Download General Chemistry I

General Chemistry I
Dr. PHAN TẠI HUÂN
Faculty of Food Science and Technology
Nong Lam University
Module 6:
Solutions
• Definitions (solutions: solvent, solute, solubility,
concentration; interactions between solvent phase and
solute molecules; properties of polar and non-polar
solvents; properties of colloid).
• Chemical equilibrium for solutions (self-ionisation of
water, solvation, pH; Acids and Bases; solution
equilibrium with precipitation).
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Solution
• A solution is defined as a homogeneous mixture of
substances in which no settling occurs.
• A solution consists of a solvent and one or more solutes,
whose proportions vary from one solution to another.
• The solvent is the medium in which the solutes are
dissolved. The fundamental units of solutes are usually ions
or molecules.
• Water is the most important solvent, and compounds
dissolved in water are said to be in aqueous solution.
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Solution
• In reality, any combination of the three states can be
considered a solution.
• Usually a solution is formed by dissolving a solid
(e.g., sugar) in a liquid (e.g., water).
• Air is a solution which is a mixture of various gases.
• Carbonated water (soda) is a mixture of a gas (CO2)
dissolved in a liquid (H2O).
• Even alloys such as gold-silver alloys are solutions
containing two solids.
A true solution is a solution which has only one solvent
with one or more solutes.
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2
The concept of solubility
• Solubility of a substance is defined as the amount of the
substance that will dissolve in a particular solvent.
• Solubilities vary tremendously.
• At one extreme, some substances form solutions in all
proportions and are said to be miscible. For example,
acetone and water can be mixed in any proportion, from
pure water to pure acetone.
• At the other extreme, a substance may be insoluble in
another. One example is common salt, NaCl, whose
solubility in gasoline is virtually zero.
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Determinants of solubility
• Many combinations display solubility that is between the two
extremes of miscible and insoluble. In other words, the
substance dissolves, but there is a limit to the amount of
solute that will dissolve in a given amount of solvent.
• Concentrations of solutions are expressed in terms of either
the amount of solute present in a given mass or volume of
solution, or the amount of solute dissolved in a given mass or
volume of solvent.
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3
Percent by mass
• Concentrations of solutions may be expressed in terms of
percent by mass of solute, which gives the mass of solute per
100 mass units of solution. The gram is the usual mass unit.
percent sulute =
mass of solute
x 100%
mass of solution
• Thus, a solution that is 10% calcium gluconate,
Ca(C6H11O7)2, by mass contains 10 grams of calcium
gluconate in 100 grams of solution. This could be described
as 10 grams of calcium gluconate in 90 grams of water.
• Unless otherwise specified, percent means percent by mass,
and water is the solvent.
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Molarity vs. Molality
• The Molarity, M, of a solution is defined as the number of
moles of the solute per liter of solution.
Molarity =
moles of solute mol
=
liter of solution
L
• The molality, m, of a solution is defined as the number of
moles of the solute per kilogram of solvent.
Molality =
moles of solute
kilograms of solvent
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4
Normality
• The normality, N, of a solution is the number of
equivalents of solute per liter of solution.
• The equivalent is usually defined in terms of a chemical
reaction. For acid-base reactions, an equivalent is the
amount of substance that will react or form 1 mole of
hydrogen (H+) or hydroxide (OH-) ions. For redox
(oxidation-reduction) reactions, an equivalent is the amount
of substance that will react or form 1 mole of electrons.
Normality =
number of equivalents
1 liter of solution
N = nM
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Exercise
• How many grams of H2O must be used to dissolve 50
grams of sucrose to prepare a 1.25 m solution of sucrose,
C12H22O11?
Ans:
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Exercise
• Hydrogen peroxide disinfectant typically contains 3.0%
by mass. Assuming that the rest of the contents is water,
what is the molality of this disinfectant?
Ans:
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Spontaneity of the dissolution process
• A process is favored by (1) a decrease in the energy of the
system, which corresponds to an exothermic process, and (2)
an increase in the disorder, or randomness, of the system.
