Chpt. 13, Solids, Liquids, and Gases Flashcards Preview

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Flashcards in Chpt. 13, Solids, Liquids, and Gases Deck (32):

intermolecular forces

forces that occur between different molecules, rather than within a single molecule


dipole-dipole forces

forces of attraction that occur between polar molecules


hydrogen bonds

extremely strong dipole-dipole force that occur between molecules with H-O, H-F, and H-N bonds; because these molecules are extremely polar due to the high difference in electronegativity between these two molecules, and because the lone pairs on the O, F, or N atoms can get really close to the hydrogen atom on another molecule, these forces are way stronger than regular dipole-dipole forces


London dispersion forces

attractions caused by temporarily induced dipoles in non-polar molecules; Van der Waals forces are almost the exact same thing as these, although there are some small differences; the process occurs as follows:

1. random movement of electrons in a molecule causes it to be temporarily polar

2. this polarity causes electrons to shift in an adjacent molecule (be driving them away from the negative spot), making it polar, too

3. the two molecules, temporarily polar, stick to each other in the manner of dipole-dipole forces

4. the electrons eventually go back to their original state, doing away with the temporary polarity


similarity between liquids and solids

states of matter in which the attraction between molecules, ions, or atoms keep them from moving away from each other


ionic solid characteristics

hard, brittle, high melting and boiling points, and conductive of electricity when melted or dissolved, generally less flammable than covalent compounds


molecular solids

crystalline solids in which molecules are stuck to each other via intermolecular forces; the melting and boiling points of these compounds will depend largely on how strong these forces are, as well as how big the molecules are, though mp; they are also generally soft, electrical insulators, and have lower energy; examples include sugar and water ice


covalent network solid characteristics

crystalline solids held together by multiple of covalent bonds; though there are exceptions, these solids have very high melting and boiling points, though lower than ionic compounds, more likely to burn than ionic compounds (C and H have more occurrence in them due to their identical electronegativities); examples include quartz, diamond, and silicon


metallic solids

malleable, ductile, conducive of electricity, highly conductive of thermal energy, luster, insoluble in most solvents (electrons cannot easily go into solution)


amorphous solids

unlike crystalline solids, there’s no long-range order in an amorphous solid; amorphous solids contain compounds as varied as glasses, rubbers, and plastics; they are the only type of solids that does not exhibit a crystal structure


atomic solid

solids that occur when noble gases are cooled to very low temperatures and lock themselves in place using very weak London dispersion forces; these solids almost never occur naturally because they only occur when temperatures are extremely low


other types of substances

alloys (interstitial or substitutional)


flammability criteria

cannot be flammable unless it contains carbon and oxygen



the state of matter in which particles are attracted to each other enough that they don’t fly apart (like a gas) or stick in place like a solid; liquids generally have lower densities than solids, but have much higher densities than gases; as a result, if a solid is melted, it will grow in volume


properties of liquids

1. cannot be compressed-- already very close to one another

2. they flow-- they have enough attraction to keep them together, but not enough to prevent them from flowing apart due to gravitational and other forces

3. they have viscosity

4. they have surface tension-- stronger the intermolecular forces of a liquid, the higher the surface tension

5. they experience capillary action


capillary action

the process by which some liquids are pulled up a thin tube, this is caused by cohesion and adhesion; the liquid is pulled up until the gravitational downward force equals the upwards force of the cohesion and adhesion; the thinner the tube, the higher the liquid will be pulled



another term for "thickness"; what it really measures is the strength of the force between the particles of a liquid


phase changes

melting (solid to liquid)
freezing (liquid to solid)
vaporization (liquid to gas)
condensation (gas to liquid)
sublimation (solid to gas)
deposition (gas to solid)



This is when you turn a solid into a liquid. This is done by adding enough energy to disrupt the forces between the particles in the solid, but not so much that the particles in the solid fly apart into a gas.



This is when a liquid turns into a solid. As you suck the energy out of a liquid, there’s less energy available for the molecules to overcome the forces between the particles. As a result, the particles lock into place as a solid. This is what happens when you turn water into ice by cooling it down.



This is when you turn a liquid into a gas by adding energy to it - this added energy breaks apart the forces between the particles, allowing them to fly wild and free as gases.



When you cool a gas enough, the particles no longer have enough energy to overcome the forces between them. At this point, the gas condenses into a liquid.



Sublimation is when a solid turns directly into a gas. This occurs when the vapor pressure of the liquid phase of the material is so high that the material goes directly from a solid to a gas without first turning into a liquid. You know how ice cream gets all disgusting and rubbery after it’s been in the freezer for a long time. That’s because the water sublimed out of it and only the rubbery other stuff is left over.



When a gas turns into a solid. This happens when the gas loses energy and goes from the gas to the solid state. When ice cream grows ice crystals when it gets rubbery in the freezer, that’s because the water was deposited as ice crystals.


evaporation and boiling

One example of vaporization is evaporation, which occurs when only a few molecules of the liquid have enough energy to become a gas – the pressure of the gas caused by the evaporation of these particles is called its vapor pressure. As the temperature of the liquid increases, evaporation increases and the vapor pressure of the liquid increases, too. When the vapor pressure of the liquid increases to where it’s the same as the atmospheric pressure, the liquid boils. An example of evaporation is when your goldfish dies because his bowl dried out.


phase changes diagrams: lines, normal melting point, and normal boiling point

1. Each line that marks the border between two phase changes denotes the conditions under which both phases of matter can stably exist.

2. The temperature at which the compound melts at a pressure of one atmosphere. In this diagram, the normal melting point is 00 C.

3. The temperature at which the compound boils at a pressure of one atmosphere. In this diagram, the normal boiling point is 1000 C.


phase change diagram triple point and critical point

1. The conditions of temperature and pressure at which all three phases of matter can stably exist. For water, the triple point is 0.06 atm and 0.010 C, which is why you’ve never seen all three phases of water in equilibrium.

2. The conditions of temperature and pressure past which it’s impossible to distinguish between the liquid and gas phase of the material. This occurs because the material has too much energy to stick together (which is true of gases) but is crammed so tightly together that intermolecular forces between the particles are strong. Under these conditions, the material is said to be a supercritical fluid.


kinetic molecular theory of gases

The kinetic molecular theory of gases is a mathematical model that’s intended to quantitatively describe the behavior of gases. Unfortunately, the math that goes into such a model is really hard to do, and requires annoyingly huge quantities of computing power. Because we usually don’t feel like buying such gigantic computers (and because the model was invented in the mid-19th century, before computers were even born), the kinetic molecular theory uses some assumptions about the particles in gases to make the whole thing a little easier to deal with.


postulates of the kinetic molecular theory

1. gas molecules are infinitely small
2. gas molecules are constantly moving, and never move preferentially in one direction
3. gas molecules undergo elastic collisions; when two hit each other, none of the energy is lost
4. the kinetic energy of a gas molecule is related to its temperature; the higher the temperature, the higher the kinetic energy
5. gases don’t experience intermolecular forces


Why are gas molecules, "infinitely small?"

Because the particles in a gas are all really far away from each other, the volume of the particles in a gas is really small compared to its overall volume. Since the gas molecules take up a fantastically small percentage of its volume, we can just make the assumption that they’re infinitely small to make the math easier.



How gas travels can be described by either diffusion or effusion.

Diffusion is when gas travels through other gas. Effusion is when gas travels through a vacuum containing no other gas molecules.


rate of diffusion

rate of gas A / rate of gas B = ∫molar mass of gas B / molar mass of gas A