Lesson 3.1
Types of Solutions
General Chemistry 2
Science, Technology, Engineering, and Mathematics
Solute, Solvent, and Solution
Solution
● a homogeneous mixture
of two or more pure
substances
● consists of a solute and a
solvent
2
How are solutions
classified?
3
Classification of Solutions
Solutions may be classified based on the:
● phase of the solvent
● saturation
● concentration
4
Types of Solutions Based on the Solvent Phase
Liquid Solutions
● solutions where the solvent is
a liquid
● most common type
● solid-liquid solution, liquidliquid solution, or a gasliquid solution
seawater
5
Types of Solutions Based on Solvent Phase
Solute
Solvent
Phase of the
resulting
solution
solid
liquid
gas
Examples
brine (salt in water)
liquid
liquid
rubbing alcohol
(ethanol in water)
carbonated drinks
6
Types of Solutions Based on the Solvent Phase
Solid Solutions
● solutions where the solvent is solid
● solid-solid solution, liquid-solid solution, or a gassolid solution
● e.g., H2 gas trapped in palladium metal
7
Types of Solutions Based on Solvent Phase
Different types of solid solutions and their examples
Solute
Solvent
Phase of the
resulting
solution
Examples
solid
brass (zinc in copper)
liquid
amalgam (mercury in
gold)
gas
solid
solid
hydrogen gas in
palladium metal
8
Types of Solutions Based on the Solvent Phase
Gaseous Solutions
● solutions where the solvent is gas
● solid-gas solution, liquid-gas solution, or a gas-gas
solution
● e.g., air
9
Types of Solutions Based on Solvent Phase
Solute
Solvent
Phase of the
resulting
solution
solid
liquid
gas
gas
gas
Examples
camphor in nitrogen
gas
water vapor in the air
air
10
Types of Solutions Based on Saturation
● classifying solutions based on the amount of solute
dissolved in a given amount of solvent at a specific
temperature
● solubility at a specific temperature must be
considered
11
Types of Solutions Based on the Solvent Phase
Unsaturated Solution
The amount of solute is less than the solute’s solubility at a
given volume and temperature.
12
Types of Solutions Based on the Solvent Phase
Saturated Solution
The amount of solute is equal to the solute’s solubility at a
given volume and temperature.
13
Types of Solutions Based on the Solvent Phase
Supersaturated Solution
● the amount of solute is greater than the solute’s
solubility at a given volume and temperature
● done by dissolving a solute at a higher temperature,
and subsequently cooling it down
● unstable → agitation causes crystallization
14
Liquid-Vapor Equilibrium
Crystallization in a supersaturated solution
15
Remember
In unsaturated solutions, the amount of
solute is less than the solubility capacity of
the solvent. In saturated solutions, the
amount of solute is equal to the solubility
capacity of the solvent. In supersaturated
solutions, the amount of solute is greater
than the solubility capacity of the solvent.
16
How is crystallization different
from precipitation?
17
Types of Solutions Based on Concentration
Concentrated Solution
● when a solution contains an excessively large amount
of solute
● most common unit of concentration: molarity (M)
● molarities > 1 M
● used as stock solutions in preparing diluted solution
18
Types of Solutions Based on Concentration
Concentrated Solution
concentrated sulfuric acid = 18 M
glacial (concentrated) acetic acid =
17 M
Highly concentrated
solutions emit fumes.
19
Types of Solutions Based on Concentration
Diluted Solution
● solution of low concentration
● prepared through the process of dilution
1. taking an aliquot (portion of the stock solution)
2. adding more solvent to lower the concentration
20
Lesson 3.2
Energy of Solution
Formation
General Chemistry 2
Science, Technology, Engineering, and Mathematics
Intermolecular Forces of Attraction: A Recall
● Particles of substance exhibit intermolecular forces
of attraction (IMFAs) and repulsion to one another.
