solutions and colloids

Solutions and Colloids I.  Solutions
    Definitions to know:
    solution:  homogenous mixture; uniform throughout
    solvent:  substance that does the dissolving; ** substance which is in greater quantity.
    solute:  substance that is dissolved;  ** substance which is in smaller quantity.
              Water is the most common solvent!  Why?
               1.  Water is a polar molecule because of the lone pairs on oxygen.
                2.  Thes lone pairs of electrons are attracted to ions on crystal surfaces.
                     This attraction seperates the ions from each other and the crystalline
                     solid dissolves.
    dissociation:  seperation of ions from each other.
    solvation:  occurs when the solvent surrounds th particles of solute.
    electrolyte:  substances that break up (ionize or dissociate) in water to produce ions.
                         They are able to conduct electric current.  They usually consist of ionic compounds.
                         (Acids and bases are electrolytes)
    nonelectrolyte:  substances that do not break apart and do not conduct electricity.
                                   They are usually covalent compounds with the exception of acids.
II.  Solubility
      Main rule of solubility -- like dissolves like
                  1.  Polar solvents dissolve polar solutes.
                  2.  Nonpolar solvents dissolve nonpolar solutes.
      hydration:  occurs when water dissolves a polar solute.
III.  Solids, Liquids, and Gases in Solutions
       Possible Solution Combinations

Solvent	Solute	Common Example
gas	gas	helium-oxygen (deep-sea diver's gas)
gas	liquid	air-water (humidity)
gas	solid	air-naphthalene (mothballs)
liquid	gas	water-carbon dioxide (soft drink)
liquid	liquid	water-acetic acid (vinegar)
liquid	solid	water-salt (seawater)
solid	gas	palladium-hydrogen (gas stove lighter)
solid	liquid	silver-mercury (dental amalagam)
solid	solid	gold-silver (ring)

        miscibility:  the ablity of two liquids to be mixed.
        example:   Water and acetic acid are miscible.  (vinegar)
                          Oil and water are immiscible.  (They don't mix)

IV.  Solution Equilibrium
       solution equilibrium:  occurs when the rates of particles leaving and returning to solution are equal.
       3 Types of Solutions
    1.  saturated:  when undissolved solute is in equilibrium with the dissolved solute.
     2.  unsaturated:  contains less than the saturated amount of solute for that temperature.
     3.  supersaturated:  contains more solute than a saturated solution can normally hold.

        solubility:  the quantity of solute that will dissolve in a specified amount of solvent at a specific temperature.
        SOLUBILITY CURVES
 
 
 
 
 

V.  Factors that Affect Rates of Solution
         1.  Agitation
              - increases solubility
              - brings solvent into contact with more of the surface area of the solute
        2.  Temperature
             - if temperature increases, solubility increases.
             - An increase in kinetic energy, increases the frequency and force of collisions of solvent and solute which breaks solute apart.
        3.  Particle Size
             - smaller particles dissolve faster because they have less suface area
VI.  Concentration of Solution
        concentrated solution:  large amount of solute in small amount of solvent
        dilute solution:  small amount of solute in large amount of solvent
                  Determining Concentration by Different Methods
                        1.  Molarity  (M)
                          Molartiy =  moles of solute
                                                liters of solvent
                           example:  What is the molarity of a solution in which 58 g of NaCl are dissolved in 1.0 L of solution?
                                           58 g NaCl  | 1 mol NaCl      =      1 mol NaCl
                                                               |  58 g NaCl
                                            Molarity =   1 mol NaCl
                                                                  1 liter
                                            Molarity = 1 M NaCl
                      2.  Molality (m)
                           Molality =   moles of solute
                                                   Kg of solvent
                            Example:  What is the molality of a solution in which 3.0 moles of NaCl is dissolved in 1.5 Kg of water?
                            Molality =  3.0 moles of NaCl
                                                 1.5  Kg of water
                            Molality =  2.0 m NaCl
                      3.  Normality (N)
                           Normality =  Molarity   x   total positive oxidation number of solute
                            Example:  What is the normality of 3.0 M of H₂SO₄ ?
                           Normality = 3.0 x total positive oxidation number
                          total positive oxidation number =  +1(2) = 2       This is because hydrogen's oxidation number is +1 and there are 2 hydrogens.
                           Normality = 3.0  x   2
                           Normality =  6.0 N
VII.  Types of Mixtures
        1.  Colloids:  mixtures composed of two phases of matter
                     Two phases are:
                      -- dispersed phase - particles are larger than particles in solution but smaller than suspensions
                      -- continous phase
              Classification of Colloids
               aerosols:  liquids and solids dispersed in gases.  Examples:  fog and smoke
               foams:  gases dispersed in liquids and solids.  Examples:  whipped cream and marshmallows
              emulsions:  liquids dispersed in other liquids or solids.  Examples:  mayonnaise - liquid emulsion
                                                                                                                                      cheese - solid emulsion
              sols:  solids dispersed in liquids or other solids.  Examples:  jelly and paint

        2.  Suspensions:  dispersed phase contains particles much larger than in colloids or solutions.
            *** Because the particles in a suspension are so large, the particles are suspended but eventally settle out upon standing.
 

VIII.  Properties of  Solutions, Colloids, and Suspensions

Solutions	Colloids	Suspensions
Do not settle out	Do not settle out	Settle out upon standing
Pass unchanged through ordinary filter paper	Pass unchanged through ordinary filter paper	Separatated by filter ordinary filter paper
Pass unchanged through a membrane	Pass unchanged through a membrane	Separated by membrane
Do not scatter light	Scatter light	Scatter light

      Tyndall Effect:  the ability of colloids and suspensions to scatter light
                                   examples:  1.  If a window blind in opened in a dark room, suspended dust particles scatter light.
                                                      2.  If a search light is used in the night air, light is scattered by suspended water droplets.
        Brownian Motion:  chaotic movement of particles in a colloid or suspension
                                           example:  Hitting two chalkboard erasers together allows you to see the chaotic movement of suspended
                                                            dust particles.

