States of Matter – Standard Deviants: The Super-Charged World of Chemistry
- ______act differently than solids or liquids.
- ______explains why gases act the way they do.
- The five assumptions of Kinetic Molecular Theory are:
- Gas molecules are ______from each other in comparison to their own dimensions.
- Gas molecules are in constant ______motion.
- Molecules move in a ______unless they collide with another molecule or with the wall of the container.
- The molecules exert ______, or the container until they collide with each other, or with the walls of the container. The average kinetic energy of the molecules of a gas is ______to the temperature, this one sounds complex but think about it this way. The movement of the molecules is the kinetic energy.
- Every time a molecule collides it exerts a ______on it.
- ______says that if you decrease the volume of a container of gases, in other words, if you make the container smaller, the number of molecules per unit of volume increases. So, if you put the same amount of gas in half the volume, the pressure of the gas doubles as a result of increasing collisions with the wall.
- Gases can also ______and ______.
- ______is the process by which one gas mixes with another.
- ______is the process by which a gas escapes from a container through a very small hole.
- The rate of diffusion or effusion of a gas is ______to the square root of its molar mass.
- The ______a gas is the slower it diffuses or effuses. This phenomenon is called Graham's law.
- An equation to describe how gases ought to behave is called the ______Equation.It is useful because it shows us the pressure, volume and temperature as well as moles of any hypothetical gas.
- The ideal gas equation is ______.
- ______represents the pressure of the gas.
- ______equals the volume.
- ______equals the number of moles.
- ______is the temperature.
- R is the ______. 8.314 J/mol*K or 0.0821 atm *L/mol
- You choose the form of the gas constant you use depending on whether you need to work with joules (a unit that measures energy) or atmospheres (a unit that measures pressure).
- The ______pressure of a gas in a mixture is the pressure exerted by one of the gases in the mixture.
- Dalton's Law of Partial Pressure says that in a mixture of gases, the total pressure exerted is the sum of the pressures that each gas would exert if it were present alone ______.
- To convert temperature from degrees Celsius to Kelvin, just add ______to the degrees Celsius.
- If your substance consists of a single type of atom, youcan look up the mass of that atom in AMUs in ______.
- The mass in grams per mole of a substance that consists of molecules is called ______.
- When you work with the ideal gas laws, there will be deviations from ideal behavior. Two situations where gases do not obey the ideal gas law are :
- The molecules are crammed into a very small space and the pressure is ______,
- There are ______forces acting between the molecules.
- Intermolecular forces are the forces of ______between molecules.
- Intermolecular forces in the ______state are the weakest.
- Intermolecular forces in the ______state are a bit stronger.
- Intermolecular forces in the ______state are the strongest.
- ______assume the shape of whatever container you put them in.
- When a hydrogen atom bonds to a strong ______atom, it forms a polar bond.
- A ______is the transformation of a substance from one state to another.
- The ______shows us what phase a substance is in at particular temperatures and pressures.
- When two opposing processes are occurring at the same time, for example, if a liquid is evaporating and condensing simultaneously, at the same rate, it is called ______.
- The ______on the diagram represent the equilibrium point between states.
- The intersection of the three lines on the diagram is called the ______.
- This point represents the ______and ______at which all three phases are at equilibrium.
- A phase diagram gives us a ______of the phases of a substance.
Movie Transcript
00:05> [MUSIC]
00:16Look, up in the sky.
00:19What the bird?
00:21> No! > The Standard Deviants present
00:23The Super-Charged World of Chemistry.
00:27> Starring Gelila Asres, Gabrielle Smith,
00:31Jeremy Klavens, Deena Rubinson, Malcolm Smith.
00:37Trevor Dean.
00:40Chuck Hinshaw.
00:43Christiana Celeste.
00:45[SOUND] Tessa Munro. [SOUND] And Chas Mastin.
00:48[SOUND]
00:52[MUSIC]
00:53> [SOUND] So,
00:59how are you
01:05guys?
01:10Who's playing?
01:12> Section A, kinetic molecular theory.
01:15> So, gases act differently than solids or liquids.
01:17Kinetic molecular theory explains why gases act the way they do.
01:23> [SOUND] All these theories.
01:26Can't chemists ever just say anything for sure?
01:28[SOUND] > You could say one thing for certain.
01:31You are getting on our nerves.
01:34> [SOUND] The kinetic molecular theory is based on these five assumptions.
01:39> Number one.
01:41A gas is composed of molecules that are far apart from each other
01:44in comparison with their own dimensions.