● London dispersion forces (LDFs)
○ exist in all molecules, polar or nonpolar
○ strength: size, surface area, and polarizability
22
Intermolecular Forces of Attraction: A Recall
● ion-ion interactions
○ between charged substances
○ strength: Coulomb’s law
● dipole-dipole interactions
○ between polar, uncharged substances
○ strength: electronegativity of atoms in a bond
23
Intermolecular Forces of Attraction: A Recall
Induced Dipole Interactions
Ions and dipoles induce formation of temporary dipoles
on nonpolar molecules.
24
Intermolecular Forces of Attraction: A Recall
Hydrogen Bonding
Polar molecules that have H atoms bonded to O, N, or F
(e.g., H2O, NH3, and HF)
25
Intermolecular Forces of Attraction: A Recall
26
IMFA and Solution Formation
Solution Process
● solute particles must be dissolved in solvent particles
● a change in IMFAs between the particles:
○ solvent-solvent interaction
○ solute-solute interaction
○ solute-solvent interaction
27
IMFA and Solution Formation
28
IMFA and Solution Formation
Solution Process
● an energy change accompanies each of these
interaction → molar enthalpies
● IMFAs present in each substances before mixing and
IMFA present in the substances after mixing
29
IMFA and Solution Formation
Solvent
molecules are
separated.
30
IMFA and Solution Formation
Solute and solvent
molecules form a
solution.
31
IMFA and Solution Formation
Which of these enthalpy
values are positive?
Which of these enthalpy
values are negative?
32
IMFA and Solution Formation
Heat of Solution, ΔHmixing
● also referred to as ΔHsoln
● sum of all the enthalpy changes associated with each
step
33
IMFA and Solution Formation
Heat of Solution, ΔHmixing
● ΔH1 + ΔH2 is always positive
● Sign of ΔHmixing depends on the magnitudes of the
three enthalpy changes.
34
IMFA and Solution Formation
Heat of Solution, ΔHmixing
● ΔHmixing > 0 is endothermic.
○ ΔH1 + ΔH2 is greater than ΔH3.
○ Solvent-solvent and solute-solute interactions are
stronger than solute-solvent interactions.
35
IMFA and Solution Formation
Heat of Solution, ΔHmixing
● ΔHmixing < 0 is exothermic.
○ ΔH1 + ΔH2 is less than ΔH3.
○ solvent-solvent and solute-solute interactions are
weaker the solute-solvent interactions.
36
IMFA and Solution Formation
Enthalpy is not the only factor to consider in solution
formation.
○ In some cases, solute still dissolves even though
ΔHmixing > 0.
○ entropy (S)
37
IMFA and Solution Formation
Entropy (ΔS)
● the inherent tendency toward disorder in highly
favorable processes
● the chaos when solute and solvent particles mix in
comparison to their initial ordered states
38
When can one say that
mixing is favorable?
39
IMFA and Solution Formation
Solubility
● the extent to which a solute dissolve in solvent at a
particular temperature
● governed by both ΔHmixing and ΔSmixing
40
IMFA and Solution Formation
Solution Process
● Solute and solvent with similar IMFAs tend to mix
favorably.
● Solute and solvent with different IMFAs do not usually
mix.
● “like dissolves like”
41
Remember
Like dissolves like. Substances with
similar intermolecular forces of
attraction dissolve to one another.
42
Dissolution of Ionic Solids in Liquids
Solvation
● the process in which solvent molecules surround an
ion or a molecules
Hydration
● the solvation process when water is the solvent
● ions are attracted to the dipole by ion-dipole
interaction
43
Dissolution of Ionic Solids in Liquids
Dissolution of Ionic Solids
● NaCl in water
○ NaCl: crystal lattice
composed of Na+ and Cl○ H2O: polar solvent
The unit cell of NaCl
44
Dissolution of Ionic Solids in Liquids
45
Dissolution of Ionic Solids in Liquids
Dissolution of Ionic Solids
Step 1: separation of ions in the NaCl crystal lattice
● Na+ and Cl- ions
● endothermic process; ΔH1 > 0
● crystal lattice energy
46
Dissolution of Ionic Solids in Liquids
Crystal Lattice Energy (CLE)
The energy released when a mole of formula units of a
solid is formed from its constituent ions in the gas phase.