Colloids - Suspensions

Explanation

Suspensions are mixtures of particles that settle out if let undisturbed. Suspensions can be filtered, while solutions cannot. Colloids are a type of mixture whose particles are held together through Brownian Motion, the erratic movement of colloid particles. Colloids cause the Tyndall Effect, or scattered light due to Brownian motion. The size of the particles found in colloids is smaller than those found in suspensions and greater than those found in solutions. One commonly known colloid is milk.

Transcript

Alright, so types of mixtures that you're going to see one is homogenous mixture is also known as solution that the same through out and the next one is heterogeneous mixture and there're different types of heterogeneous mixtures where the particles are really big like solids or sands mixtures or things like that. But then there're ones that are very small such colloids and suspensions. Let's talk about those, suspensions are mixtures containing particles that settle out if left undisturbed meaning that like the particles are so large they have really big particles, they're bigger than 10 of the negative 6 which might seem quite small actually but compared to like atoms or compared to other particles typically in a solution which tenth to the negative ninth meters, they're actually quite large. Since that have large particles ad they have nothing to withhold them together they can be filtered, they actually can be separated out so types of suspensions that you'll see, that you'll come across that you might know if they're suspension or colloids or solutions even.

Blood if you leave blood, left undisturbed it will actually separate out, you can actually filter it out to separate it aerosols, corn start in water those kind of things are your types of suspensions. Another type of heterogeneous mixture is colloids, colloids are mixtures containing intermediate size particles held together through Brownian motion and Brownian motion we'll also get to in a second is what distinguishes suspensions versus colloids. So different types and examples of colloids would be milk where the particles are kind of big but not as big paint and fog where they actually stay together they don't filter out. So let's go and talk about what Brownian motion is, so Brownian motion is the erratic movement of colloid particles. So let's say take this picture for example we have out 2 larger colloidal particles like they could be proteins, they could be whatever and they're typically repulsed by each other, they have a repulsion but they might be attracted in other ways like other things within the solution, not the solution, the colloid will be attracted to it. So they have this like constant erratic movement of these particles. So how can we, if we're looking at 2 a suspension versus a colloid how can we like notice that one has Brownian motion one doesn't.

Well there's this thing called the "Tyndall effect" and the tyndall effect is that if you shine light through a colloid you're going to get a scattering of that light. So for example fog we know is a colloid, the reason you don't put your high beams on when you're driving through fog is because the Brownian motion of colloidal particles within the fog will shine that light right back in your eyes and actually like affect your driving negatively rather than positively. You'll actually see less of the road than if you put your regular low beam on or fog lights which are located lower which will then like light up the bottom of your driving. So that's the reason why, good example of real life example of fog and the tyndall effect.

But let's look at it here in the lab, let's talk about, so one of these guys is a colloid and one of these guys is a suspension. So if I shine my light through it, my light source which is this red laser beam, one of them should actually not be able to go through, the light should not be able to go through and the other one should. So let's actually test it out, so we have here if we put this through this side actually can't see the light source on the other side this is the colloid.

The colloid will stop the light it'll scatter the light and not allow it to go straight through. And this one you kind of probably already noticed that is a suspension is already starting to settle out but let's just test it out using the tyndall effect. Putting light through you can actually see, notice the light actually does go completely all the way through and it's got this like quite a little bit but it definitely like allows the light to go straight through whereas the colloid doesn't at all. So this tyndall effect through the Brownian motion and colloids and suspensions.

Freezing Point Depression

Explanation

Freezing point depression is a colligative property of solutions. Solutions freezing points are lower than that of the pure solvent or solute because freezing, or becoming solid, creates order and decreases entropy. Solutions have high entropy because of the mix of solvent and solute, so it takes more energy to decrease their entropy to the same point.

Transcript

Alright when dealing with solutions, you're going to be coming across a colligative properties and one of the colligative properties that you're going to see is freezing point depression and that says in a solution solute particles interfere with attractive forces among the solvent particles. And this prevents solution into entering a solid state. So essentially what they're saying is, because a liquid has like all these extra particles in it to make a solution, and it's not a pure solvent those get in the way of the intermolecular forces that make it a solid, a solid. Like the hydrogen bonding, dipole-dipole interaction and the London dispersion forces, those extra particles that are there kind of get in the way and they actually help to like lower the freezing point to get, to push those particles out so it'll be a pure solvent when it's actually frozen.

Freezing point states that the particles are no longer have sufficient kinetic energy to overcome intermolecular attractive forces so when those particles are there those attractive forces are not necessary. So like they're going to have to just push those particles out so they can actually have the inner particle attractive forces present. So let's actually talk about different substances and their freezing points. So we're talking about water which is a universal solvent and we know that water freezes typically at 0 degrees Celsius at normal freezing point. Now for every molal of substance I'm going to put within that water in a pure substance is actually going to drop the temperature the freezing point temperature by 1.86 degrees Celsius.

Benzene freezes at 5.5 degrees Celsius well higher than water and for every mol of substance that you have for a kilogram of solution the boiling point is going to drop even more 5.12 degrees Celsius and so on and so forth. So if you were to look at this, this is very similar to boiling point elevation but this set formula is exactly the same but there's some slight differences. So the change in temperature of the freezing point is equal to the constant that we had discussed, time similarity of this solution that we're dealing with, time is a Ben Hoff Factor and a Ben Hoff Factor is how much the particle actually separates in solution. So we're talking about ionic compounds, they separate into solution depending on how many particles they have or how many ions they have in that but molecular compounds don't at all. So let's actually put that in action.