01:46So, most of the volume a gas occupies is really empty space.
01:51> I am just a gas in space.
01:56> Two, gas molecules are in constant random motion.
02:00Each molecule continues to move in a straight line
02:02unless it collides with another molecule, or with the wall of the container.
02:06> 3. The molecules exert
02:09no force on each other, or the container until they collide with each other, or
02:13with the walls of the container.
02:16So, if a gas molecule is just cruising around in space,
02:19minding its own business and not running into anything,
02:23it's not exerting any force on its fellow molecules or the container.
02:27> Hey, I'm just cruising, I'm not bothering anybody.
02:30> Oh yeah, until you ran into me, you idiot.
02:33> The average kinetic energy of the molecules of a gas is proportional
02:38to the temperature, this one sounds complex but think about it this way.
02:43When you heat up a pot of water the molecules start moving really fast
02:47then they escape from the pot or boil off.
02:50The movement of the molecules is the kinetic energy.
02:54As the temperature goes up, the molecule move faster.
02:58The more heat there is, the faster the molecules move.
03:01So, the kinetic energy is proportional to the temperature.
03:05> And number five, every time a molecule collides with the wall,
03:11it exerts a force on it.
03:14> [LAUGH] [NOISE] Uh-oh, we did that?
03:20> [SOUND] And here's an effect of this rule number five.
03:25If you decrease the volume of a container of gases, in other words, if you make
03:29the container smaller, the number of molecules per unit of volume increases.
03:34So, if you put the same amount of gas in half the volume, the pressure of the gas
03:38doubles as a result of increasing collisions with the wall.
03:42This phenomenon is called Boyle's Law.
03:44> [SOUND] Half the size, twice the pressure.
03:49[SOUND] Boyle's Law [SOUND] > Now
03:54let's go over diffusion and effusion 2 important properties of gases.
03:58> [SOUND] Gases can also diffuse and effuse.
04:02Diffusion is the process by which one gas mixes with another.
04:06> Effusion is the process by which a gas escapes from a container through a very
04:11small hole, like if you open a helium-filled balloon and
04:14let the helium escape.
04:16The helium gas effuses into the air and once the molecules escape through
04:20effusion, we can say they feel well, effusive.
04:24> The rate of diffusion or effusion of a gas
04:27is inversely proportional to the square root of its molar mass.
04:31What the heck does this mean?
04:33Well, simply this.
04:35The heavier a gas is the slower it diffuses or effuses.
04:40This phenomenon is called Graham's law.
04:44> Not only am I snack food, but I made a law.
04:47Once again, my law states the heavier a gas is, the slower it effuses or
04:53diffuses [LAUGH] > [NOISE] Section B.
04:58The ideal gas equation.
05:00> Chemists, wacky idealists dreamers that they are,
05:03came up with an equation to describe how gases ought to behave.
05:07It is called Ideal Gas Equation and it is useful because it shows us the pressure,
05:13volume and temperature as well as moles of any hypothetical gas.
05:18We use this equation constantly to describe and
05:21understand the behavior of gases.
05:23Of course, the equation is hypothetical since it real life there
05:26are always factors that throw a monkey wrench into the works and
05:29squeeze the numbers but hey, that's why it's called ideal.
05:34You can remember the equation PC equals nRT
05:37because it looks kind of like pervnert.
05:40> In the ideal gas equation, P represents the pressure of the gas.
05:45V equals the volume, n equals the number of moles we're working with,
05:51and R is the gas constant, and T is the temperature.
05:55R, the gas constant is one of those deals you can just look up in your book.
06:01It's 8.314 Joules per mole times Kelvin, or 0.0821 atmospheres
06:10times liters per moles times Kelvin Either one works in the equation.
06:17You can choose which form of the gas constant you use depending on whether you
06:21need to work with joules, a unit that measures energy or atmospheres,
06:26a unit that measures pressure.
06:29So, the gas constant doesn't change and that's why it's called a constant.
06:33You can use the ideal gas equation to find the partial pressure of a gas.
06:38> The partial pressure of a gas in a mixture
06:40is the pressure exerted by one of the gases in the mixture.
06:44Dalton's Law of Partial Pressure says that in a mixture of gases,
06:49The total pressure exerted is the sum of the pressures that each
06:54gas would exert if it were present alone under the same conditions.
06:59So, the total pressure equals the pressure of the first gas plus
07:03the pressure of the second gas plus the pressure of the third, and so on.
