Download II. Acids and Bases

I. Measuring Concentration
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Molarity (M): the number of moles of solute
dissolved per liter of solution; also known as molar
concentration
Molality (m): the ratio of the number of moles of
solute dissolved in one kilogram of solvent; also
known as molal concentration
We will always use molarity in this class.
M = # of moles of solute / # of liters of solution
To calculate the mass of a solute when given the
volume and concentration of solution, use the
following setup:
To calculate the volume of a solution when given
concentration and mass of solute, use the following
setup:
II. Acids and Bases
1. Acidic solutions taste sour, turn blue
litmus paper pink, and produce more
hydrogen ions than hydroxide ions.
2. Basic solutions taste bitter, feel
slippery, turn pink litmus paper blue,
and produce more hydroxide ions than
hydrogen ions.
II. Acids and Bases
3. Arrhenius model: states that an acid is a
substance that contains hydrogen and
ionizes to produce hydrogen ions in
aqueous solution (ex: HCl). A base is a
substance that contains a hydroxide group
and dissociates to produce a hydroxide ion
in aqueous solution (ex: NaOH).
4. HCl(g)  H+(aq) + Cl-(aq)
5. NaOH(s)  Na+(aq) + OH-(aq)
II. Acids and Bases
6. Although the Arrhenius model is useful
in explaining many acidic and basic
solutions, it has some shortcomings.
7. For example, ammonia (NH3) does not
contain a hydroxide group, yet it
produces hydroxide ions in solution
and is a well known base.
II. Acids and Bases
8. Bronsted-Lowry model: states that an acid
is a hydrogen-ion donor and a base is a
hydrogen-ion acceptor
9. For example, when a molecule of acid, HX,
dissolves in water, it donates an H+ ion to a
water molecule. The water molecule acts as
a base and accepts the H+ ion.
HX(aq) + H2O(l) <--> H3O+(aq) + X-(aq)
10. In the example above, the water molecule
becomes an acid (H3O+) by accepting the
H+ ion.
II. Acids and Bases
11. In the example above, the acid (HX) becomes
a base because it can now accept a positive
hydrogen ion.
12. Conjugate acid: the species produced when
a base accepts a hydrogen ion from an acid
13. Conjugate base: the species that results
when an acid donates a hydrogen ion to a
base
14. Amphoteric: water and other substances that
can act as both acids and bases
III. Naming Acids and Bases
1. Bases are named like other ionic
compounds (e.g., NaOH = sodium
hydroxide)
2. Two types of acids: binary and oxyacids
3. Binary acid contains hydrogen and one
other element.
hydro + nonmetal root + ic acid
(e.g., HCl = hydrochloric acid)
III. Naming Acids and Bases
4. Oxyacid contains hydrogen and an
oxyanion (polyatomic ion with oxygen
in it)
a. Name the polyatomic ion.
b. Replace “ate” with “ic”, “ite” with “ous.”
c. Change nonmetal root for pronunciation.
d. Add “acid” to the name.
(e.g., H2SO4 = sulfuric acid)
IV. Acid/Base Strength
1. Strong acids/bases: acids/bases that
dissociate (ionize) completely in aqueous
solutions
2. Weak acids/bases: acids/bases that
dissociate (ionize) only partially in aqueous
solutions
3. Reactions involving weak acids/bases reach
a state of equilibrium where the forward and
reverse reactions occur at equal rates. This
is represented by a double arrow.
IV. Acid/Base Strength
4. The equilibrium constant, Keq, provides
a measure of the degree of ionization
of an acid/base.
5. For the following reaction:
HCN(aq) + H2O(l) <-> H3O+(aq) + CN-(aq)
Keq =
[H3O+][CN-]
[HCN][H2O]
IV. Acid/Base Strength
Keq [H2O] = Ka =
[H3O+][CN-]
[HCN]
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The concentration of liquid water in the
denominator of the expression is considered to be
constant in dilute aqueous solutions, so it can be
combined with Keq to give a new equilibrium
constant, Ka.