• The energy change that accompanies a dissolution process is
called the heat of solution, ∆Hsolution. It depends mainly on
how strongly solute and solvent particles interact.
• A negative value of ∆Hsolution designates the release of heat.
• The main interactions that affect the dissolution of a solute
in a solvent follow:
– Weak solute–solute attractions favor solubility.
– Weak solvent–solvent attractions favor solubility.
– Strong solvent–solute attractions favor solubility.
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Spontaneity of the dissolution process
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Like dissolves Like
• Solubility is a complex phenomenon that depends on the
balance of several properties.
• The general features of solubility are summarized by the
expression like dissolves like.
• Substances that dissolve in each other usually have similar
types of intermolecular interactions.
• One substance dissolves in another if the forces of
attraction between the solute and the solvent are similar to
the solvent–solvent and solute–solute interactions.
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7
Dissolution of liquids
• Water and methanol are alike in that both substances contain
O-H groups that form hydrogen bonds readily. When these
liquids are mixed, H2O... H2O hydrogen bonds and CH3OH...
CH3OH hydrogen bonds break, but H2O...CH3OH hydrogen
bonds form.
• The net result is that the degree of hydrogen bonding in the
solution is about the same as in either of the pure liquids,
making these two liquids miscible.
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Dissolution of liquids
• The intermolecular interactions of
octane and cyclohexane are alike.
• Octane and cyclohexane have low
polarities, so these molecules in the
pure liquids are held together by the
dispersion forces caused by their
polarizable electron clouds.
• Dispersion forces in solutions of
octane and cyclohexane are about
the same as in the pure liquids. So
these two liquids are miscible.
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8
Dissolution of liquids
• Water and octane are not alike and nearly insoluble in each other.
• Octane does not form hydrogen bonds, so the only forces of
attraction between water molecules and octane molecules are
dispersion forces.
• Because hydrogen bonds are stronger than dispersion forces, the
cost of disrupting the hydrogen-bonding network in water is far
greater than the stability gained from octane–water dispersion
forces.
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Dissolution of liquids
• Some liquids can interact with other
substances in multiple ways. Acetone, for
instance, has a polar CO bond and a
three-carbon bonding framework.
• The bonding framework is similar to that
of a hydrocarbon, so acetone mixes with
cyclohexane and octane.
• The polar CO group makes acetone
miscible with other polar molecules such
as acetonitrile .
• The polar oxygen atom in acetone has
lone pairs of electrons that can form
hydrogen bonds with hydrogen atoms of
ammonia or water.
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9
Exercise
• Give a molecular explanation for the following trend in alcohol
solubilities in water:
n-Propanol
n-Butanol
n-Pentanol
n-Hexanol
CH3CH2CH2OH
Completely miscible
CH3CH2CH2CH2OH
1.1 M
CH3CH2CH2CH2CH2OH
0.30 M
0.056 M
CH3CH2CH2CH2CH2CH2OH
Strategy
• Solubility limits depend on the stabilization generated by
solute–solvent interactions balanced against the destabilization
that occurs when solvent–solvent interactions are disrupted by
solute.
• Intermolecular interactions involving water and alcohol
molecules must be examined.
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Solubility of solids: network solids
• Network solids such as diamond, graphite, or silica cannot
dissolve without breaking covalent chemical bonds.
• Because intermolecular forces of attraction are always
much weaker than covalent bonds, solvent–solute
interactions are never strong enough to offset the energy
cost of breaking bonds.
• Covalent solids are insoluble in all solvents, although they
may react with specific liquids or vapors.
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Dissolution of solids: metalic solids
• Metals do not dissolve in water, because they contain
extensive delocalized bonding networks that must be
disrupted before the metal can dissolve.
• A few metals react with water, and several react with
aqueous acids, but no metal will simply dissolve in water.
Likewise, metals do not dissolve in nonpolar liquid
solvents.
Zn(s) + 2H3O+ (aq) Î Zn2+(aq) + H2(g) + 2H2O(l)
• The aqueous medium dissolves the metal by a chemical
reaction that converts the insoluble metal into soluble
cations. The solution produced when zinc reacts with
aqueous HCl is an aqueous solution of ions, not a solution
of Zn metal in water.