07:09So, you would use the PERVNERT ideal gas law to find the pressure of each gas,
07:15and then use Dalton's partial pressure law to find the total.
07:18[SOUND] > Here's an example.
07:20Suppose we have a one liter flask containing 5 grams of carbon dioxide and
07:252 grams of oxygen molecules, and the temperature is 25 degrees Celsius.
07:30We have to find the partial pressures of each of the gases, and
07:34the total pressure in the flask.
07:35We can use the ideal gas equation PV = nRT, but first we have to rearrange it so
07:40it's more useful to us.
07:42We'll set up the equation so it finds the pressure.
07:45For P, which is what we're looking for.
07:47Use some algebra to rearrange the problem.
07:50> Algebra?
07:53> It's not that bad, just divide the whole deal by the volume.
07:57So the equation now looks like this.
07:59Pressure = the quantity of the number of moles times the gas
08:04constant times the temperature divided by the volume.
08:07[NOISE] > Pressure.
08:11Don't tell me about pressure.
08:12I feel pressure from above.
08:14Pressure from below.
08:15[NOISE] > Lets start by converting our
08:17temperature from degrees Celsius to Kelvins.
08:20That's really easy.
08:22Just add 273.15 to the degrees Celsius.
08:27And we get 298.15 Kelvins.
08:31> Remember, to go from Celsius to Kelvins just add 273.15.
08:39[SOUND] > Now we have to determine how many moles
08:44of each gas are in the flask.
08:46> This is the equation we'll use for the conversion, n,
08:49the number of moles, equals 1 mole of your substance over the weight of 1
08:54mole of your substance in grams, times the mass of your sample.
08:59And here's a quick reminder about finding the weight of your substance in
09:02grams per mole.
09:04If your substance consists of a single type of atom,
09:07you can look up the mass of that atom in AMUs in the periodic table.
09:12Then, since one mole of any type of atom
09:15has a mass in grams equal to the atom's atomic mass measured in AMU's,
09:19you can literally just change your units from AMUs to grams per mole.
09:24The mass in grams of compound that's made up of crystals,
09:27like NaCl, is called the formula weight.
09:31The formula weight is just the sum of the masses of the atoms in the compound's
09:35empirical formula.
09:36The mass in grams per mole of a substance that consists of molecules
09:40Is called molecular weight.
09:42To find molecular weight, take the mass of each of the molecule's component atoms and
09:48AMU's, and add them up.
09:50So for instance, to find the mass in grams per mol of the oxygen
09:54molecule in our pressure example, just find the mass of oxygen,
09:59multiply it by two, and slap the unit grams per mole on the end.
10:03So here we go, converting our sample of oxygen from grams to moles so
10:08we can find its partial pressure in the flask.
10:12n, the number of moles of oxygen molecules equals 1 mol of
10:17O2 over the weight of 1 mol of O2 in grams,
10:2132 grams, times our sample, 2 grams of O2.
10:27Multiplying our sample, 2 grams, by 1 mol of O2 over the mass
10:32in grams of a mol of O2, is like multiplying by 1.
10:36So it changes units without changing amounts.
10:40See, if you weighed out 1 mol of oxygen molecules, it would weigh 32 grams.
10:46So in effect, 1 mol of O2 and 32 grams of O2 are the same thing.
10:53That's why it's like multiplying by 1.
10:551 mole of O2 over 32 grams of O2 multiplied by our 2 gram sample
11:02gives us 0.0625 moles of oxygen molecules.
11:07[MUSIC]
11:11Remember, the ideal gas law, PV=nRT is rearranged to solve for pressure.
11:17Pressure equal the quantity of the number of moles times the gas constant
11:22times the temperature divided by the volume.
11:25> Now let's plug the numbers into our equation to find the pressure
11:29of the oxygen in the flask.
11:31The pressure equals the quantity 0.0625 moles times the gas constant.
11:37Now you need to choose one of the two gas constants.
11:40Since we want pressure, and since the total flask is measured in liters,
11:44use the one in atmospheres and liters.
11:46That's 0.0821 liters times atmospheres per moles times Kelvin.
11:52Those two quantities are multiplied by the temperature,
11:56which we determined was 298.15 Kelvins.
12:00And all that is divided by the volume, 1 liter.
12:03Now here comes a little more algebra.
12:05[SOUND] > Relax.
12:07It's not hard.
12:10> The denominator in the gas constant,
12:12mols times Kelvin, shifts down to the denominator of our equation.
12:17So we'll cross out all the terms in the equation and
12:20multiply the terms in the numerator.