Alright so what a freezing point of a 0.029 molal of NaCl aqueous solution so we know it's aqueous and the aqueous tells us that our solvent is water. So we're going to say our delta T are changed in temperature to freezing point is going to equal to the constant of water which is 1.86 degrees Celsius for every molar. And the molar solution is 0.029 and because it's NaCl I know it's ionic for every one molal it's going to actually separate into 2 substances Na plus and Cl minus. So we're actually multiplying this by 2, we have substances when it's in solution. So when you multiply all of these together you get 0.11 degrees Celsius and we're going to say alright our original freezing point is 0 it's going to lower by 0.11 and so our new freezing point is 0.11 degrees Celsius negative because it dropped that much. So we can actually like talk about, when you think about when it snows outside and the reason that you put salt on the roads there isn't even salt on the roads that's actually going to lower the freezing point so there's not going to be sheets of ice on your drive way or on the roads or on the side walks. So that's why they use salt and they actually use calcium chloride typically which is actually better than sodium chloride because this actually breaks up into 3 particles so it'll drop the freezing point 3 times much as another solute would. So this is an example of frizzing point depression.

Boiling Point Elevation

Explanation

Boiling point elevation is a colligative property of solutions. Solutions boiling points are higher than that of the solute or solvent because the vapor pressure of solutions is lower. A boiling point is when the vapor pressure of the solution becomes equal to the external pressure, so when the initial vapor pressure is lower, it takes more heat to elevate the vapor pressure to the same point.

Transcript

Alright so one of the properties that you're going to see in class, is we're going to talk about boiling point elevation and that is the temperature difference between the boiling points of a solution and that of a pure solvent. Now as you can probably guess from the title it's actually going to increase from the solvent to the solution, solution is going to have a higher point than the solvent. And why is that? Let's define point first. Boiling point is when vapor pressure of whatever substance you're dealing with is going to equal the atmosphere pressure of the pressure around it. Okay great let's say we have a pure solvent, let's say we have a beaker of water. We know that the surface of water, there's particles that are going to escape into vapor pressure into vapor back and forth back and forth and there're going to have some sort of pressure up here.

And as we heat the substance that pressure is going to raise due to kinetic energy increasing and it's going to eventually reach the atmosphere pressure and then boil okay. So let's over to a solution that you'll put something in it salt, sugar or whatever your solute maybe and to make that solute to actually dissolve these water molecules are going to surround it. So these water molecules are actually really attracted to the solute that's in the solution. So what's going to happen is they'd rather be around these dots or these solute particles than escaping into vapor, so it's already going to start off at a lower vapor pressure. So they actually have to go increase kinetic energy a lot more to reach that vapor pressure to equal the atmosphere pressure because they're already starting in a lower place. So that temperature is going to have to be a lot higher or depending on what the solvent is to that atmosphere pressure and then boil.

Okay so let's talk about different solvents and how the solutes are actually affected by that. So let's about water, we know water boils at a 100 degrees Celsius and it's found for every molar of solute that you put in there it's going to increase the boiling point temperature by 0.512 degrees Celsius. Benzene boils at 80.1 degrees Celsius and every molar or every mole per kilogram of Benzene you put in is going to raise the temperature 2.53 degrees Celsius or boiling point temperature that much and so on and so forth. So basically, in a nutshell, we have a formula here called the boiling point elevation formula and this is a change in temperature of the boiling point that delta T is going to the changing temperature times that constant that you see for every molar that's in the amount of degrees Celsius is going to increase.

Times on molar, in the morality which is mole of the solute divided by the kilograms of the solvent and we're going to multiply it by the Ben Hoff Factor or how much it's actually dividing out for a number of electrolyte actually dissolve and separate themselves in water. So if you have 1 molar of let's say NaCl you're actually going to have 2 molars of particles. So that is what we call a Ben Hoff Factor. So let's actually look at that in real life. It says what is the new boiling point of a 0.029 molar of an NaCl solution? Okay so we put salt in water and we're going to know the new boiling point of this water. So we can just look at our formula and say our change in temperature is going to equal the constant and now we're dealing with water because we know it's an aqueous solution, so our solvent is going to be water so when we look at out table and for every molar it's going to raise 0.512 to degrees Celcius that's what I'm going to say 0.512 degrees Celsius per molar.

We know our molarity is 0.29 and out Ben Hoff Factor is going to be 2 because for every molar of this is going to actually break up, for every 1 mole of NaCl is going to break up into 2 moles of particles. So our Ben Hoff Factor in this case is going to be 2, okay so we multiply all these together and we should get 0.030 degrees Celsius. This is not our new boiling point, this is just a change in boiling point, so our new boiling point is, original boiling point was 100 degrees Celsius, it's going to change or increase because or increase because we know it's elevating by 0.030 degrees Celsius. And so our new boiling point is 100.030 degrees Celsius, the new boiling point of this solution. Okay and this is why you can think about when you're cooking and you're staring some salt into a pot of water not only are you for flavor you're actually increasing the boiling point of that water. So when you're dealing with like let's say pasta you actually can cook pasta a little bit faster because your boiling point is actually a little bit higher than 100 degrees Celsius and you can cook it at something a little bit higher if you put more in, the more the higher the higher the temperature is going to be. But don't put so much salt in there because it's going to be over salted and over flavored. But either way that is boiling point elevation.

Vapor Pressure Lowering

Explanation

Vapor pressure lowering is a colligative property of solutions. The vapor pressure of a pure solvent is greater than the vapor pressure of a solution containing a non volatile liquid. This lowered vapor pressure leads to boiling point elevation.

