to work with, it took much longer for scientists to understand how gases work. In fact, it wasn t until the 19 th century that people really began to get an idea of how the particles in a gas interact with one another. This growing understanding eventually became 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: o 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.

9 o 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. These collisions are elastic, which means that there is no loss of energy during the collisions. The collisions of gas particles with the sides of the container they re in is called pressure. o Gas particles don t experience intermolecular forces: Because the gas particles are so small and 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 takes place.

10 o The kinetic energies of gas particles are proportional to their temperatures (in Kelvin). This should make sense: All this says is that the more hot the particles in a gas are, 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 (0 o C). As a result, we have to use a temperature scale where the zero point energy of a molecule corresponds to zero. Review converting degrees Celsius to Kelvin. This relationship is mathematically described by: 1 KE = mv 2 2 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: o You can compress a gas because all you re doing is just squishing the particles together. o Gases expand because the particles move all over the place unless you stop them from doing so.

11 Gases diffuse: Diffusion is when two things put in the same container mix together. o 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. o Because of ½ mv 2, 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. o Graham s Law of effusion: Rate MW A = Rate MW B o 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) A B Phase Changes: Up until now, we ve talked about the three phases of matter. However, from our own personal experience we know that solids can become liquids and that liquids can become gases. How does this work? Solid Liquid transformations:

12 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: o 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: o 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: 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.

13 Evaporation is the process in which only a very few of the molecules in a liquid have enough energy to vaporize. o The pressure of the molecules that have become a gas is called the vapor pressure of the liquid. o 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) This explains why a hot shower steams up the bathroom while a cold shower doesn t. When the vapor pressure of the liquid becomes equal to the atmospheric pressure, this means that the average energy of all the particles has become greater than the amount of energy needed to enter the vapor phase. o Vaporization/boiling: When the molecules in a liquid have gotten enough energy to break free of the intermolecular forces that hold them together. o Boiling point: The temperature at which this happens. Condensation: When enough energy has been removed from a gas that intermolecular forces again hold them together as a liquid.

14 Condensation and vaporization are the reverse process of one another and happen at the same temperature. Solid Gas transformations: 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). o This is why ice cubes in the freezer get smaller over time. o This is how things are freeze-dried. Deposition is when things go from the gas phase directly to the solid phase. o This is why frost sometimes builds up on the sides of plastic bags or on the sides of a freezer. Phase diagrams: How we figure out what s going on As with many things in chemistry, we like to take all of this information and put it into a chart that makes it possible to see what happens at a glance. Phase diagrams: Tables that show you what phase changes occur at different temperatures and pressure.

15 o Obviously, phase changes take place when the temperature changes that s how we normally boil water and melt ice. o 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.

16 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 atm and 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. [Ask them some questions about the phase diagram things along the lines of what happens as we move from point A to point B? and such.]