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Dissolution of solids: molecular solids
• At the opposite extreme, molecular solids contain individual
molecules bound together by various combinations of
dispersion forces, dipole forces, and hydrogen bonds.
Conforming to “like dissolves like,” molecular solids
dissolve readily in solvents with similar types of
intermolecular forces.
• Nonpolar , for instance, is soluble in nonpolar liquids such as
carbon tetrachloride .
• Many organic compounds are molecular solids that dissolve
in organic liquids such as cyclohexane and acetone.
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Dissolution of solids: molecular solids
• Hydrogen bonding makes sugars such as sucrose and glucose
highly soluble in water.
• When glucose dissolves, hydrogen bonds between water and
glucose replace the hydrogen bonds lost by the water
molecules of the solvent. This balance means that the energy
requirements for solution formation are small, and glucose is
quite soluble in water.
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Dissolution of solids: molecular solids
• Hydrogen bonding allows water to dissolve materials that
form hydrogen bonds.
• On the other hand, naphthalene, a similarly sized solid
hydrocarbon limited to dispersion forces, is nearly insoluble
in water.
Î The best solvent for a molecular solid is one whose
intermolecular forces match the forces holding the molecules
in the crystal.
– For a solid held together by dispersion forces, good
solvents are nonpolar liquids such as carbon tetrachloride
(CCl4) and cyclohexane (C6H12).
– For polar solids, a polar solvent such as acetone works
well.
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Example: solubility of vitamin
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Solubility of solids: ionic solids
• The ability of an ionic solid to go into solution depends most
strongly on its crystal lattice energy, or the strength of
attractions among the particles making up the solid.
• Crystal lattice energies are always negative:
M(g) + X(g) Î MX(s) + energy
• If the solvent is water, the energy that must be supplied to
expand the solvent includes that required to break up some
of the hydrogen bonding between water molecules.
• The third major factor contributing to the heat of solution is
the extent to which solvent molecules interact with particles
of the solid. The process in which solvent molecules
surround and interact with solute ions or molecules is called
solvation. When the solvent is water, the more specific term
is hydration.
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Solubility of solids: ionic solids
• A cluster of water molecules surrounds each ion in solution.
Notice how the water molecules are oriented so that their
dipole moments align with charges of the ions. The partially
negative oxygen atoms of water molecules point toward
cations, whereas the partially positive hydrogen atoms of
water molecules point toward anions.
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Solubility of solids: ionic solids
• Hydration energy is defined as the energy change involved
in the (exothermic) hydration of one mole of gaseous ions.
Mn+(g) + xH2O Î M(H2O)xn+ + energy (for cation)
Xy- (g) + rH2O Î X(H2O)ry- + energy (for anion)
• Hydration is usually highly exothermic for ionic or polar
covalent compounds, because the polar water molecules
interact very strongly with ions and polar molecules.
• The overall heat of solution for a solid dissolving in a liquid
is equal to the heat of solvation minus the crystal lattic
energy.
∆Hsolution = (heat of solvation) - (crystal lattice energy)
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Solubility of solids: ionic solids
• Magnitudes of crystal lattice and hydration energies generally
increase with increasing charge and decreasing size of ions
(ionic charge densities increase).
• Hydration energies and lattice energies are usually of about
the same magnitude for low-charge species, so the
dissolution process is slightly endothermic for many ionic
substances.
• As the charge-to-size ratio (charge density) increases for ions
in ionic solids, the magnitude of the crystal lattice energy
usually increases more than the hydration energy. For ex.
aluminum fluoride, AlF3; magnesium oxide, MgO; and
chromium(III) oxide, Cr2O3 are very endothermic and not
very soluble in water.
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Ionic radii, charge/radius ratios, and hydration
energies for some cations
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Dissolution of gases
The only gases that dissolve appreciably in water are:
• (1) those that are capable of hydrogen bonding (such as HF),
• (2) those that ionize (such as HCl, HBr, and HI),
• (3) those that react with water (such as CO2).