Transcript

Alright so one of the gas properties that we're going to be discussing is vapor pressure lowering. So and we're talking about a pure solvent like water, at the surface of water it will evaporate and go into the vaporous state. So the water molecules at the surface are going to go back and forth, back and forth in equilibrium into the vapor around it and back to liquid state. So what we call vapor pressure is going to have most of pressure of gas on top of it. But let's say we put something in it and make it a solution, okay so I put some sort of like particle in this, that made of that will be able to dissolve and what happens when something dissolves is that the water molecules are attracted with whatever it is that's being dissolved and they would rather be around that particular particle than escape into the vaporous state. So actually the number of particles that are going to escape into the vapor is actually going to be lowered. Because depending on how many particles of solution I have, so the vapor pressure this pure solvent is greater than the vapor pressure of a solution containing a non-volatile liquid. So why did I say non-volatile liquid? When something is volatile is actually able to evaporate very quickly. So if I have something that's volatile in here that will actually make it escape a lot easier.

But the fact that these are non-volatile meaning that they like to be in liquid rather than gaseous state, then they're going to make the vapor pressure within this solution lower that it was originally. Okay, so let's talk about what the different particles we could put in there. That actually makes a difference, so we have sodium chloride, we have calcium chloride and then we have sugar molecules. Now the main difference between these guys is, these guys are electrolytes meaning they'll separate in solution and this guy is not, it's not electrolyte meaning it'll stay together. So when I say 1 molar of sugar I mean 1 molar of particle. So here when I have this, this actually will end up being 1 molar okay fine. So when I put this in solution, this will break up into 3 different ions, 1 calcium and 3 chloride ions. So when I put 1 molar in it actually ends up being 3 molar of particles.

Here I have sodium chloride this is 2 particles, 1 sodium ion and 1 chloride ion so when I put 1 molar in there it's actually going to break up into 2 [moles] so this one it's going to actually affect the vapor pressure than most, then this guy and lastly this guy. So it actually does make a difference of what you put in solution and how much you put in there and how much the vapor pressure is actually going to lower.

Colligative Properties

Explanation

Colligative properties are the properties of a solution as a whole and depend on the concentration. The colligative properties include freezing point depression, boiling point elevation, vapor pressure lowering and osmotic pressure.

Transcript

Alright. Let's talk about colligative properties solutions and colligative properties are a collection of physical properties of the solution that are affected by the number of solute particles. So properties of solutions are going to be different from the properties of the actual solvent by itself. And the things that, there are four things that are going to be affected by this. One's going to be vapor pressure, and actual vapor pressure of the solution versus its solvent will lower, the other one, another one is boiling point and the boiling point of the solution will be higher than the boiling point of the solvent alone. The freezing point of the solution will be lower than the actual solvent by itself and the osmotic pressure is going to increase, compared to a solvent by itself. And this will depend on the molality, molality not molarity, molality of the solution. And molality is moles of solute over kilograms of solvent, and don't forget the density of water is one gram per millilitre. So we can just change number of litres to kilograms, they'll be the sa-, it should be the same conversion because the density of water is one. And don't forget that the unit for that the symbol for molality is a small m.

So thi- this will depend, this molality is what's going to show how much all these guys are affected. So let's talk about different types of things that should be in a solution. So you have electrolytes and as you should remember electrolytes are things that conduct electricity in solution or ionic compounds and ionic compounds as you know will break up in solutions. So for example we have sodium chloride. When you put sodium chloride in water, it will actually break up into sodium ion and a chloride ion. So if you put one molal of sodium chloride in solution, you'll actually get two molals of product, of particle in your so- in your solution. So it will actually break up and you know that your solute is going to be an electrolyte.

Also if it's an electolyte that has like calcium chloride that has more than you know just two particles those will break up and actually three particles of every one. So if it's one molal of calcium chloride you actually need three molals of particles which will affect this more than sodium chloride will affect it even if they were the same amount. Also, non-electrolytes, non-electrolytes do not break up in solutions so however amount you put in solution is the amount that you actually, the amount of particles will actually be in the solution.

So, for example here is table sugar which we know is a molecular compound. If we put one molal of a non-electrolyte, it will just be the same. These are not going to break up into its atoms. So it will stick together. It will break up from each other but it'll stick together in terms of the compound. So given one molal of a non-electrolyte, you are going to end up with one molal of particles of non of the non-electrolyte. So this is actually going to play a big part in dealing with the co- colligative properties, the solutions, boiling point, freezing point and vapor pressure and osmotic pressure. And we'll talk about others in each separate video.

Molarity - Molality

Explanation

Molarity and molality are units of concentration. Molarity measures concentration in terms of moles per liter. A one molar solution has one mole of solvent for every one liter of solution. Molality, on the other hand, measures concentration in terms of kilograms per liter. A one molal solution has one kilogram of solvent for every one liter of solution. Molarity and molality are important in reactions in aqueous solutions and affects reaction rates.

Transcript

Alright. So we're going to talk about different ways you can express concentration, and concentration is the measurement of how much solute dissolves in a specific amount of solvent or solution. So like let's for example say you're talking about getting juice from concentrate. Concentrate means it's very highly concentrated have a little bit of so- sorry a lot of solute compared to the little bit of solvent. But if you want to make it less concentrated, you're going to make it more dilute. The opposite of concentrated. You're going to add more solvent making it the measurement basically a ratio of solute to solvent.