47
Dissolution of Ionic Solids in Liquids
Crystal Lattice Energy (CLE)
● reflects the strength of IMFA present in solid
○ higher CLE, strong IMFA holding the particles
together
● the energy needed to break the solid → ΔH1 = -CLE
○ always positive
48
Dissolution of Ionic Solids in Liquids
Dissolution of Ionic Solids
Step 2: separation of solvent molecules
● H2O molecules (H-bonding)
● endothermic process; ΔH2 > 0
49
Dissolution of Ionic Solids in Liquids
Dissolution of Ionic Solids
Step 3: interaction of solute and solvent particles
● Na+ interact with partially negative O atoms.
● Cl- interact with partially positive H atoms.
50
Dissolution of Ionic Solids in Liquids
Solvation Shells
● form when the solvent molecules surround ions (or
atoms or molecules) of the solute in a specific
arrangement
Hydration Shells
● the solvation shell when water is the solvent
51
Hydration Shells
52
Dissolution of Ionic Solids in Liquids
Dissolution of Ionic Solids
Step 3: interaction of solute and solvent particles
● IMFAs formed between ions and H2O molecules;
ΔH3 < 0
● hydration energy
53
Dissolution of Ionic Solids in Liquids
Hydration Energy, ΔH3
● the energy released when one mole of formula units
becomes hydrated
For a generic crystal lattice MmXn:
54
Dissolution of Ionic Solids in Liquids
Hydration Energy
Increases with charge density of ion
● charge density - the ratio of the ion’s charge to its
size or volume
● higher charge density of ion, higher magnitude of
hydration energy
55
Dissolution of Ionic Solids in Liquids
Charge densities and hydration energies of some ions
Ion
Radius, Å Charge/Radius
Hydration Energy,
kJ/mol
K+
1.33
0.75
–351
Ca2+
0.99
2.02
–1650
Cu2+
0.72
2.78
–2160
Al3+
0.50
6.00
–4750
56
Dissolution of Ionic Solids in Liquids
Heat of Solution, ΔHmixing, of Ionic Solids in Liquids
● ΔHmixing > 0 is endothermic
○ |–CLE| > |HEcation + HEanion|
● ΔHmixing < 0 is exothermic
○ |–CLE| < |HEcation + HEanion|
57
Dissolution of Liquids in Liquids
Dissolution of liquid solute in a liquid solvent: similar to
general processes
Solvation process: highly dependent on IMFAs present in
solute and solvent particles
58
Dissolution of Liquids in Liquids
Benzene (C6H6) and
toluene (C7H8)
● exhibit only LDFs;
similar IMFAs
● miscible liquids
59
Dissolution of Liquids in Liquids
Immiscible liquids
● do not mix well and form distinct layers
● exhibit distinct IMFAs
Ideal solution
● IMFAs of solute and solvent are ‘compatible’
● ΔHmixing = 0
60
Which among toluene, oxalic
acid, and benzaldehyde is the
most readily soluble in water?
Why?
61
Dissolution of Gases in Liquids
● dissolution of gases in liquids: similar to general
processes
● “like dissolves like”
○ polar gases (e.g., NO and NO2) are soluble in polar
solvents (e.g., H2O)
○ nonpolar gases (e.g., H2S) are soluble in nonpolar
solvent (e.g., hexane)
62
Dissolution of Gases in Liquids
Some polar gases enhance their solubility by reacting with
water.
63
Dissolution of Gases in Liquids
Few nonpolar gases are soluble in water because few can
react with water.
64
Dissolution of Gases in Liquids
Gases have very weak IMFAs.
○ energy needed to break IMFAs often negligible
Dissolution of gases in liquid is always exothermic.
65
Why is carbon dioxide, CO2,
soluble in water? How
important is this phenomenon?
