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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