Transcript Slide 1

Chapter 18: Solutions
18.1 PROPERTIES OF SOLUTIONS
•Solution Formation – Recall that Solutions are
homogenous (physical) mixtures
•Solute dissolved particles in a solution (lesser
amount of the two parts) are dissolved in the
Solvent (the greater amount of the two parts)
forming a solution.
• How well a material (solute) dissolves in a
solvent is described by a substance’s Solubility.
• There are three factors that determine how fast a
substance dissolves are: stirring (agitation),
temperature and surface area.
•Think about ways you
have used these factors to
better dissolve your sugar
into your tea.
Temperature effects Solubility
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Temperature effects Solubility
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Salt Solubility Simulation by PhET & Colorado University
Water: The Super Solvent
•Water is an excellent solvent. It is used to
dissolve most of the substances you have
encountered in your life. In fact it is the
solubility of the water in your body, blood and
cells that accounts for most of the chemical
processes that allow you to stay alive.
• Most of the water on
Earth is not pure, but
rather is present in
solutions.
Water: The Super Solvent
• Water is such a versatile solvent that it is
sometimes called the universal solvent.
• Water is difficult to keep pure because it is an
excellent solvent for a variety of solutes.
• Its ability to act as a solvent is one of its
most important physical properties.
• As you will see, it is again the attraction of
water molecules for other molecules, as
well as for one another, that accounts for
these solvent properties.
Water Dissolves Many Ionic Substances
• Salt, like a great many ionic compounds,
is soluble in water. The salt solution is
also an excellent conductor of electricity.
• This high level of electrical conductivity is
always observed when ionic compounds
dissolve to a significant extent in water.
• Your model of water and its interactions
explains why salt and many other ionic
compounds dissolve in water and why the
solutions conduct electricity.
A Model of the Dissolving of NaCl
• Remember that ionic solids are composed of a
three-dimensional network of positive and
negative ions, which form strong ionic bonds.
A Model of the Dissolving of NaCl
• The process by which the charged particles
in an ionic solid separate from one another
is called dissociation.
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A Model of the Dissolving of NaCl
• You can represent the process of dissolving
and dissociation in shorthand fashion by the
following equation.
Water Dissolves Many Covalent Substances
• Water is not only good at dissolving ionic
substances. It also is a good solvent for
many covalent compounds.
• Consider the covalent substance sucrose,
commonly known as table sugar, as an
example.
Water Dissolves Many Covalent Substances
• You have probably observed that this
substance, with the formula C12H22O11,
dissolves in water. In fact, it is highly
soluble.
• It’s possible to dissolve almost 200 g of
sugar in 100 mL of water.
Water Dissolves Many Covalent Substances
• Take a look at the molecular structure of a
sucrose molecule.
• Notice that the
structure has a
number of O—H
bonds.
Water Dissolves Many Covalent Substances
• As you learned earlier, if a molecule contains
O—H bonds, it will
tend to be polar and
it can form
hydrogen bonds.
Water: The Molecular View
• Because water is so much a part of life, its
properties are easy to take for granted.
• If you step back a bit and examine water
scientifically, you will find that it is unusual
among the compounds found on Earth.
• Water is most often thought of as a liquid.
Water: The Molecular View
• However, solid water, called ice, and gaseous
water, called steam or water vapor, also exist
in large quantities on Earth.
• Water is the only substance on Earth that
exists in large quantities in all three states.
Geometry of the Water Molecule
• The arrangement of electrons about the central
oxygen in the water molecule relates to its
three-dimensional geometry.
• There is a large
electronegativity
difference between the
covalently bonded
hydrogen and oxygen.
Geometry of the Water Molecule
• Therefore, the electron pair is shared unequally.
• Because of the molecule’s bent structure, the
poles of positive and
negative charge in the
two bonds do not
cancel, and the water
molecule as a whole is
polar.
Like Dissolves Like
• Although water dissolves an enormous
variety of substances, both ionic and
covalent, it does not dissolve everything.
• The phrase that scientists often use when
predicting solubility is “like dissolves like.”
• The expression means that dissolving
occurs when similarities exist between the
solvent and the solute.
Concentrated Versus Dilute
• Chemists never apply the terms strong and
weak to solution concentrations.
• As you’ll see in
the next chapter,
these terms are
used in chemistry
to describe the
chemical behavior
of acids and bases.
• Instead, use the terms concentrated and dilute.
Unsaturated Versus Saturated
• Another way of providing information about
solution composition is to express how much
solute is present relative to the maximum
amount the solution could hold.
• If the amount of solute dissolved is less than
the maximum that could be dissolved, the
solution is called an unsaturated solution.
Unsaturated Versus Saturated
• Such a solution, which holds the maximum
amount of solute per amount of the solution
under the given conditions, is called a
saturated solution.
Unsaturated Versus Saturated
• An interesting third category of solution is
called a supersaturated solution.
• Such solutions contain more solute than the
usual maximum amount and are unstable.
• They cannot permanently hold the excess
solute in solution and may release it suddenly.
Unsaturated Versus Saturated
• Supersaturated solutions, as you might
imagine, have to be prepared carefully.
• Generally, this is done by dissolving a
solute in the solution at an elevated
temperature, at which solubility is higher
than at room temperature, and then slowly
cooling the solution.
Effect of Temperature on Solubility
• Temperature has a significant effect on
solubility for most solutes.
• The solubilities of some solutes, such as
sodium nitrate and potassium nitrate,
increase dramatically with increasing
temperature.
Effect of Temperature on Solubility
• Other solutes, like NaCl and KCl, show
only slight increases in solubility with
increasing temperatures.