12:23That leaves us with 1.53 atmospheres, the pressure of the oxygen in the flask.
12:28> On to carbon dioxide.
12:30We have to convert from grams to moles again.
12:33The number of moles of carbon dioxide equals one mole of carbon dioxide over
12:38the molecular weight of one mole of carbon dioxide, 44 grams,
12:42times the weight of our sample, 5 grams.
12:45That equals 0.114 moles.
12:48> Now we're ready to find the pressure of the carbon dioxide in the flask.
12:52The pressure, P, equals the quantity of 0.114
12:57moles times the gas constant, 0.0821 liters,
13:02times atmospheres per moles, times Kelvin, times the temperature,
13:08198.15, all divided by one liter.
13:14Again, we'll move the denominator of the gas constant to the denominator of
13:18the equation and cancel out the units.
13:20That gives us 2.79 atmospheres,
13:23the pressure of the carbon dioxide in the flask.
13:28Finally, to find out the total pressure in the flask,
13:31just add up the individual partial pressures.
13:341.53 atm + 2.79
13:39atm = 4.32 atm.
13:43So the total pressure is 4.32 atmospheres.
13:48> And now a bit more about moles.
13:51There is a simplified version of the equation to convert grams to moles
13:55that your esteemed chemistry professor will probably expect you to know.
14:00It's really quite simple.
14:02> Here's the equation you've been using to convert from grams to moles.
14:06Like we said, it uses the unit factor method to change units from grams to
14:10moles, by multiplying the weight of your sample by one mole of the sample,
14:15over the molecular or formula weight of the sample in grams per mole.
14:20> Grams to moles, grams to moles.
14:24> Now here's the simplified version.
14:26You can just ignore the 1 mole of substance in the numerator
14:29because it doesn't do anything for the values in your equation.
14:33Just take a shortcut.
14:34Put the mass of your sample right over the molecular or
14:37formula weight in grams per mole of your substance, and calculate it that way.
14:42> So, for instance, this is how we converted carbon
14:45dioxide from grams to moles in the pressure calculation we just did.
14:49Let's do it again using the simplified version of the equation.
14:52> We'll put the mass of the sample of carbon dioxide, 5.00 grams,
14:58over the molecular weight of carbon dioxide, 44.0g/mol.
15:03So cancel out the grams and divide 5.00g by 44.0g/mol and
15:10you get the same answer you got before, 0.114mol.
15:16[SOUND] > Enough about the moles already.
15:20Get back to the gases.
15:24> When you work with the ideal gas laws,
15:26keep in mind that there will be deviations from ideal behavior.
15:29> In real life, gases don't obey the law.
15:33Two excuses that gases have for not obeying the ideal gas law are one,
15:37when the molecules are crammed into a very small space and the pressure is high.
15:42The molecules take up a significant fraction of the empty space, and
15:45this can skew your results.
15:48And two, there are intermolecular forces acting between the molecules.
15:52Now, what the heck are intermolecular forces?
15:55Stay tuned.
15:56[MUSIC]
15:59So the ideal gas equation describes the pressure, volume, and
16:04temperature of any hypothetical gas.
16:06You can rearrange the ideal gas equation to find the partial pressures of
16:11gasses in a container.
16:13Then, if you add up the partial pressures,
16:16you get the total pressure of the gases in the container.
16:20It's important to remember that the ideal gas equation
16:23describes hypothetical gases that behave exactly like they're supposed to.
16:27In real life,
16:28gases just don't always behave like the ideal gas equation says they should.
16:33Part III, States of Matter.
16:36Section A, Intermolecular Forces.
16:40Intermolecular forces are the forces of attraction between molecules.
16:44They are not as strong as ionic and covalent bonds, but
16:47they definitely have an effect on states of matter.
16:50The intermolecular forces operating on gas molecules are not that strong,
16:55and they allow gas molecules to move away from each other.
16:59The intermolecular forces between molecules in a liquid are a bit stronger,
17:04so liquids hold together and
17:05assume the shape of whatever container you put them in.
17:08Solids have strong intermolecular forces at work
17:11between the molecules locking the molecules in place.
17:15> [SOUND] Intermolecular forces are responsible for
17:18the condensation of gas molecules into their liquid state.
17:22Gas molecules normally move pretty fast, but as the temperature decreases,
17:27the speed, or kinetic energies of the molecules, also decrease.
17:32As gas molecules slow down, they begin sticking together.
17:36And when they finally do stick together, they form a liquid.