Unit 10: Solutions Chapter 15-16

This tutorial is designed to help students understand scientific measurements.

Objectives for this unit appear on the next slide.

◦ Each objective is linked to its description.

◦ Select the number at the front of the slide to go directly to its description.

Throughout the tutorial, key words will be defined.

◦ Select the word to see its definition.

Objectives

18 Define solubility including the terms soluble/insoluble and miscible/immiscible; unsaturated, saturated, and supersaturated; solute, solvent, and aqueous

19 Define conductivity including how electrolytes and dissociation affect conductivity

20 Describe heterogeneous mixtures including suspensions, colloids, and emulsions

21 Explain how mixtures can be separated by six methods

22 Explain saturation points regarding liquids and gases and read a solubility curve chart

23 State Henry’s Law and give examples of the law

24 Define concentration in terms of molarity, molality, and solve problems

25 Solve dilution problems involving molarity

26 Identify and define colligative properties including calculations for boiling point elevation and freezing point depression

18 Defining a Solution

Two chemicals can often be mixed together to form a solution. ◦ In this case, the chemical that has the larger

quantity is known as a solvent. If water is the solvent, the solution is known as aqueous.

◦ The chemical that has the smaller quantity is known as a solute.

◦ It is stated that a solute will dissolve in the solvent This is not limited to a solid dissolved into a liquid.

Gases can be dissolved in liquids (pop).

Liquids can be dissolved in liquids (gasoline).

Gases can be dissolved in gases (air in the room).

Etc.

Defining a solution

When identifying solutes, they are often

labeled as:

◦ Soluble: will dissolve in a certain solvent

◦ Insoluble: will not dissolve in a certain solvent

If it is a liquid in a liquid, the solutes are

labeled as:

◦ Miscible: will dissolve in a certain liquid

◦ Immiscible: will not dissolve in a certain liquid

Defining a solution

When a solution is made, the solute

dissolves in the solvent.

The amount of solute can vary.

◦ Solvents can only hold so much solute.

◦ Once too much is added, the remaining solute sinks to the bottom.

The terms saturated, unsaturated, and

supersaturated explain this situation.

Unsaturated

A solution in which more solute can be

added.

Saturated

A solution in which the solution is holding

as much solute as possible with any extra

settling to the bottom.

Supersaturated

A unique case where a solution has more

solute dissolved than in should be allowed

to.

19 Conductivity

The conductivity of a solution is

dependent on the presence of

electrolytes.

◦ An electrolyte is an ion produced when a solid dissociates into a solvent.

◦ If electrolytes are present, the solution will conduct electricity.

Dissociation

Ions are found in solution when a

compound dissociates during the

dissolving process.

Each compound will break into a distinct

number of parts.

◦ Covalent molecules remain as one.

◦ Ionic molecules break into their ions.

Dissociation

When an ionic compound is placed in

water, the partially positive hydrogens of

the water molecule will attract the anion.

The partially negative oxygen of the water

molecule will attract the cation.

These attractions will pull apart the

compound thus separating it into ions.

Dissociation

Dissociation Factors

Each compound will break into certain

numbers of parts when they dissociate.

◦ This is the dissociation factor.

For covalent molecules, the molecule will

not break apart so the dissociation factor

is always 1.

For ionic compounds, the compound will

break into its ions so it is determined by

adding the ions together.

Dissociation Factors

Consider NaCl

◦ There is one Na+ and one Cl-.

◦ Therefore, the dissociation factor is 2.

Here are some more examples:

Molecule/

Compound

Dissociation

Factor

Reason

CH4 1 Covalent

CaCl2 3 Ionic

Ca(NO3)2 3 Ionic (polyatomic ions

remain together)

20 Heterogeneous Mixtures

Not all mixtures will make solutions.

Some are not uniform and are called

heterogeneous mixtures.

Two common varieties of heterogeneous

mixtures are:

◦ Suspensions: will separate if not agitated

Surfactants and micelles will make up suspensions

These describe soaps and detergents.

◦ Colloids: will not separate

21 Separations

With the different homogenous (solutions)

and heterogeneous mixtures that exist, it is

important to be able to separate the parts.

