Download Properties of Solutions

Properties
of
Solutions
Brown, LeMay Ch 13
AP Chemistry
Edited significantly LJC 2015
CaCl2 (aq)
1
13.1: Types of Solutions
Example
Air (g in g)
Soda (g in l)
H2 in Pt (g in s)
Alcoholic
beverages (l in l)
Solvent
Solute
Sea water (s in l)
Brass (s in s)
3
13.1: Types of Solutions
Example
Air (g in g)
Soda (g in l)
H2 in Pt (g in s)
Alcoholic
beverages (l in l)
Solvent
N2
Solute
O2
H2O
Pt
CO2
H2
H2O
C2H5OH
Sea water (s in l)
H2O
Brass (s in s)
Copper
(55% – 90%)
NaCl
(one of many salts)
Zinc
(10% – 45%) 4
Solutions
• Solutions are homogeneous mixtures of two or
more pure substances.
• In a solution, the solute is dispersed uniformly
throughout the solvent.
Solutions
How does a solid dissolve
into a liquid?
What ‘drives’ the dissolution
process?
What are the energetics of
dissolution?
How Does a Solution Form?
1. Solvent molecules attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy (DH) changes with each interaction broken or
formed.
Ionic solid dissolving in water
How Does a Solution Form?
1. Solvent molecules attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy (DH) changes with each interaction broken or
formed.
How Does a Solution Form?
The ions are solvated
(surrounded by solvent).
If the solvent is water, the
ions are hydrated.
What intermolecular force
is present?
How Does a Solution Form?
The ions are solvated
(surrounded by solvent).
If the solvent is water, the
ions are hydrated.
What intermolecular force
is present?
ION-DIPOLE
Energy Changes in Solution
To determine the enthalpy
change, we divide the
process into 3 steps.
1.
2.
3.
Energy Changes in Solution
To determine the enthalpy
change, we divide the
process into 3 steps.
1. Separation of solute
particles.
2. Separation of solvent
particles to make
‘holes’.
3. Formation of new
interactions between
solute and solvent.
Energy Changes in Solution
Predict: which of these
steps would likely be
endothermic?
Which would be
exothermic?
Enthalpy Changes in Solution
Start
The enthalpy change of the overall process
depends on DH for each of these steps.
Enthalpy changes during dissolution
DHsoln = DH1 + DH2 + DH3
The enthalpy of
solution, DHsoln,
can be either
positive or
negative.
DHsoln (MgSO4)= -91.2 kJ/mol --> exothermic
DHsoln (NH4NO3)= 26.4 kJ/mol --> endothermic
Why do endothermic processes
sometimes occur spontaneously?
Some processes, like
the dissolution of
NH4NO3 in water, are
spontaneous at room
temperature even
though heat is
absorbed, not
released.
Enthalpy Is Only Part of the Picture
There is also ENTROPY.
Entropy is a measure of:
• Dispersal of energy in the
system.
• Number of microstates
(arrangements) in the
system.
• OR, put simply, degree of
disorder
(more on this in chap 19)
Which beaker is the favored
energy state?
Enthalpy Is Only Part of the Picture
There is also ENTROPY.
Entropy is a measure of:
• Dispersal of energy in the
system.
• Number of microstates
(arrangements) in the
system.
• OR, put simply, degree of
disorder
(more on this in chap 19)
(b) has greater entropy,
 is the favored state
Entropy changes during dissolution
What were the three steps of
solution formation?
1.
2.
3.
What would the entropy
change during each step
be?
SAMPLE EXERCISE 13.1 Assessing Entropy Change
In the process illustrated below, water vapor reacts with excess solid sodium
sulfate to form the hydrated form of the salt. The chemical reaction is
Does the entropy of the system increase or decrease?
Dissolution vs reaction
Ni(s) + HCl(aq)
NiCl2(aq) + H2(g)
dry
NiCl2(s)
• Dissolution is a physical change—you can get back the original
solute by evaporating the solvent.
• If you can’t, the substance didn’t dissolve, it reacted.
Degree of saturation
• Saturated solution
Solvent holds as much
solute as is possible at
that temperature.
Undissolved solid
remains in flask.
Dissolved solute is in
dynamic equilibrium with
solid solute particles.
Degree of saturation
• Unsaturated Solution
Less than the maximum
amount of solute for that
temperature is dissolved
in the solvent.
No solid remains in flask.
Degree of saturation
• Supersaturated
Solvent holds more solute than is normally possible at
that temperature.
These solutions are unstable; crystallization can often
be stimulated by adding a “seed crystal” or scratching
the side of the flask.
Degree of saturation
How can you tell whether a solution is unsaturated,
saturated or supersaturated?
How much solute can be dissolved in a solution?
SOLUBILITY is how much solute can be dissolved in a
solution.
More on this in Chap 17
(solubility products, p 739)
Factors Affecting Solubility
What do you observe?
How does this illustrate “Like Dissolves Like”?
Factors Affecting Solubility
Example: ethanol in water
Ethanol = CH3CH2OH
The stronger the
intermolecular
attractions
between solute
and solvent, the
more likely the
solute will
dissolve.
Intermolecular forces = H-bonds; dipole-dipole; dispersion
Ions in water also have ion-dipole forces.
Factors Affecting Solubility
Which would be more soluble in water:
cyclohexane or glucose?
Factors Affecting Solubility
Glucose (which has hydrogen bonding) is very
soluble in water.
Cyclohexane (which only has dispersion forces) is
not water-soluble.
Factors Affecting Solubility
• Which vitamin is fat soluble?
• Which vitamin is water soluble?
Factors Affecting Solubility
• Vitamin A is soluble in nonpolar compounds (like
fats).
• Vitamin C is soluble in water.
Which
vitamin is
water-soluble
and which is
fat-soluble?
Gases in Solution
What do you observe about the solubility of
gases in water? Why? (Be CL.EV.ER)
Gases in Solution
• In general, the solubility of gases in water increases
with increasing mass.
Why?
• Larger molecules have stronger dispersion forces.
Gases in Solution
Increasing
pressure above
solution forces
more gas to
dissolve.
• The solubility of
liquids and solids
does not change
appreciably with
pressure.
• But, the solubility of a
gas in a liquid is
directly proportional
to its pressure.
Henry’s Law
Sg = kPg
where
• Sg is the solubility of the
gas;
• k is the Henry’s law
constant for that gas in
that solvent;
• Pg is the partial pressure
of the gas above the
liquid.
Temperature
What do you observe
about the solubility of
solids as T increases?
Temperature
• What do you observe
about the solubility of
gases as T increases?
• Is this the same or
different from what you
observed for solids?
Carbonated soft drinks are
more “bubbly” if stored in
the refrigerator.
Warm lakes have less O2
dissolved in them than
cool lakes.
Practice: what IMF’s are
involved in making this
solution?
NaCl (s) → Na+ (aq) + Cl- (aq)
Practice: what IMF’s are
involved in making this
solution?
NaCl (s) + H2O (l) → Na+ (aq) + Cl- (aq)
• Ion-dipole interactions
> H-bonds (H2O···H2O)
< Ionic bonds (Na+ Cl-)
• The increase in disorder also drives the dissolving process.
http://phet.colorado.edu/en/simulation/soluble-salts
Ways to Express Concentration (Card #2)
• Mass Percent
• No longer on AP Exam!
mass solute
%
100
mass solute  mass solvent
• Mole Fraction: commonly used for gases
mol A
XA 
total mol
• Molarity: commonly used for solutions