So the different ways you can actually measure that, several different ways are the ways that we're going to talk about today. One being percent by mass and then we're talking about like solid solutions like metal alloys and things like that or you're doing percent by mass and that's mass of a solute over the mass of a solution. Obviously percentage means multiply by 100. And you can also do percent by volume and these are things we're talking about liquids. And you might even see symbols like this, excuse me. You might even see symbols like on like a, on like a product. This means percent by mass. This means percent by volume. You might see like 70 v v that means percent by volume then times the solute over volume solution. Molarity is very very common. You're going to see this a lot in Chemistry class. Talking about mols of solute versus litres of solution. Another way to like abbreviate that or shorten that is you might see a big m. When you see this big m you know it's mols over litres. Another one is molality and they are very easily confused but molality is mols of solute over kilograms of solvent. Very very different. These are just used in different ways. This is little m. Don't mix it up with metres but it is little m.

Lastly we're going to talk about this mol ratio. Mol ratio is an easy way to talk about the ratio of mols of solute versus mols of solute plus solvent. It's an easy way to go from, you know to kind of like go from any one of these. It's a good like converter. So these are the five different ways you're probably going to see and the way you're going to actually use these a lot of times it's through when you're talking about dilutions.

Let's go to dilute solution and what that means. So this could be of highly concentrated solution. Let's say this is acid and 12 molar acid is actually extremely concentrated. If you've got your hands on it like any skin on this, it would burn you to no end. So let's see how this is my stock solution meaning that I'm a Chemistry teacher and the back when, before I do lab demos, this is what I'm going to use. So I don't want to use that with you guys, I don't want to use that with my students because it could hurt them. So I'm going to dilute it. Meaning I'm going to add water. Okay. It means I'm adding liquid but I'm changing the concentration. I'm keeping the same number of mols, my volume is getting bigger. So my concentration is getting smaller. So the mols of this I didn't change the amount of substance I had, I just changed the water. So the number of mols of whatever it is my volume is, is exactly the same. So another way to dilute it, if I rearrange my molarity solution, my molarity equation, I find that the mv because mols are the same, mv of the first one, molarity times volume of the solution times equals the molarity times the volume of the second solution because the mols are the same. Let's do a problem based on that.

Let's say I have a, what is the volume of a two molar solution of a calcium chloride which you use to make a 0.5 litre solution of 0.3 molar calcium chloride. Okay. So, basically I want to make this, I want to make a 0.3 molar solution. I want to make 0.5 litres of it. I only have two molar solution. What I'm I going to do? Well, I take that two molar solution. I'm going to multiply to figure out the volume is. Multiply by 0.3 molar, that's my second molarity times my volume. And I do the math and I end up with, sorry. I end up with 0.075 litres.

So what I'm going to do, is I'm going to take my 0.075 litres or 75 millilitres of my 0.2 solution and add 25 millilitres of water to make my 0.3 molar solution of that. Because I know the number of mols and of calcium chloride's the same in both scenarios. So concentration, the ways that they can use it is numerous but one of the ways is that they can use it in diluting solutions.

Solvation

Explanation

Solvation, also called dissolution, is the process of surrounding solute with solvent. It involves evening out a concentration gradient and evenly distributing the solute within the solvent.

Transcript

Alright. Let's talk about solvation. Solvation is the process the process sorrounding solute particles with the solvent. So when we are dealing with solutions the solvation actually is what's actually happening within the solution. So we know that the universal solvent is water and we're dealing with water as a solvent. We're going to call that hydration. So you might hear that as well.

Alright. So the phrase like dissolves like, what does that mean? It might be something that your Chemistry teacher might have talked about in class. That actually means, we're dealing with polarity. So we're going to talk about water as a univ- because it's a universal solvent.

Water is a polar substance. Meaning that it has a negative end and a positive end. And so, we want to, anything that's polar has a charge, can dissolve within water. So when we're taking something that's ionic, we know ionic compounds are held together through electrostatic forces and they actually do have positive and negative ions within them, that's what makes the ionic compounds. So when you drop it into water, what happens? Well, these the ions are actually going to separate from themselves and they're going to be surrounded by the water particles. So that oxygen is negatively charged, they're going to surround, they kind of like attack this positive cation and that this is actually the solvation process. And the negative ones are going to be surrounded by the hydrogen. So the more that they're actually like pulling off and actually going away, the more the ions you're exposing and eventually all these will be exposed in water making the whole thing dissolved. Okay? The same thing happens within a molecular solution but you have to make sure the molecular solution is a polar substance. If it's non-polar, this actually won't be able to happen. And what happens is this is sugar, this is table sugar and the table sugar has oxygen and hydrogen bound together. And we know that hydrogen bound together we know will create hydrogen bonding. And so when water comes close to it, this actually will, it will surround it here and here. All over the place. You'll see this bond or this attraction for water with sugar. So when the sugar separates from itself, each sugar molecule separates from itself, water will actually surround it making it dissolvable or dissolve in the solution. So this is the process of solvation. Okay. It's only polar molecules can do this, do this, it's non-polar. they don't have water is not attracted to it at all. So it's not going to surround it at all, so it's not going to dissolve. So that's the idea of the like versus like, sorry, like dissolves like means.

So what kind of factors affect solubility rate? Well, meaning how fast can it dissolve? So if we have the more surface area exposure that this compound has to the water or whatever it's being dissolved in, but they will obviously dissolve faster. So if you think about like a sugar cube versus granular sugar, the sugar cube will take a little bit longer to dissolve than the granular sugar which will be much quicker because it's more exposed. Stirring or agitating it obviously that's why you're - when you're thinking about dissolving something you stir it, that makes sense. That's actually making it more exposed as well. And then heating. When you heat something up, that creates more kinetic energy and things are moving around much more. So, you're able to dissolve this solvation processes if you are able to at a quicker rate.