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Rates of dissolution and saturation
• When a solid is placed in water, some of its particles solvate and
dissolve. The rate of this process slows as time passes because
the surface area of the crystals gets smaller and smaller.
• At the same time, the number of solute particles in solution
increases, so they collide with the solid more frequently. Some
of these collisions result in recrystallization.
• The rates of the two opposing processes become equal after
some time. The solid and dissolved ions are then in equilibrium
with each other.
• After equilibrium is established, no more solid dissolves without
the simultaneous crystallization of an equal mass of dissolved
ions.
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Rates of dissolution and saturation
• The solubilities of many solids increase at
higher
temperatures.
Supersaturated
solutions contain higher-than-saturated
concentrations of solute.
• The saturated solution is cooled slowly,
without agitation, to a temperature at which
the solute is less soluble. At this point, the
resulting supersaturated solution is
metastable (temporarily stable).
• A supersaturated solution produces crystals
rapidly if it is slightly disturbed or if it is
“seeded” with a dust particle or a tiny
crystal.
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Effect of temperature on solubility
• LeChatelier’s Principle: A system at equilibrium, or
changing toward equilibrium, responds in the way that
tends to relieve or “undo” any stress placed on it.
• Many ionic solids dissolve by endothermic processes.
Their solubilities in water usually increase as heat is
added and the temperature increases.
Endothermic: reactants + heat Î products
• For example:
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Effect of pressure on solubility
• Changing the pressure has no
appreciable effect on the solubilities
of either solids or liquids in liquids.
• The solubilities of gases in all
solvents increase, however, as the
partial pressures of the gases
increase.
• Henry’s Law: The pressure of a gas
above the surface of a solution is
proportional to the concentration of
the gas in the solution.
The relationship is valid at low
Pgas = kCgas
concentrations and low pressures.
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Colloids
• Particles whose dimensions are between 1 nanometer and 1
micrometer, called colloids, are larger than the typical
molecule but smaller than can be seen under an optical
microscope.
• When a colloid is mixed with a second substance, the colloid
can become uniformly spread out, or dispersed, throughout
the dispersing medium. Such a dispersion is a colloidal
suspension that has properties intermediate between those of
a true solution and those of a heterogeneous mixture.
Mixture
suspension
colloidal dispersion
solution
Example
sand in water
starch in water
sugar in water
Approximate Particle Size
larger than 10,000 Å
10–10,000 Å
1–10 Å
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Types of colloids
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Aqueous solutions: an introduction
• Approximately 3/4 of the earth’s surface is covered with water.
• The body fluids of all plants and animals are mainly water.
Î Many important chemical reactions occur in aqueous (water)
solutions, or in contact with water.
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Electrolytes and extent of ionization
• Solutes that are water-soluble can be classified as either
electrolytes or nonelectrolytes.
• Electrolytes are substances whose aqueous solutions
conduct electric current.
• Strong electrolytes are substances that conduct electricity
well in dilute aqueous solution.
• Weak electrolytes conduct electricity poorly in dilute
aqueous solution.
• Aqueous solutions of nonelectrolytes do not conduct
electricity.
• Electric current is carried through aqueous solution by the
movement of ions. The strength of an electrolyte depends
on the number of ions in solution and also on the charges
on these ions.
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Electrolytes and extent of ionization
• Three major classes of solutes are strong electrolytes: (1) strong
acids, (2) strong bases, and (3) most soluble salts. These
compounds are completely or nearly completely ionized (or
dissociated) in dilute aqueous solutions, and therefore are strong
electrolytes.
• Dissociation refers to the process in which a solid ionic
compound, such as NaCl, separates into its ions in solution:
• Ionization refers to the process in which a molecular compound
separates or reacts with water to form ions in solution:
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Acid–Base
The Arrhenius theory (1884)
• An acid is a substance that contains hydrogen and produces
H in aqueous solution.
• A base is a substance that contains the OH (hydroxyl) group
and produces hydroxide ions, OH, in aqueous solution.
• Neutralization is defined as the combination of H ions with
OH ions to form H2O molecules.
