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.