Acid ionization constant (Ka): the value of the
equilibrium constant expression for the ionization of
a weak acid
IV. Acid/Base Strength
8. The value of Ka indicates whether
reactants or products are favored at
equilibrium.
9. The weakest acids have the smallest
Ka values because their solutions have
the lowest concentrations of ions and
the highest concentrations of unionized acid molecules.
IV. Acid/Base Strength
10. Base ionization constant (Kb): the
value of the equilibrium constant
expression for the ionization of a base
11. CH3NH2 (aq) + H2O(l) <->
CH3NH3+(aq) + OH-(aq)
12. Kb =
[CH3NH3+] [OH-]
[CH3NH2]
V. pH and pOH
1. Ion Product Constant for Water (Kw): the
value of the equilibrium constant expression
for the self-ionization of water
H2O(l) <-> H+(aq) + OH-(aq)
2. Keq = [H+][OH-]
[H2O]
3. Keq [H2O] = [H+][OH-]
4. Kw = [H+][OH-]
5. In pure water at 298 K, [H+] and [OH-] are
both equal to 1.0 x 10-7 M.
V. pH and pOH
Therefore, the value of Kw is 1.0 x 10-14.
This is also the product of [H+] and [OH-] for other
solutions.
8. Because concentrations of H+ ions are often small
numbers expressed in scientific notation, chemists
adopted an easier way to express H+ ion
concentration using a pH scale based on common
logarithms.
9. pH: the negative logarithm of the hydrogen ion
concentration of a solution
10. pH = -log [H+]
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V. pH and pOH
11. If [H+] equals 1.0 x 10-7 M, pH is calculated
as follows:
pH = -log (1.0 x 10-7)
pH = -(log 1.0 + log 10-7)
pH = -[0 + (-7)] = 7.00
12. pH > 7 is basic/alkaline
pH = 7 is neutral
pH < 7 is acidic
V. pH and pOH
13. Sometimes chemists find it convenient
to express the basicity, or alkalinity, of
a solution on a pOH scale that mirrors
the relationship between pH and [H+].
14. pOH: the negative logarithm of the
hydroxide ion concentration of a
solution
15. pOH = -log [OH-]
V. pH and pOH
16. pOH > 7 is acidic
pOH = 7 is neutral
pOH < 7 is basic
17. pH + pOH = 14.00
VI. Neutralization
1. Neutralization Reaction: a reaction in
which an acid and a base react in aqueous
solution to produce a salt and water.
2. Salt: an ionic compound made up of a
cation from a base and an anion from an
acid.
3. Neutralization is a type of double
replacement reaction.
VI. Neutralization
4. Titration: a method for determining the
concentration of a solution by reacting a
known volume of the solution with a solution
of known concentration.
5. If you wished to find the concentration of an
acid solution, you would titrate the acid
solution with a solution of base of known
concentration.
6. The solution of known concentration is
called the standard solution.
VI. Neutralization
7. Measured volumes of the standard solution
are added slowly and mixed into the
solution in the beaker.
8. This process continues until the reaction
reaches the stoichiometric point
(equivalence point), which is when the
moles of H+ ion from the acid equal the
moles of OH- ion from the base.
9. A buret is used to dispense the standard
solution.
VI. Neutralization
10. For convenience, chemists often use a chemical
dye rather than a pH meter to detect the
equivalence point of an acid-base titration.
11. Acid-base indicators: chemical dyes whose colors
are affected by acidic and basic solutions
12. End point: the point at which the indicator used in
a titration changes color
13. The equation used in a titration is as follows:
MAVA = MBVB
M = molarity
A = acid
V = volume
B = base
VI. Neutralization
14. Buffer: a solution that resists change in pH
when limited amounts of acid or base are
added
15. A buffer is a mixture of a weak acid and its
conjugate base or a weak base and its
conjugate acid.
16. The mixture of ions and molecules in a
buffer solution resist changes in pH by
reacting with any hydrogen ions or
hydroxide ions added to the buffered
solution.
VI. Neutralization
17. Buffer capacity: the amount of acid or
base a buffer solution can absorb
without a significant change in pH