H+ (aq) + OH- (aq) Î H2O (l) (neutralization)
The hydronium ion (hydrated hydrogen ion)
• The hydrated hydrogen ion is the species that gives aqueous
solutions of acids their characteristic acidic properties.
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Acid–Base
The Brønsted–Lowry theory (1923)
• An acid is defined as a proton donor and a base is defined as a
proton acceptor.
• An acid–base reaction is the transfer of a proton from an acid to
a base.
• The ionization of hydrogen chloride, HCl, in water is an acid–
base reaction in which water acts as a base or proton acceptor.
• We can describe Brønsted–Lowry acid–base reactions in terms
of conjugate acid–base pairs. These are two species that differ
by a proton.
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Acid–Base
The Lewis theory (1923)
• An acid is any species that can accept a share in an electron
pair. A base is any species that can make available, or
“donate,” a share in an electron pair.
• These definitions do not specify that an electron pair must
be transferred from one atom to another—only that an
electron pair, residing originally on one atom, must be
shared between two atoms. Neutralization is defined as
coordinate covalent bond formation.
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The autoionization of water
• Water is said to be amphiprotic; that is, H2O molecules
can both donate and accept protons.
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Chemical equilibrium
• Reactions that do not go to completion and that can occur in
either direction are called reversible reactions.
rate f = k f [C ] [D ]
c
d
rate r = k r [A ] [B ]
a
b
• At equilibrium, ratef = rater
[C ]eq [D ]eq
kf
=
a
b
k r [A ]eq [B ]eq
c
d
Kc
[C ] [D ]
=
[A ] [B ]
c
d
eq
a
eq
b
eq
eq
(For any pure liquid or pure solid, the activity is taken as 1.)
• Both kf and kr are constant, so kf/kr is also a constant and
given a special name and symbol the equilibrium constant, Kc
or simply K.
⇒ − ΔG o = RTlnK
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The autoionization of water
• The equilibrium constant is known as the ion product for
water and is usually represented as Kw.
Kw = [H3O+][OH-]
• Careful measurements show that, in pure water at 25°C,
[H3O+]=[OH-]= 1.0 x 10-7 mol/L
Kw = [H3O+][OH-]= 1.0 x 10-14 (at 25°C)
• Although this expression was obtained for pure water, it is
also valid for dilute aqueous solutions at 25°C.
• This is one of the most useful relationships chemists have
discovered. It gives a simple relationship between H3O+ and
OH- concentrations in all dilute aqueous solutions.
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Exercise
•
Calculate the concentrations of H3O+ and OH- ions in a
0.05 M HNO3 solution.
Ans:
50
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The pH and pOH scales
• In common chemical applications, the concentration of
hydronium ion, the measure of a solution´s acidity, ranges
from fairly large (1 M) to very small (10-14M).
• The pH and pOH scales provide a convenient way to
express the acidity and basicity of dilute aqueous solutions
that avoids using very small numbers.
• The pH and pOH of a solution are defined as
pH = - log [H3O+]
or
[H3O+] = 10-pH
pOH= - log [OH-]
or
[OH-] = 10-pOH
• It is convenient to describe the autoionization of water in
terms of pKw.
pKw = -log Kw
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The pH and pOH scales
[H3O+][OH-]= 1.0 x 10-14 (at 25°C)
Î pH + pOH = 14
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Exercise
• Calculate [H3O+ ], pH, [OH -], and pOH for a 0.015 M
Ca(OH)2 solution.
Ans:
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The pH and pOH scales
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Strong and weak acids
• Strong acids ionize (separate into hydrogen ions and stable
anions) completely, or very nearly completely, in dilute
aqueous solution.
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Strong and weak acids
• Weak acids cannot dissociate completely. They undergo the
same type of dissociation as that of strong acids and bases,
but the extent of dissociation is very little compared to
strong acid or strong base dissociations (usually less than
5%) .
• Acid-ionization constant (Ka):
[H O ][A ]
=
+
Ka
-
3
[HA]
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Strong and weak acids
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Exercise
• Find the degree of ionization of 0.1 M solution of acetic
acid (CH,COOH). Also find the pH of the solution.