There are six main ways to separate

mixtures:

◦ Decanting

◦ Filtration

◦ Evaporation

◦ Centrifuge

◦ Distillation

◦ Chromatography

Decanting

Decanting is a

separation technique

where the solid is

allowed to settle to the

bottom.

The liquid is then

poured off leaving two

parts.

Image from: http://www.sciencequiz.net/jcscience/jcchemistry/practicals/decanting.htm

Filtration

Filtration occurs by pouring a mixture through a porous paper.

The larger particles are caught in/above the paper.

The smaller particles pass through to the collection container.

Image from: http://en.wikipedia.org/wiki/File:Vacuum-filtration-diagram.png

//upload.wikimedia.org/wikipedia/commons/d/da/Vacuum-filtration-diagram.png

Evaporation

Evaporation is the

removal of the liquid

from a solution.

The liquid is heated

past its boiling point

driving it away.

The solute then sinks

to the bottom.

Image from: http://www.school-for-champions.com/science/evaporation.htm

Centrifuge

The centrifuge is an instrument in which the mixture is separated based on its densities.

The sample is spun rapidly forcing the more dense portion to the bottom of the tube.

After using a centrifuge, the sample is generally decanted.

Image from: http://www.daviddarling.info/encyclopedia/C/centrifuge.html

Distillation

Distillation is used when two or more liquids are present with distinct boiling points.

The mixture is heated to just above the boiling point of one liquid.

As it boils, the gas is funneled down a cooling tube causing the gas to condense.

It is then collected in an additional flask.

This allows both liquids to be kept.

Image from: http://glossary.periodni.com/glossary.php?en=distillation

Chromatography

Chromatography is a widely used technique to separate mixtures.

It is typically used for dyes and organic molecules.

The sample is added to a mobile phase which is passed over a stationary phase.

The stationary phase will interact with one part of the mixture thus slowing it down.

As this occurs, the mobile phase pulls the other part of the mixture further down the column.

Image from: http://www.waters.com/waters/nav.htm?cid=10048919&locale=en_US

Chromatography

There are several types of

chromatography.

Here are a few:

◦ Gas chromatography

◦ High Pressure Liquid Chromatography

◦ Ion Chromatography

◦ Size-Exclusion Chromatography

◦ Thin-Layer Chromatography

Separation Recap

Technique

Separates

Heterogeneous

Mixtures

Separates

Homogenous

Mixtures

Notes

Decanting Yes No

Filtration Yes No Uses porous paper;

requires different sizes of

particles

Evaporation Yes Yes Will lose the liquid

portion of sample

Centrifuge Yes No

Requires the use of

another technique

(typically decanting to

finish the separation)

Distillation Yes Yes Requires distinct boiling

points; keeps both liquids

Chromatography Not Preferred Yes Mostly used with dyes

and organics

22 Solubility

Solubility describes how much solute a

solution is allowed to hold during certain

conditions.

The solubility of certain substances can

be shown using a solubility curve.

◦ The curve shows how a solute dissolves in a solvent given certain conditions.

Solubility Curves

The line represents the saturation point at each temperature. ◦ For example: at 20°C, sugar is

saturated when 200 grams are added to 100 grams of water.

◦ Any point below the line is unsaturated.

◦ Any point above the line is saturated.

The curves often show more than one solute for each solvent. ◦ In this case, water is the solvent.

Image taken from: http://www.btinternet.com/~chemistry.diagrams/solubility_curves.htm

http://www.btinternet.com/~chemistry.diagrams/solubility_curves.htm

23 Henry’s Law

Henry’s Law explains a phenomena that

occurs when a gas is dissolved in a liquid.

The gas in the liquid is directly

proportional to the partial pressure of

the gas above that liquid.

If the gas above the liquid is removed, gas

that is dissolved will escape to fill this

space.

Henry’s Law Example - POP

One of the key components of pop is the carbonation.

The carbonation is caused by dissolving carbon dioxide in the liquid.

When the lid is removed from a bottle of pop, the gas above the liquid is removed.

◦ This will cause some gas in solution to escape and fill the space above the liquid meaning less carbon dioxide is dissolved.

◦ After time, this will cause the pop to go flat.

Henry’s Law Example-POP

The pop starts

with some gas

dissolved and

some in the space

above the liquid.

Once the can is

opened, gas escapes

out the top of the can.