I come across words like saturated, sorry. Unsaturated, saturated and supersaturated, when something, only a certain amount of substance is able to dissolve in, let's say 100 grams of water. So if it's not to that point, there's actually a point. It's called saturation point and if you get if you don't get there, if you can still do continous to dissolve more sugar or whatever we're talking about in that water, that solution is what we're going to call unsaturated. Meaning you can continously add more solute into the solvent, that solution until you get to the saturation point and once you hit the saturation point, it's going to be saturated. Meaning you can't add any more in there. And supersaturated, what's that? That means like if you heat up the substance the more, the hotter the substance is, the more it's able to dissolve something. So if you heat up the substance and then put in like a certain amount of sugar and then cool the substance back down, that certain amount of sugar is actually too much that cool temperature, but it will stay in solution because you'd have originally put it there. And if you agitate it just a bit, it will start crystallizing out and that's how like things like rock candy is formed or things like that. It's from supersaturated solutions.

And gas solutions are actually a little bit unique too. Let's think about soda as an example. Soda is a typical like gaseous solution within a liquid. And so we talk, let's talk about pressure. So if you have a two litre container of soda, how are you going to make sure that the gas doesn't escape from that soda and make it a non-gaseous solution. Well, you probably want to keep it cold. The colder a solution is, the less movement those gas particles will have to escape. So the colder a solution is the more it will stay in solution and not only that. The more pressure you have in that solution the more those gas particles will want to stay in that liquid. they're not going to escape. So those kind of things are unique within gases.

So the more, actually, the higher the temperature of the gas is, the solution with the gas is, the worse it is for that gaseous solution. It will actually like not be able to dissolve as much. So it's kind of the opposite as you would think to other types of solutions. So it is unique in that way.

So these are the properties and the reasons things dissolve in other solutions and it's called solvation.

Types of Solutions

Explanation

Solutions are homogeneous mixtures. Different types of solutions have solvents and solutes in different phases. Solutes are dissolved in the solvent. In a solution in which carbon dioxide is dissolved in water, the water is the solvent and the carbon dioxide is the solute. Two important concepts in studying chemical solutions are solution concentration and solubility equilibrium. Properties of solutions as a whole are called colligative properties.

Transcript

Alright. Let's talk about the different types of solutions that you'll see. Don't forget a solution is actually a homogenous mixture meaning that there are actually things within a solution not bonded together. They're just kind of, they're attracted to each other in a way that this makes it the same throughout.

Different types of words that you'll see when you're dealing with solutions are words like solute and solvent. A solute is actually dissolved within the solvent. So whatever is being dissolved is a solute, what it's being dissolved in is a solvent. The universal solvent that is that you'll come across is water and that makes sense since most things are dissolved in water, but there are other types of solutions that you'll see too. One being gas dissolved in gas and that think, you know think about the air we breathe is a mixture of gases. Gas dissolves in liquid and soda is a great example of that, the ga- carbon dioxide bubbles. They're actually gaseous or dissolved in liquid of the soda.

Then we have the liquid dissolved in liquid. Juice is a good example of that. some fresh squeezed juice or something along those lines. You can have solid dissolved in liquid which is like our sugar. Our sugar is going to be our solid form, dissolved in our liquid tea, that would be our. They do actually see this actually quite often solid dissolving in liquid.

Solid dissolving in solid. Steel or any type of metal alloy actually is a mixture of solids.

Gas dissolved in solids like foams or marshmellows. So these are different types of mixtures that you'll see in different mediums that they're actually in, and there are different words that we actually may come across too. Soluble versus insoluble and when you're dealing with double replacement reactions you'll have a precipitate. Precipitates are insoluble meaning they come out of solution. If something is soluble, it means that it's actually dissolved within the solution it doesn't come out, it's still within the solution.

You may see the word miscible versus immiscible. This miscible you can kind of sounds like mixable. That's exactly what it means. If something's miscible, they actually it's something they can mix together. If they're immiscible like oil and water, they are actually not able to mix together. So these two, these four words you might see quite often when you're dealing with types of solutions.

Chemical Solutions-Osmosis

Chemical Solutions

Osmosis

Explanation

Osmosis is the diffusion of water molecules across a semi-permeable membrane from an area of high concentration to an area of low concentration. In cells, osmosis occurs across the cellular membrane to keep a cell from becoming flaccid (not enough water) or turgid (too much water).

Transcript

So osmosis is the diffusion of water through semi permeable membrane. And this is something that a lot of teachers will spend a little bit of time on but then they'll hammer you with on a test. Because it is really kind of unusual when you first approach it a lot of kids don't really get what it means, but to the teachers it's pretty obvious. So let me see if we can help you figure this out. So what this is all about is a special case of diffusion, you know diffusion is but why is it specifically water? And that's because water has the ability to go through a number of different membranes. This black line here that I drew in I'm using it to represent some kind of membrane. Well it's a cell membrane, a special kind of plastic or dialysis tube whatever.

You'll notice that there're small holes in it, I've used red to represent sodium ions they're too large to fit through these holes I've used blue to represent water. It's easily small enough to fit through so what happens is that like all other molecules water does diffusion and you know diffusion is the movement from an area of high concentration to an area of low concentration. So what we see here is that the water starts to diffuse from there's 90% water to where there's 70% water why because that's what molecules do. So the water tends to go like this now does all of the water go no. if I look at just this one area here if this is 90% then let's suppose I'll make up a number of 90 molecules of water going that way at the exact same time 70 of the water molecules over here are going back. 90 go this way 70 go that way there's a net difference of regaining 20 on this side. Now the sodium it trys to diffuse but like my friend 2 ton Tony he can't fit through the opening of the door way so he gets he bounces off and stays over here. This sodium tries to move but it bounces off because it can't fit through.