• A few solutes, like cerium(III) sulfate,
Ce2(SO4)3, decrease in solubility as
temperature increases.
Molarity
• Concentration units can vary greatly.
• They express a ratio that compares an
amount of the solute with an amount of the
solution or the solvent.
• For chemistry applications, the concentration
term molarity is generally the most useful.
• Molarity is defined as the number of moles
of solute per liter of solution.
• Molarity = moles of solute/liter of solution
Molarity
• Note that the volume is the total solution
volume that results, not the volume of
solvent alone.
• Suppose you
need 1.0 L of
the salt solution
mentioned
above.
Molarity
• In order to be at the same concentration as the
salt in the patient’s blood, it needs to have a
concentration of 0.15 moles of sodium
chloride per liter of solution.
• In other words, it
must have a
molarity of 0.15.
Molarity
• To save space, you refer to the solution as
0.15M NaCl, where the M stands for
“moles/liter” and represents the word molar.
• Thus, you need 1.0
L of a 0.15-molar
solution of NaCl.
How are you going
to prepare it?
Molarity
• Assuming you’re making an aqueous solution,
you need to know only three things when
working quantitatively: the concentration, the
amount of solute, and the total volume of
solution needed.
Preparing 1 L of an NaCl Solution
• How would you prepare 1.0 L of a 0.15M
sodium chloride solution?
• First, determine the mass of NaCl to add to a
1.0-L container.
• The 0.15M solution must contain 0.15 moles
of NaCl per liter of solution.
Preparing 1 L of an NaCl Solution
• The proper setup, showing the conversion
factors, is as follows.
Preparing 1 L of an NaCl Solution
• Then carry out cancellations and calculate
the answer.
Preparing 1 L of an NaCl Solution
• The result means you need to measure 8.8 g
of NaCl, add some water to dissolve it, and
then add enough additional water to bring
the total volume of the solution to 1.0 L.
Preparing a Different Volume of a Glucose
Solution
• How would you prepare 5.0L of a 1.5M
solution of glucose, C6H12O6?
• You need to determine the number of grams
of glucose to add to a 5.0-L container.
Preparing a Different Volume of a Glucose
Solution
• The 1.5M solution must contain 1.5 mol of
glucose per liter of solution.
• The proper setup, showing the conversion
factors, is as follows.
Preparing a Different Volume of a Glucose
Solution
• Cancel units and carry out the calculation.
Preparing a Different Volume of a Glucose
Solution
• The mass of glucose required is 1400 g.
• Weigh this mass, add it to a 5.0-L container,
add enough water to dissolve the glucose,
and fill with water to the 5.0-L mark.
Calculating Molarity
• You add 32.0 g of potassium chloride to a
container and add enough water to bring the
total solution volume to 955 mL. What is
the molarity of this solution?
• You are given that there are 32.0 g of solute
per 955 mL of solution, so this relationship
can be expressed in fraction form with the
volume in the denominator.
Calculating Molarity
• Therefore, the initial part of the setup is
as follows.
Calculating Molarity
• Determine that the molar mass of KCl is
74.6 g/mol by adding the atomic masses of
K and Cl and applying the unit grams/mole
to the sum.
• The conversion factor that must be used to
convert from grams to moles of KCl is
1mol KCl/74.6 g KCl.
Calculating Molarity
• Next, to convert milliliters to liters, given
that there are 1000 mL solution/L solution,
use that conversion factor in the setup.
Calculating Molarity
• Cancel units and carry out the calculation,
using the setup just developed.
Freezing-Point Depression
• A solution always has a lower freezing
point than the corresponding pure solvent.
• If you are interested only in aqueous
solutions, this means that any aqueous
solution will have a freezing point lower
than 0°C.
• The amount that the freezing point is
depressed relative to 0°C depends only
upon the concentration of the solute.
Boiling-Point Elevation
• You have just learned that the freezing point
of a solution is lower than the freezing point
of the pure solvent.
• It turns out that the boiling point of a
solution is higher than the boiling point
of the pure solvent.
Boiling-Point Elevation
• For aqueous solutions this means that the
solution boiling point will be greater than
100°C, assuming standard atmospheric
pressure.
• The solute must also be nonvolatile; that is
not able to evaporate readily.
Molality
• The molality (m) of a solution is equal to
the number of moles of solute per kilogram
of solvent.
Calculating Molality
• What is the molality of a solution that contains
16.3 g of potassium chloride dissolved in 845
g of water?
• Convert the mass of solute to moles.
Calculating Molality
• The solvent mass, 845 g, must be expressed
in kilograms.
Calculating Molality
• Substitute the known values into the equation
for molality and solve.
Question 1
A solution is made by dissolving 17.0 g of
lithium iodide (LiI) in enough water to make
387 mL of solution. What is the molarity of
the solution?
Answer
0.328M
Question 2
Calculate the molarity of a water solution of
CaCl2, given that 5.04 L of the solution that
contains 612 g of CaCl2.
Answer
1.09M
Question 3
What is the molality of the solution formed by
mixing 104 g of silver nitrate (AgNO3) with
1.75 kg of water?
Answer
0.350m
Question 4
Suppose that 5.25 g of sulfur (S8) is dissolved
in 682 g of the liquid solvent carbon disulfide
(CS2). What is the molality of the sulfur
solution?
Answer
0.0300m
Colligative Properties
Property whose magnitude depends
solely on the concentration of
particles, NOT on the nature of the
particles
Pressure effects Solubility
Vapor Pressure Lowering
Boiling Point
temperature where
vapor pressure of
solvent equals the
atmospheric pressure
Boiling PointElevation
Freezing Point Depression