(The acid-ionization constant of acetic acid is 1.8 x 1 0-5)
Ans:
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Strong bases, insoluble bases, and weak bases
• Most common bases are ionic metal hydroxides in the solid state.
• Strong bases are soluble in water and are dissociated completely
in dilute aqueous solution.
• Common strong bases
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Strong bases, insoluble bases, and weak bases
• Other metals form ionic hydroxides, but these are so
sparingly soluble in water that they cannot produce strongly
basic solutions. They are called insoluble bases or
sometimes sparingly soluble bases. Typical examples
include Cu(OH)2, Zn(OH)2, Fe(OH)2, and Fe(OH)3.
• Common weak bases are molecular substances that are
soluble in water but form only low concentrations of ions in
solution. The most common weak base is ammonia, NH3.
• Just like the acid-ionization constant, there is also the baseionization (Kb)
[NH ][OH ]
+
Kb =
-
4
[NH ]
3
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Dissociation of polyprotic acids
• Some acids have two or more protons that can be released
upon dissociation. Such acids are called polyprotic acids.
• Sulfuric acid is a polyprotic acid that can lose two protons
in solution.
• The first ionization is complete because sulfuric acid is a
strong acid.
Dissociation constant, Kal = very large
• In the second ionization, an equilibrium exists because
hydrogen sulfate ion (HSO4-) is not as strong as H2SO4.
Dissociation constant, Ka2 = 1.2 x l0-2
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Differentiating acidic and basic salts
• Salt solutions can be acidic, neutral, or basic.
• A salt of a weak acid and a strong base gives a basic aqueous
solution.
NaCN is a salt of a weak acid (HCN) and a strong base (NaOH).
• A salt of a weak base and a strong acid gives an acidic aqueous
solution.
Zn(NO)3 is salt of a weak base (Zn(OH)3) and a strong acid (HNO3).
• A salt of a strong base and a strong acid gives a neutral aqueous
solution.
NaCl is a salt of a strong acid (HC1) and a strong base (NaOH).
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Common-ion effect
• When a salt is added to a solution containing either the same
cation or anion, there will be changes in the solubility
because of what is commonly known as common-ion effect.
• The phenomenon can be best explained in terrns of Le
Chatelier's principle.
• For ex. adding magnesium fluoride (MgF2) to a solution of
sodium fluoride (NaF).
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Solubility guidelines
for common ionic compounds in water
• Compounds whose solubility in water is less than about 0.02
mol/L are usually classified as insoluble compounds.
• No gaseous or solid substances are infinitely soluble in water.
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Solubility guidelines
for common ionic compounds in water
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Bonding, solubility, electrolyte characteristics, and
predominant forms of solutes in contact with water
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Solubility product principle
• In equilibria that involve slightly soluble compounds in water,
the equilibrium constant is called a solubility product
constant, Ksp ,or solubility product.
Ksp = [Mz+]y[Xy-]z
• In general, the solubility product expression for a compound is
the product of the concentrations of its constituent ions, each
raised to the power that corresponds to the number of ions in
one formula unit of the compound. The quantity is constant at
constant temperature for a saturated solution of the compound.
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Ion product
• Ion product (reaction quotient, Qsp) is the product of the
concentrations of the ions from the compound (solute) in a
solution, each concentration raised to a power equal to its
coefficient in the balanced equation. In other words, the
expression for the ion product is the same as that of Ksp.
• By comparing the ion product of a compound against its Ksp,
we can predict whether or not precipitation is likely to occur.
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Exercise
• If 100 mL of 0.00075 M sodium sulfate, Na2SO4, is mixed
with 50 mL of 0.015 M barium chloride, BaCl2, will a
precipitate form? (Ksp for BaSO4 = 1.1 x 10-10).
• Ans:
69
Buffers
• A solution which resists changes in pH when small amounts of
acid or base are added to it is called a buffer solution.
