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Define the terms heterogeneous and homogeneous . - PowerPoint PPT Presentation


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Warm up. Define the terms heterogeneous and homogeneous . What do you think is the difference between dissolving and melting?. Warm up. Heterogeneous: not uniform in composition Parts are usually visually distinct Homogeneous: uniform in composition

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slide1

Warm up

  • Define the terms heterogeneous and homogeneous.
  • What do you think is the difference between dissolving and melting?
slide2

Warm up

  • Heterogeneous: not uniform in composition
    • Parts are usually visually distinct
  • Homogeneous: uniform in composition
    • Parts are not visually distinct due to being well mixed at the microscopic or submicroscopic level
slide3

Warm up

  • Melting: one substance changing phase from solid to liquid
  • Dissolving: two substances mixing to form a solution
solutions

Solutions

Chapter 15 (write the red!)

slide5

Objectives

  • Distinguish between homogeneous and heterogeneous mixtures and explain how dissolving is different from melting.
  • Define solute, solvent, solvation, dissociation, electrolyte, and aqueous.
  • Identify and compare the nine different solute-solvent combinations.
  • Compare solutions, suspensions, and colloids.
slide6

Defining solutions

  • A little review! Chemists often work with mixtures.
  • Mixtures consist of two or more substances, physically combined, each of which retains its properties.
  • Mixtures tend to be
  • Parts of heterogeneous mixtures are too large to mix uniformly.

HETEROGENEOUS.

slide7

Mixtures

  • When particles making up a mixture are small, they can be uniformly mixed or intermingled.
  • There is a(n) _________ relationship between particle size and uniformity of a mixture.
  • When the particles are small enough and thoroughly mixed, the result is a HOMOGENEOUS mixture.

inverse

slide8

The terminology of solutions

  • A solution is a homogeneous mixture – one substancedissolved in another. These two parts have special names:
    • Solute: the substance being dissolved. Generally, the solute is present to the lesser extent; if the two substances were originally in two different phases, the solute is the one that changes phase.
    • Solvent: the substance in which the other is dissolved, and is generally present to a greater extent.
slide9

The terminology of solutions

  • The process of dissolving (ordissolution) also has two “parts”:
    • Solvation: the process by which solvent molecules are attracted to and associate with solute molecules
    • Dissociation: the process by which an ionic compound splits into its component ions

**Do not confuse dissociation with ionization!

slide10

The terminology of solutions

  • Example: salt dissolving in water
    • Solute: salt (NaCl)
    • Solvent: water
    • Solvation: water molecules being attracted to and surrounding salt ions (salt is solvated by water)
    • Dissociation: salt ions breaking apart from one another (the salt dissociates)
slide11

The terminology of solutions

  • Some solutions conduct electricity
  • Electrolytes: compounds that conduct electricity in aqueous solution OR in the molten state.
    • All ionic compounds are electrolytes because they dissociate into ions, which carry charge.
  • A compound that does not conduct electricity when either aqueous or molten is a nonelectrolyte
    • Sugar is a covalent compound, and does not form ions when dissolved.
slide12

The terminology of solutions

  • If a solution is made in which the solvent is water, it is known as an aqueous solution.
slide13

The terminology of solutions

  • When the solvent in a solution is an alcohol, it is known as a tincture. A tincture of iodine is a solution of iodine (solid) in alcohol.
slide14

The terminology of solutions

  • Two liquids are “miscible” when they can mix in all proportions and form a homogeneous solution (ex: water and milk)
  • Opposite: immiscible.
slide15

There are 9 types of solutions

Alloys (steel, brass, etc.)

Kool Aid, salt water

Smoke

Dental Amalgam

Antifreeze, rubbing alcohol

Fog

Lava

Coke, 7 Up, Pepsi

Air

slide16

Careful!

