• Define the terms heterogeneous and homogeneous.

• What do you think is the difference between dissolving and melting?

Warm up

1

• 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

Warm up

2

• Melting: one substance changing phase from solid to liquid

• Dissolving: two substances mixing to form a solution

Warm up

3

SolutionsChapter 15 (write the red!)

• 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.

Objectives

5

• 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.

Defining solutions

HETEROGENEOUS.

6

• 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.

Mixtures

7

inverse

• A solution is a homogeneous mixture – one substance dissolved 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.

The terminology of solutions

8

• 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!

The terminology of solutions

9

• 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)

The terminology of solutions

10

• 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.

The terminology of solutions

11

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

The terminology of solutions

12

• 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.

The terminology of solutions

13

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

• Opposite: immiscible.

The terminology of solutions

14

Solute Solvent Example

Solid Solid

Liquid

Gas

Liquid Solid

Liquid

Gas

Gas Solid

Liquid

Gas

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

15

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

Careful!

16

Comparison

17

Property Solution Colloid/Emulsion

Suspension

Particle Type

Effect of Light

Effect of Gravity

Filtration

Uniformity

Examples

Atoms, ions, small molecules

No scatteringStable, no separation

Particles not retained on filters

HomogeneousKool Aid

• 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 by

filtration• Considered

heterogeneous

Suspensions

18

• 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.

Colloids

19

• Colloids whose particles are in the liquid phase.

Emulsions

20

Comparison

21

Property Solution Colloid/Emulsion

Suspension

Particle Type Atoms, ions, small molecules

Large molecules (1 to 100 nm)

Fine particles (>100 nm)

Effect of Light

No scattering

Scattered Scattered

Effect of Gravity

Stable, no separation

Separates over long periods of time

Unstable, forms sediment

Filtration Particles not retained on filters

Particles not retained on filters

Particles retained on filters

Uniformity Homogeneous

Borderline Heterogeneous

Examples Kool Aid Whipped cream/milk

Silt water

• 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?

Logistical questions

22

• 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.

How do you know if something WILL dissolve?

23

• 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.

Like dissolves like

24

Yes! Both are polar.No! Salt is polar, CBr4 is not.

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

Homework

25

• 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?

Logistical questions

26

• 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.

How do we measure solution concentration?

27

• 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.

Comparing concentrations

28

29

• 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.

Molarity

30

• 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”.

Molality

31

• 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.

Mole Fraction

32

• 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.

Mass Percent

33

• 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?

What if you need to change the concentration?

34

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

0.050 L = 0.80 moles HCl

• 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 102 mL)

What if you need to change the concentration?

35

• 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.

Dilution Calculations

36

• 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

Warm up

37

• 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.

Warm up

38

• 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

Do it!

39

• So you would need to take your 50. mL HCl 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 mL HCl. That is not always true because molecules do not always pack the same way.

Methods of dilution

40

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

Methods of dilution

41

• 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).

Warm up

42

• 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?

Logistical questions

43

• 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?

Logistical questions

44

• 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?

What affects how FAST something dissolves?

45

• Heat– Note: Gases work differently

• Surface area• Stirring (agitation)• Concentration

Factors that affect the rate of solution formation

46

• 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?

How about how MUCH can dissolve?

47

• Nature of solute and solvent (remember: like dissolves like)

• Temperature (different for gases)• Pressure (only for gases)

Factors that affect solubility

48

• 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

Why do gases work differently?

49

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

Heating Curves

50

• 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.

Explanation

51

• 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.

Explanation

52

• 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

Phase diagrams

53

• 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.

Why do gases work differently?

54

• 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.

Cooling Curves

55

• 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.

How do we describe the amount dissolved compared to the amount

you CAN dissolve?

56

• 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 an unsaturated solution – less solute is dissolved than the solvent is capable of solvating.

Imagine that

57

• 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

Imagine that

58

• 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.

The song that never ends

59

• 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.

Solution equilibrium

60

• 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.

Under special conditions

61

Solubility CurvesReading and Interpreting

• 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.

Graphs in general

63

• 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.

Solubility

64

• A solubility curve is 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.

Solubility curves

65

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

For each curve on the graph

66

• 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?

Solubility curves

67

39 g KCl

~441 g NaNO3

• 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?

Solubility curves

68

4 g SO2

Most = KI; Least = KClO3

• 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?

Solubility curves

69

Unsaturated!

Supersaturated!

• 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?)

Homework

70

Colligative Properties

• A colligative 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!

What is a colligative property?

72

• 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).

How do you get that number?

73

• 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

Jacobus Henricus van’t Hoff II

74

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

What are the colligative properties?

75

• 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?

Vapor Pressure

76

• 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.

Vapor Pressure Lowering

77

• 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?

Freezing Point

78

• 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.

Freezing Point Depression

79

• 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

Freezing Point Depression

80

• 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?

Real Life: Salting Icy Roads

81

• 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?

Boiling Point

82

• 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!

Boiling Point Elevation

83

• 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?

Real Life: Cooking pasta, antifreeze

84

• 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?

Real Life: Hot springs

85

• 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.

Boiling Point Elevation

86

• 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

Sample Calculations

87