And that means that ultimately we see the water moving from this side to that side while the sodium stays and as this will continue going until the water wound up diluting out the sodium on this side equilibrating its concentrations. That's osmosis now as I said teachers love to ask a bunch of questions about this so I'm going to give you the 3 basic questions that they'll generally ask about osmosis. Now to do this I'm going to be introducing some vocabulary, this stuff about tonic. Now tonic is a root word that means pull I think in Latin I'm not sure it's all Greek to me. But what I'm going to do is I've drawn here 3 different circumstances this little black circle here that's in my beaker of water let's say that's a red blood cell alright, all of these red blood cells have 80% water 20% other stuff. I don't care what it is it's just not water, so they're all the same however I plunked them into 3 different kinds of solutions.

This is pure water as you can see its 100% water, 0 stuff, this is 80% water 20% stuff, this is 60% water 0% stuff. Let's take a look at this one right here every second some of the water molecules move out but some of the water molecules move in, what's the net change? Well 80 move in as 80 move out so there's no change, the salt and other proteins and other things that cannot pass the membrane they stay so we're only looking at the water because this is osmosis. Because we see no change because both sides have the same amount of pull tonic ability we call them isotonic. Isotonic means that it has the same concentration of water and solutes let me use the red pen to represent solutes alright. What about this situation, well 80% inside 100% outside 80 move out 100 move in.

What's going to be the overall change, water is going to continue moving in and in and in this is going to make the cells to go, swell up and a red blood cell doesn't have the ability to do anything else like start shoving some of the salt out so it's going to pop. Because this outside water solution doesn't have any ability to pull water into it in fact it's pushing water out it's called hypotonic. Because it has lower than the normal amount of pull compared to the cell. It's always in comparison, it's kind of like saying taller, somebody can be taller or shorter than someone else. You can ever just say that person is taller, so hypotonic says, means it has less concentration of solute more water. Okay if we take a look at this one this has 80% water on the inside, 60% water on the outside because there's a ton of other stuff say salt or whatever. So we have our water move out 80 move out 60 move in we have a net movement out of the cell. So in this case instead of popping sometimes it's called liaising.

Here it starts to shrivel, if it's a plant cell they don't say that plant shrivelled they say that plant wilted. And if it's a red blood cell and the other word that you'll see beside shrivel is crenate I don't know why people keep coming up with these words just to confuse you guys. So this one, this solution outside the cell has a greater pull and excessive pull and it pulls all the water out tonic means pull. What is excessive what's the root with that means excessive? Well you have known that hypo meant below as in hypothermia. What means excessive hyper like your brother, he is hyper active so this has a hyper excessive pull. So hypertonic solution has more solute less water, so that's how you do osmosis alright so in an isotonic solution a cell will just stay the same it'll be an equilibrium. A cell in a hypotonic solution will swell up and if it's a plant cell its wall will keep it from bursting but animal cells will typically pop especially red blood cells. And a hypertonic solution a cell will shrivel or wilt.

Phase Diagrams

Explanation

Phase diagrams graphically depict the state of matter in varying temperatures and pressures. The x-axis of a phase diagram is always temperature while the y-axis is always pressure. There is a point on a phase diagram called the triple point at which all three phases of matter exist simultaneously.

Transcript

So we're going to talk about phase diagrams, phase diagrams are graphical interpretation of pressure versus temperature.

okay so basically you can fig- you can figure out if you don't know the temperature or pressure of any substance you can figure out which phase it is using a phase diagram so and we also can see what pressure and temperature is needed for different phase changes going from solid to liquid to vapour so let's like let's look at this graph in itself and try understand what this is saying, so we have 3 major lines here we have one going up this way and we have one coming down here so this, this separation here between the solid and vapour, this point where if any time they crosses this line, we're going to have sublimation occuring. Any time it crosses this line here it's going from a solid to a liquid or liquid to solid either way, we're going to have either freezing or melting going on. Notice that one atmosphere which is that we know one atmosphere in 0 degree Celsius is our normal freezing point and melting point.

Any time that you have temperature pressure within this region your substance will be so- will be a liquid anything down here is starting to be a vapour or gas. We have the reason why we call this normal is because one atmosphere we call this normal atmosphere pressure and we also know this is our boiling point at 100 degrees Celsius we know the water at a 100 degrees.

Okay there are a couple other things in here that you might not know, one is when they're all in a centric here that was what we called the triple point actually there is a there is a temperature and pressure at which all three phases will coexist together and we're going to call that the triple point.

There's also something up here called the critical point, the critical point is a place where anything any pressure or temperature that's higher than the critical point there's no way liqu- the liquid phase will exist it will only the substance will only exist in the gaseous phase so this actually kind of tells us a lot of things in terms of subst- which where the substance is in terms of the state of matter that is in.

Alright, so let's look at something else, this was water, so let's look at something else, it's another phase diagram we can talk about each substance has its own phase diagram and this is carbon dioxide and there is a couple differences in carbon dioxide than there was in water one main difference if you look at both of them, one main difference is this line here okay notice this line versus this line okay water is very unique in the sense that it will have its solid form is more dense sorry lesser dense than it's liquid form so as we continue to add pressure to the solid formula it will actually change to a liquid state and if we continue to add pressure you notice it'll go up its going to eventually change into liquid state because liquid state actually more dense than the solid state whereas in most substances if you increase pressure it's going to stay in it's solid state because the solid state is actually more dense than liquid state so you're going to see a line looks like this.