Solutions should not be confused with suspensions, colloids, or emulsions

slide17

Comparison

Atoms, ions, small molecules

No scattering

Stable, no separation

Particles not retained on filters

Homogeneous

Kool Aid

slide18

Suspensions

    • Finely ground particles (larger than 100 nm), when placed in a solvent, can become “suspended”
  • Usually visible to the naked eye (scatters light)
  • Will settle out due to gravity in time
  • Can be purified byfiltration
  • Considered heterogeneous
slide19

Colloids

  • A subset of suspensions
  • Particles or molecules too small to be seen with ordinary microscopes (between 1 and 100 nm) become suspended in the solvent; scatters light, cannot be filtered
  • Under the influence of gravity, colloidal particles may take months, years, or even centuries to settle out.
  • Also considered heterogeneous.
slide20

Emulsions

  • Colloids whose particles are in the liquid phase.
slide22

Logistical questions

  • When we make and use solutions, there are some logistical questions that need to be answered:
    • Will the solute dissolve in the solvent?
    • How do you measure the concentration of a solution? How do you calculate the amount of solute and solvent to combine?
    • How do you change the concentration of a solution?
    • How do you make something dissolve faster?
    • How can you get more of a solute to dissolve in the same amount of solvent?
slide23

How do you know if something WILL dissolve?

    • Watch the video again
    • Salt dissolves in water because they are both polar. If water were not polar, or if salt were not polar (ionic), the water molecules would not be attracted to the salt ions.
  • This is true of all substances; polar things dissolve in polar substances, and nonpolar things dissolve in nonpolar substances.
  • Like dissolves like.
slide24

Like dissolves like

  • Will these dissolve?
    • CuCl2 in water
    • NaCl in CBr4
  • Why are water and oil immiscible?
    • Water is polar, oil is nonpolar.
  • Why does soap dissolve in water?
    • Lipids form micelles, surrounding nonpolar items with polar ends pointing outwards, allowing them to be washed away.

Yes! Both are polar.

No! Salt is polar, CBr4 is not.

slide25

Homework

  • Work on your research project papers; have the introduction and methods sections drafted by end of the month. We will peer-review around then.
slide26

Logistical questions

  • When we make and use solutions, there are some logistical questions that need to be answered:
    • Will the solute dissolve in the solvent?
    • How do you measure the concentration of a solution? How do you calculate the amount of solute and solvent to combine?
    • How do you change the concentration of a solution?
    • How do you make something dissolve faster?
    • How can you get more of a solute to dissolve in the same amount of solvent?
slide27

How do we measure solution concentration?

  • What exactly IS concentration?
    • Concentration is measured as the amount of solute per amount of solvent or solution.
    • For a liquid solution, we most commonly measure molarity, but there are many other units we can use.
slide28

Comparing concentrations

  • Concentrated: A LOT of solute is dissolved in the solution.
  • Dilute: A little solute is dissolved in the solution.
    • Note: these terms are very ambiguous! They are usually used in the relative sense – one solution is more dilute or concentrated than another.
slide30

Molarity

  • Molarity is the most common way to measure the concentration of a solution.
  • Molarity is expressed in M, which is equal to the moles of solute per liters of total solution.
    • A “3 molar (or 3 M) solution” of HCl is a solution that contains three moles of HCl per liter.
slide31

Molality

  • Molality (m) is calculated as moles of solute per kilograms of solvent.
    • A solution containing 3 moles of solute per 1 kg of solvent is a “3 molal solution”.
slide32

Mole Fraction

  • The mole fraction of a substance is measured as the moles of solute divided by the total moles of all parts of the solution. Note, there is no unit for this quantity.
slide33

Mass Percent

  • The mass percent is measured as mass of the solute divided by the total mass of the solution, multiplied by 100.
    • Note that since the units will cancel out, it does not matter what unit of mass you use – as long as they are the same.
slide34

What if you need to change the concentration?

  • Many solutions (like acids) are sold and shipped at high concentration for efficiency, and need to be diluted by the buyer to lower concentrations for lab work.
    • Let’s say you start with 50. mL of 16 M HCl. How many moles of HCl are in this sample?

16 M HCl = x moles HCl x = 0.050 L x 16 M

0.050 L = 0.80 moles HCl

slide35

What if you need to change the concentration?

  • If you were to use the same number of moles of HCl to make a 1.0 M solution, what would be its volume?