66
Lesson 3.3
Factors Affecting
Solubility
General Chemistry 2
Science, Technology, Engineering, and Mathematics
Macroscopic View of Solution Formation
Solution formation: dynamic equilibrium between two
opposing processes—dissolution and crystallization
● forward reaction: dissolution process
● reverse reaction: crystallization process
68
Macroscopic View of Solution Formation
An illustration showing significant stages before dynamic equilibrium. The lengths
of the arrows represent the rate of dissolution (pink) and crystallization (black). 69
Factors Affecting Solubility
●
●
●
●
Nature of solute and solvent
Temperature
Pressure (or volume)
Mechanical changes
70
Factors Affecting Solubility
Solute-Solvent Interactions
● the nature of solute and solvent
● “like dissolves like”
● solute and solvent molecules with the same IMFAs are
soluble to one another
● stronger interaction between solute and solvent
molecules with similar IMFAs → more soluble
71
Factors Affecting Solubility
Solute-Solvent Interactions
● Ionic solids are more soluble in polar solvents than
nonpolar solvents.
○ Ion-dipole interactions are stronger than ioninduced dipole interactions.
72
Factors Affecting Solubility
Solute-Solvent Interactions
● Polar molecules are soluble
in polar solvents.
○ dipole-dipole interactions
● Nonpolar molecules are
soluble in nonpolar
solvents.
○ LDFs
Oil forms layer with water because
their IMFAs are ‘incompatible.’
73
Factors Affecting Solubility
Solute-Solvent Interactions
Compounds with H-bonds are soluble in polar solvents.
74
Factors Affecting Solubility
Solute-Solvent Interactions
● solubility → depends on the amount of solute and
solvent present in solution
75
Dissolution of Ionic Solids in Liquids
Solubilities of some alcohols in water and hexane in mol ROH/100 g solvent at 20 oC
Solubility in
water, H2O
Solubility in
hexane, C6H14
Methanol, CH3OH
∞
0.12
Ethanol, CH3CH2OH
∞
∞
Propanol, CH3CH2CH2OH
∞
∞
Butanol, CH3CH2CH2CH2OH
0.11
∞
Pentanol, CH3CH2CH2CH2CH2OH
0.030
∞
Hexanol, CH3CH2CH2CH2CH2CH2OH
0.0058
∞
Alcohol
76
Suppose you want a drug that is
soluble in water. How will you
increase its solubility in water if
you can change its chemical
structure?
77
Factors Affecting Solubility
Temperature
● the most evident factor that
affects solubility
● Most solids are soluble at
higher temperatures.
Solubilities of some ionic compounds
as a function of temperature
78
Factors Affecting Solubility
Temperature
● the steeper the slope of the
curve = the greater the effect
of temperature on the
solubility
Solubilities of some ionic compounds
as a function of temperature
79
Which of the nitrates found in
the figure from the previous
slide would you use if you want
a higher concentration of
nitrate in solution at 10 oC?
Explain your answer.
80
Factors Affecting Solubility
Temperature
● Solubility of gases decreases
as temperature increases.
Solubilities of some gases as a
function of temperature
81
Remember
The solubility of most solids increases
with temperature. The solubility of
gases decreases with temperature.
82
Factors Affecting Solubility
Pressure (or Volume)
● does not significantly affect the solubility of solids and
liquids
● greatly affects the solubility of gases
○ increasing the pressure increases the amount of
dissolved gas in solution
83
Factors Affecting Solubility
The effect of pressure on solubility of gases
84
Factors Affecting Solubility
Henry’s Law
The relationship between pressure of the gas over the
solution and solubility of the gas
cgas
P
kH
solvent
=
=
=
concentration of gas (M)
pressure of gas (atm)
Henry’s constant at a specific gas-
combination and temperature (M/atm)
85
Factors Affecting Solubility
Mechanical Changes
● stirring
○ improves dissolution
○ adds heat to the solution
○ increases the probability of contact between solid
solute particles and solvent particles
86
Factors Affecting Solubility
Mechanical Changes
● increase in surface area
○ improves dissolution
○ e.g., granules are more soluble than blocks.
87