Varies with T
mol solute
M
Liters solution
• Molality: commonly used for colligative properties
Does not vary with T
No longer on AP exam!

mol solute
m
kg solvent
42
Solubility Vocabulary (Card #3)

http://www.youtube.com/watch?v=VTmfQUNLlMY
 Saturated: a solution that is in equilibrium with
undissolved solute (appears as solution and
crystals)
 Solubility: the amount of solute needed to
form a saturated solution
 Unsaturated: a solution containing less than
the saturated amount (appears as solution
only)
 Supersaturated: a solution containing more
than the saturated amount, yet appears
unsaturated.
43
Solubility
Solubility: go to the temperature and up
to the desired line, then across to the Yaxis. This is how many g of solute are
needed to make a saturated solution of
that solute in 100g of H2O at that
particular temperature.
At 40oC, the solubility of KNO3 in 100g of
water is 64 g. In 200g of water, double
that amount. In 50g of water, cut it in
half.
Supersaturated
If 120 g of NaNO3 are added to 100g of
water at 30oC:
1) The solution would be
SUPERSATURATED, because there is
more solute dissolved than the
solubility allows
2) The extra 25g would precipitate out
3) If you heated the solution up by
24oC (to 54oC), the excess solute would
dissolve.
Unsaturated
If 80 g of KNO3 are added to 100g of
water at 60oC:
1) The solution would be
UNSATURATED, because there is less
solute dissolved than the solubility
allows
2) 26g more can be added to make a
saturated solution
3) If you cooled the solution down by
12oC (to 48oC), the solution would
become saturated
13.4: Factors Affecting Solubility
(Card #4)
1.
“Like dissolves like.”
•
Miscible: liquids that mix (polar or ionic solute with polar
solvent, or nonpolar with nonpolar)
•
Immiscible: liquids that do not mix (polar or ionic solute with
nonpolar solvent)
•
Covalent network solids do not dissolve in polar or nonpolar
solvents.
47