1.0 M = 0.80 moles HCl

x L solution

x = 0.80 L = 800 mL (Sig figs: 8.0 x 102mL)

slide36

Dilution Calculations

  • The calculation for dilutions is very simple. Since you are not changing the number of moles of the solute, the product of the volume and molarity of the initial solution is equal to the product of the volume and molarity of the diluted solution.
  • Notice that the volumes can be in L or mL – as long as they match.
slide37

Warm up

  • Devise a procedure for determining the molarity and molality of an unknown sodium chloride solution.
    • Start with a known volume (in L) of the solution
    • Measure the mass of the solution
    • Boil the solution until only the salt is left
    • Mass the salt alone and convert to moles
slide38

Warm up

  • Devise a procedure for determining the molarity and molality of an unknown sodium chloride solution.
    • Molarity: divide the moles of salt by the volume of the solution
    • Molality: subtract the final mass of the salt from the initial mass of the solution to find the mass of water lost. Convert this into kg. Divide the moles of salt by the kg of water lost.
slide39

Do it!

  • How would you use your data to calculate the mass % and mole fraction?
    • Mass %: Divide the mass of the salt by the original mass of the solution, multiply by 100%
    • Mole fraction: convert mass of salt and mass of water lost into moles and divide moles of salt by total moles
slide40

Methods of dilution

  • So you would need to take your 50. mLHCl and add enough water to make 800 mL of the diluted solution.
  • Most people assume that this means you can measure out 750 mL of water and just add it to the 50 mLHCl. That is not always true because molecules do not always pack the same way.
slide41

Methods of dilution

  • The most common method of making a dilution requires a volumetric flask.
slide42

Warm up

  • Turn in your calculations from Tuesday’s mini-lab. I will award 2 EC pts to the group with the fewest errors and 1 EC pt to the runners-up.
  • List the ways in which you can make a solute dissolve in the solvent faster (hint: there are 4).
slide43

Logistical questions

  • When we make and use solutions, there are some logistical questions that need to be answered. We have answered:
    • Will the solute dissolve in the solvent?
    • How do you measure the concentration of a solution?
    • How do you calculate the amount of solute and solvent to combine?
    • How do you change the concentration of a solution?
slide44

Logistical questions

  • We still need to know:
    • How do you make something dissolve faster?
    • How can you get more of a solute to dissolve in the same amount of solvent?
    • How do you describe concentrations relative to the solute’s identity?
slide45

What affects how FAST something dissolves?

  • Watch the video again – what has to happen in order for the salt to dissolve?
  • Just like a chemical reaction, forming a solution is all about particle collisions! The more contact there is between the solute and solvent particles, the faster something will dissolve.
    • How do you get particles to hit each other more often?
slide46

Factors that affect the rate of solution formation

  • Heat
    • Note: Gases work differently
  • Surface area
  • Stirring (agitation)
  • Concentration
slide47

How about how MUCH can dissolve?

  • The solubility of a substance is an important property – it is defined as the extent to which a substance can be dissolved in a given solvent (how MUCH of it you can get to dissolve in a certain amount of solvent?) at a specific temp.
  • What affects the solubility of a substance?
slide48

Factors that affect solubility

  • Nature of solute and solvent (remember: like dissolves like)
  • Temperature (different for gases)
  • Pressure (only for gases)
slide49

Why do gases work differently?

  • When heating a solid or a liquid, what happens?
    • Particles move faster and farther apart; particles eventually move far enough apart that we observe a phase change – usually solid to liquid
    • The phase change itself requires extra energy
slide50

Heating Curves

  • If you heat ice and graph the temperature of the substance vs. time, the graph looks like this:
slide51

Explanation

  • 1st section on the left: the solid ice is heating up. The water molecules are vibrating harder and faster, until they hit 0 degrees Celsius.
  • 2nd section: energy is still being added, but the temperature doesn’t change! The energy is all being used for the solid-to-liquid phase change.
  • 3rd section: the water is all in liquid form now, and the particles speed up, increasing in temp again.
slide52

Explanation

  • 4th section: we see the temperature again stop changing, even though we are still adding energy. That energy is all being used for the liquid-to-gas phase change.
  • 5th section: the water has all boiled off and become steam at this point, so the temperature can once again increase as the molecules speed up.
slide53

Phase diagrams

  • The states of matter/changes thereof can also be represented by a phase diagram. Here, the pressure is also considered as a factorthat affects the phase(closed system).
  • Beyond the critical pt, the substance is alwaysa gas
slide54

Why do gases work differently?