To return to the

correct relationship,

some gas will leave

the solution to fill

the space again.

24 Concentration

Since solutions can have more or less

solute, it is important to know how much

solute is dissolved.

This is measured with concentration.

Concentration is typically given in either

molarity or molality.

Molarity

Molarity provides a relationship between

the moles of solute and liters of solution.

Its mathematical relationship is:

Molarity = 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒

𝐿𝑖𝑡𝑒𝑟𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

For example:

Assume each red dot represents one mole.

Assume there is 500 ml of water.

The molarity of this solution would be:

Molality

Molarity provides a relationship between

the moles of solute and kilograms of

solvent.

Its mathematical relationship is:

Molality = 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒

𝑘𝑖𝑙𝑜𝑔𝑟𝑎𝑚𝑠 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡

For example:

Assume each red dot represents one mole.

Assume there is 450 grams of water.

The molality of this solution would be:

Solving concentration problems

Assume you 200 ml have a 5 molar solution and you want to know how many moles this would be.

◦ 5 molar can be written as 5 𝑚𝑜𝑙𝑒𝑠

1 𝑙𝑖𝑡𝑒𝑟

◦ 200 ml would have to be converted to liters

So 200 ml x 1 𝑙𝑖𝑡𝑒𝑟

1000 𝑚𝑙= 0.200 liters

◦ Dimensional analysis will allow moles to be calculated.

0.200 liters x 5 𝑚𝑜𝑙𝑒𝑠

1 𝑙𝑖𝑡𝑒𝑟= 1 mole

25 Dilutions

Often, it is necessary to take a

concentrated solution and make it less

concentrated.

◦ This is called a dilution.

To make a dilution, a portion of the

concentrated amount is taken and more

solvent is added.

Dilutions

Suppose a 2 molar solution was desired but only a 4 molar solution was present.

◦ Assume that water was the solvent.

To make the 2 molar solution, take a sample of the concentrated and double the water.

4 molar solution 2 molar solution

Dilutions

The previous problem can be calculated mathematically as well.

Assume a 500 ml sample of a 2 molar solution was desired.

If this was the case, we would need 1 mole of solute.

◦ 0.500 liters x 2 𝑚𝑜𝑙𝑒𝑠

1 𝑙𝑖𝑡𝑒𝑟= 1.0 moles

To get 1 mole of solute from the 4 molar solution, we would need 250 ml.

◦ 1.0 moles x 1 𝑙𝑖𝑡𝑒𝑟

4 𝑚𝑜𝑙𝑒𝑠= 0.25 𝑙𝑖𝑡𝑒𝑟𝑠

Once we take the 250 ml, we would add an additional 250 ml of water to bring the solution to 500 ml.

Dilutions

The calculation used on the previous slide

can be condensed to the following:

M1V1 = M2V2

It is important to realize that the volume

of the concentrated is only the volume

needed. Enough solvent would have to be

added to reach the desired concentration.

Desired concentration

Desired volume

Volume of

concentrated required

Molarity of concentrated sample

26 Colligative Properties

Colligative properties are physical

properties that are affected by the

amount of solute rather than the actual

identity of that solute.

Two of the more common properties that

are affected by the solute are:

◦ Freezing points

◦ Boiling points

Freezing Point Depression

The freezing point of the solvent will decrease based on the amount of solvent present.

The change in the freezing point is calculated using the equation:

∆T = kfdm

∆T = change in temperature

Kf = freezing constant (find on the PT)

d = dissociation factor

m = molality

Boiling Point Elevation

The boiling point of the solvent will increase based on the amount of solvent present.

The change in the boiling point is calculated using the equation:

∆T = kbdm

∆T = change in temperature

Kb = boiling constant (look on back of PT)

d = dissociation factor

m = molality

Colligative Properties

Recall that water boils at 100°C and freezes at 0°C.

If we dissolve NaCl into water (assume 2 m) the freezing a boiling point will change.

Freezing Boiling ∆T = kfdm ∆T = kbdm

∆T = 1.86°C/m x 2 x 2 m ∆T = 0.52°C/m x 2 x 2 m

∆T = 7.44 °C ∆T = 2.08°C

Therefore, the new freezing point is -7.44 °C and the new boiling point is 102.08 °C.

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