• A buffer solution is usually a mixture of a weak acid and its
conjugate base or a weak base and its conjugate acid. A buffer
solution contains equilibrium amounts of acid and base species.
• If a strong acid is added to this buffer solution, the H+
concentration increases and they react with the base (HCO3)thereby decreasing the H+ concentration and maintaining the
initial pH.
• On the other hand, if a strong base is added, the OH- ions
supplied by it will react with the H+ so that more of H2CO3 ,
will dissociate thereby restoring the initial pH.
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Buffers
• Buffers do not have unlimited capacity to resist pHchange. The buffer capacity of a buffer depends on the
nature of the buffer and the amount of acid and conjugate
base present in the solution.
• The Henderson-Hasselbalch equation can be used to relate
the pH of a buffer and the concentrations of base and acid.
• For a weak acid-conjugate base buffer:
pH = pK a + log
[conjugate base]
[weak acid]
• For a weak base-conjugate acid buffer.
pOH = pK b + log
[conjugate acid]
[weak base]
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Exercise
• Calculate the concentration of H3O+ and the pH of a buffer
solution that is 0.10 M in CH3COOH and 0.20 M in
NaCH3COO.
Ans:
72
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Titration curves of acids and bases
• Consider the reaction involving 50 ml of a 0.1 M solution of
HC1, titrated with 0.1 M solution of NaOH.
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Titration curves of acids and bases
• Acid-base titrations are reactions by which we can determine
the amount of acid or base present in a solution.
• This is done by reacting the solution with a base or acid (of
known concentration), and by measuring the volume of the
known acid or base used up in the process.
• The equivalence point denotes the point at which equivalent
amounts of acid and base have reacted.
• To know the equivalence point, we usually add an indicator
which will change its color close to the equivalence point.
• An indicator is an organic dye; its color depends on the
concentration of H3O+ ions, or pH, in the solution. By the
color an indicator displays, it “indicates” the acidity or
basicity of a solution.
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Indicator
• Methyl red is red at pH 4 and below; it is yellow at pH 7
and above. Between pH 4 and pH 7 it changes from red to
redorange, to orange, to yellow.
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Indicator
• Bromthymol blue is yellow at pH 6 and below; it is blue at
pH 8 and above. Between pH 6 and 8 it changes from yellow
to yellowgreen, to green, to blue-green, to blue.
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Indicator
• Phenolphthalein is colorless below pH 8 and bright pink above
pH 10. It changes from colorless to pale pink, to pink, to
bright pink in the pH range 8 to 10.
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Indicator
• The pH at which the color change occurs is characteristic of
each indicator. For an acid-base reaction, the indicator is
chosen based on the pH at which the equivalence point is
expected to occur.
• When the indicator is in an acidic solution, the equilibrium
shifts to the left (LeChatelier's principle), and the predominant
species is HIn making the indicator show yellow color.
• In a basic solution, the equilibrium shifts favoring the forward
reaction and the predominant species is In-(blue color).
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Some Common Indicators
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Summary
After you have studied this module 6, you should be able to
• Describe the solution.
• Express concentrations of solutions in terms of molality and
mole fractions.
• Describe the factors that favor the dissolution process.
• Describe the dissolution of solids in liquids, liquids in
liquids, and gases in liquids.
• Describe how temperature and pressure affect solubility.
• Recognize and describe colloids: the Tyndall effect, the
adsorption phenomenon, hydrophilic and hydrophobic
colloids.
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40
Summary
• Recognize strong electrolytes and calculate concentrations
of their ions.
• Recognize and classify acids (strong, weak), bases (strong,
weak, insoluble), and salts (soluble, insoluble); use the
solubility guidelines.
• Understand the autoionization of water.
• Understand the pH and pOH scales and how they are used.
• Use ionization constants.
• Describe how polyprotic acids ionize in steps.
81
Summary
• Explain the common ion effect.
• Understand the solution equilibrium with precipitation.
• Recognize buffer solutions and describe their chemistry.
• Describe how to prepare a buffer solution of a specified
pH.
• Carry out calculations related to buffer solutions and their
action.
• Explain what acid–base indicators are and how they
function.
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