  • Solids and liquids cannot escape the solution without undergoing a phase change to gas. Generally, the solvent will evaporate at a lower temperature before the solute gains enough energy to do this.
  • In the case of a gaseous solute, no phase change is needed. If the particles gain enough kinetic energy, they will simply escape.
slide55

Cooling Curves

  • Next class, we will construct the cooling curve for phenyl salicylate. You will see the same features (a cooling period, followed by a plateau during the phase change, followed by more cooling), in reverse.
slide56

How do we describe the amount dissolved compared to the amount you CAN dissolve?

    • What does it mean for a towel to be saturated?
    • What does it mean for a color to be more saturated?
  • Saturation indicates fullness
    • For example, a towel so wet that it cannot hold any more water is saturated, a color with a high concentration of pigment is a saturated color, etc.
slide57

Imagine that

    • When a small amount of solute is added to a solvent, molecules of solute rapidly dissociate and become solvated by the solvent molecules. Generally, more solutecan still be dissolved.
  • At this point, you have anunsaturated solution – less solute is dissolved than the solvent is capable of solvating.
slide58

Imagine that

  • Keep adding solute. Eventually, the solvent molecules are all occupied and no more solute can be dissolved. You now have a saturated solution – one that has as much solute dissolved in it as possible.
  • Think of it like a hotel
    • When there are empty rooms still available, the hotel is “unsaturated”
    • When all the rooms are booked, the hotel is “saturated” with guests
slide59

The song that never ends

  • If no more solute can be dissolved, does that mean that the process of solvation stops?
  • Again, think of it like a hotel
    • Hotels never stop booking – they just book guests for later dates when they run out of rooms. In order for new guests to come in, prior guests have to leave.
slide60

Solution equilibrium

  • In a saturated solution, the solute is constantly going back and forth between dissolving and precipitating. The rate of solvation = rate of precipitation.
  • This is known as solution equilibrium.
slide61

Under special conditions

  • If a major catastrophe happens at or near a hospital and there are a lot of casualties, what might the hospital do?
  • A solution can be made to be supersaturated – that is, it can be made to have more solute dissolved in it than is theoretically possible under normal conditions.
    • This often requires special preparation conditions and/or techniques. You may see this during the cooling curves lab.
solubility curves

Solubility Curves

Reading and Interpreting

slide63

Graphs in general

  • Remember the good graphing practices
    • Graphs should have a title and axes with detailed labels (including units) – always look here first to see what the graph is about
  • Generally the only feature of a graph that we are interested in is the line with the data on it.
  • But the empty space on a graph has a meaning! It’s just not usually something we are interested in.
slide64

Solubility

  • Remember, solubility is defined as how much of a solute you can dissolve in a given amount of solvent. It changes depending on the temperature of the system.
  • Water, being the universal solvent, is the most common solvent used in this measurement.
  • So more specifically: solubility is commonly defined as the total grams of solute that can be dissolved in 100 g of water.
slide65

Solubility curves

  • A solubility curveis a graph that shows the maximum amount of solute that can be dissolved in 100 g of water at any given temp.
    • The curve tells you when a solution is saturated!
  • Many solutes can be shown on the same graph.
slide66

For each curve on the graph

  • On the line = saturated
  • Under the line = unsaturated
  • Above line = supersaturated
slide67

Solubility curves

  • At 40 degrees Celsius, how many grams of KCl can dissolve in 100 g of waterunder normal conditions?
  • At 80 degrees Celsius, howmany grams of sodiumnitrate can dissolve in 300 gof water?

39 g KCl

~441 g NaNO3

slide68

Solubility curves

  • At 50 degrees C, how many grams of sulfur dioxide willmake a saturated solution in100 g of water?
  • At 20 degrees C, which solute is the most soluble?Which is the least soluble?

4 g SO2

Most = KI; Least = KClO3

slide69

Solubility curves

  • If I dissolve 75 g potassiumnitrate in 100 g of water at60 degrees C, what type of solution have I made?
  • If you dissolve 90 g of ammonium chloride in 100 g of water at 80 degrees C, what type of solution would you have?

Unsaturated!

Supersaturated!

slide70

Homework

  • Worksheet on solubility curves
  • Dress for lab
  • Continue writing research project papers; final drafts due May 22nd/23rd
  • Monday: Finish chapter 15, Cooling curves lab
  • Wednesday: review and test (intro chapter 18?)