Alright notice carbon dioxide we're going to call that dry ice carbon di- the reason we call it dry ice is because here we have some dry ice actually with us, let me get some out for you alright the reason why we call it dry ice is because we have it a solid form but yeah it's sublimate meaning it goes straight from solid to gas okay this is used sometimes in like movie studios or to make that gas the gas appearance the gaseous appearance, so if you notice this goes straight from the solid state skip skipping the liquid state going straight to the solid state I mean straight to the gaseous state and we call this sublimation. Alright so we're going to actually demonstrate with triple point will look like by putting dry ice in this container and seeing if we can make it liquid somehow so let me try and put my safety goggles on because safety is always first and we are going to put some little pieces of, we're going to, this is really cold so I really shouldn't be doing that but it's not going to hurt me unless hold it for too long, little pieces [IB] alright so what I'm going to do with this oops I'm going to, alright so what I'm going to do with this is I'm going to close this up and so now I have my dry ice in here and my dry ice is sublimating as you saw before and so it's going straight to a gaseous state meaning and also making a pressure inside this container much higher okay so if we look back and if you remember back at our graph you notice this is continue still sublimating, okay and what we want to do is we want to continue the to make the pressure high enough where it will sublimating and actually go to a triple point where we can see all these in the same time and how we're going to do that is going to lower the temperature and notice from the side we notice that inside this container is really cold as you can see from the impressions of the bottle and also as you can imagine its getting really pressurized in here because the gas is continuing to sublimate and this is going to increase the pressure inside the bottle so hopefully we'll get to the triple point, we can shake it up a little bit as you can see like on the side over here maybe you can see like the liquid inside here so now all three phases are existing making this the triple point. We can see in the bottom like maybe some liquid here, liquid CO2.

Alright so now we're going to open the bottle and see what happens whoo and this is all the gas that's escaping there's a tons of pressure in here they're 5 atmospheric pressure which is 5 times the regular atmosphere that we actually living in at sea level so actually let's go back here and just make sure we understand what will happen is here is our normal temperature around here we know that [IB] sorry carbon dioxide is in gaseous state what we did we lowered the, the temperature by putting it inside the container and make increase of pressure as it sublimates inside the pressure and we're able to create the triple point where all threes substances coexist, so this graph actually tells us a lot of things about the different types of substances in different types of forms that they can actually have in pressure and temperature and how they're related.

Phase Change

Explanation

Phase changes are the transformations from one state of matter to another due to thermodynamics. The processes of phase change between solid and liquid are called melting and freezing. Phase changes between liquid and gas are vaporization and condensation. Phase changes between gas and solid are deposition and sublimation. Phase changes can be spontaneous or non-spontaneous.

Transcript

We're going to talk about phase changes, going from different forms of matter; for the solid, liquid gases and how they interact with each other and how they change from one phase to another. So let's use this as an examp- as a good diagram to show us how these things interact with each other.

Alright, so if we're going to get from solid to a liquid we're actually call that melting which I know we've heard that word many times before and that actually requires energy, we need heat to melt something, so we're going to call that endothermic process meaning that it requires energy or requires heat for that reaction or that that phase that to occur. So if we're going to from liquid the opposite from liquid to solid, we're going to call that freezing which I know we've heard many times before too. That actually releases some sort of energy, it's going to be an exothermic process, meaning it releases energy. These guys are opposite of each other, melting and freezing, of the same thing.

Let's go over to liquid and gases. If we're going from liquid to gas, we're going to call that vaporization; we're going to vaporize that particular liquid. That actually, requires energy as well. We need heat or some sort of energy to make that happen. So we we're going to call that endothermic process. The opposite will be cond- condensation; when we're are condensing something from a gas down to liquid and that's an exothermic process meaning that's going also to release some sort of energy.

There are rare instances where substances will go straight from the gas phase to the solid phase. It doesn't happen as often as you probably know but they do happen with different substances so if we're going from the gas phase down to the solid phase, we're actually going to release that sort of energy because we know gas is in high has higher energy than solid phase so we're going to release that energy we're going to call that process deposition. The opposite would be sublimation going from solids to a gas we've seen this probably before when ice or solid CO2 maybe iodine crystals they go from this solid phase straight to the gas phase skipping over the liquid phase that actually it releases sort releases some sort of energy and that we call it endothermic process.

Alright so let's actually look at this in a different way this you might see more often in class. This is actually a different a graph describing all those things that we just talked about. Alright so on the x ax- sorry on the y axis we have temperature on the x axis we got x axis we're going to have energy okay so we know in this case we're going to talk about water the phase change of water and we know that below 0 degree Celsius that is in solid phase okay? So as we increase energy, our temperature of that solid is going to increase until we hit 0 degree Celsius which we know it is melting and freezing point s the increase if we increase energy it's going to melt and if we're going from liquid to solid it's going to start freezing but notice the temperature it's not changing even though we're increasing temperature why is that? Well that energy that we're pumping into the into this solid molecule this substance is actually being used to break up those intermolecular forces that are holding it together in a solid so here's the picture water and these blue dots are the hydrogen bonds that are holding it together in a solid so because solids have more hydrogen bonds than liquids that energy is going to be used to break up some of those bonding some of those forces that are holding it together. Then as you go from 0 degree Celsius to 100 degree Celsius we're going to be in a liquid phase all that and all the energy is going to be used to increase the temperature of that particular liquid in this case water. And here we have the same thing we have this plateau and 100 degree Celsius we know that is it's vaporization point or it's con- condensation point again it's flat and again that energy is being used to break apart more of these hydrogen bonds once at 100 degree Celsius these bonds are going to be pretty wear because their energy is being used to break them apart and have them flow around allover the place and then up at higher temperatures above it's always going be in a gaseous phase.

If we go straight from the solid to a gas which water doesn't do if it were to we would call that sublimation going to skipping this liquid phase completely, if we're going from gas to solid we're going to call it deposition we're going to complete again skipping that liquid phase so this actually cycle talks about the different phase changes that substances tends to undergo.