70

slide72

What is a colligative property?

  • Acolligative property is a property of a solution that changes based on the number of solute particles dissolved in the solvent, and not on the chemical properties or identities of those particles.This is an extensive property!
slide73

How do you get that number?

  • The number of dissociated particles produced by a solute depends on the type of bonding in the solute.
  • Covalent solutes do not break apart
  • Ionic solutes break up into their respective ions (polyatomic ions do not break up!)
    • How many particles would be produced by dissolving 1 molecule of NaCl? Mg3(PO4) 2? Glucose? NH3?
  • This is known as the van’t Hoff factor (i, unitless).
slide74

JacobusHenricusvan’t Hoff II

  • Dutch physical and organic chemist
  • First winner of the Nobel Prizein chemistry (1901)
  • His work in chemical kinetics, equilibrium, osmotic pressure,and stereochemistry helped tofound the field of physical chemistry
slide75

What are the colligative properties?

  • Vapor Pressure
  • Freezing Point
  • Boiling Point
  • Osmotic Pressure
slide76

Vapor Pressure

  • The pressure exerted by the vapor in dynamic equilibrium with its liquid in a closed system.
  • Does the solvent’s vapor pressure increase or decrease when you dissolve a solute in it?
slide77

Vapor Pressure Lowering

  • Dissolving a solute in a solvent lowers the solvent’s vapor pressure.
  • Solvent molecules are attracted to solute particles, forming solvation “shells”. This keeps them from escaping into the gas phase.
slide78

Freezing Point

  • The freezing point of a solution is the temperature at which the solution makes the phase change from liquid to solid.
  • Does the solvent’s freezing point increase or decrease when you dissolve a solute in it?
slide79

Freezing Point Depression

  • When a substance freezes, the particles have to move into a highly ordered structure to form a solid.
  • If the solvent is attracted to the solute (making solvation shells), they have a harder time forming the solid structure.Therefore, you have to lower the tempfurther than usual to get the solution tofreeze.
slide80

Freezing Point Depression

  • The temperature difference between the freezing point of the pure solvent and the freezing point of the solution is known as the “freezing point depression”, and is given the symbol .
    • i is the van’t Hoff factor
    • Kf is the cryoscopic constant, different for each solvent
slide81

Real Life: Salting Icy Roads

  • In the winter, what do we often do to combat icy roads and sidewalks?
  • Why would magnesium chloride be better than sodium chloride for melting ice?
  • Would an ocean freeze over at a higher or lower temperature than a lake?
slide82

Boiling Point

  • The temperature at which a solution makes the phase change from liquid to gas.
  • The boiling point is also defined as the temperature at which the vapor pressure of the liquid equals atmospheric pressure.
  • Would the solvent’s boiling point increase or decrease if you dissolve a solute in it?
slide83

Boiling Point Elevation

  • Remember, adding a solute causes the vapor pressure of the solution to be lower than the vapor pressure of the pure solvent.
  • In order to make the vapor pressure of the solution equal to the atmospheric pressure, more kinetic energy needs to be added. More kinetic energy means higher temperature!
slide84

Real Life: Cooking pasta, antifreeze

  • When you cook pasta, what do you add to the boiling water?
  • Antifreeze is a commonly used solution that keeps your car engine from freezing up. What purpose does it serve in the summer?
slide85

Real Life: Hot springs

  • At a hot spring, the temperature of the spring water is often reported at >100 degrees, but the steamy water is completely still. How could this be?
  • At the same temperature, would a lakeside resort or an oceanside resort experience more humidity?
slide86

Boiling Point Elevation

  • The temperature difference between the boiling point of the pure solvent and the boiling point of the solution is known as the “boiling point elevation”, and is given the symbol .
    • i is still the van’t Hoff factor.
    • Kb is the ebullioscopic constant, different for each solvent.
slide87

Sample Calculations

  • If 20 g of sodium chloride is dissolved in 1 L of water, what will the freezing point of the solution be? The new boiling point?
  • If 20 g of magnesium chloride is dissolved in 1 L of water, what will the freezing point of the solution be? The new boiling point?

For water:

Kf = -1.86 K·kg/mol

Kb = 0.512 K·kg/mol