Acids

Strong Acids: Two types of strong acids, with
examples that you should memorize, are
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

1.The hydrohalic acids HCl, HBr, and HI
2.Oxoacids in which the number of O atoms exceeds the number
of ionizable protons by two or more, such as HNO3, H2SO4, and
HClO4; for example, in H2SO4, 4 O’s − 2H’s = 2
Weak acids: There are many more weak acids
than strong ones. Four types, with examples, are

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1.The hydrohalic acid HF
2.Acids in which H is not bonded to O or to a halogen, such as
HCN and H2S
3.Oxoacids in which the number of O atoms equals or exceeds
by one the number of ionizable protons, such as HClO, HNO2,
and H3PO4
4.Carboxylic acids (general formula RCOOH, with the ionizable
proton shown in red), such as CH3COOH and C6H5COOH
Bases

Strong bases: Water-soluble compounds
containing O2− or OH− ions are strong bases. The
cations are usually those of the most active
metals:



1.M2O or MOH, where M = Group 1A(1) metal (Li, Na, K, Rb, Cs)
2.MO or M(OH)2, where M = Group 2A(2) metal (Ca, Sr, Ba)
[MgO and Mg(OH)2 are only slightly soluble in water, but the
soluble portion dissociates completely.]
Weak bases: Many compounds with an electron-
rich nitrogen atom are weak bases (none are
Arrhenius bases). The common structural feature
is an N atom with a lone electron pair:


1.Ammonia
2.Amines R-NH2
Bronsted-Lowry Acid-Base Model





This model focuses on the nature of acids and bases
and the reactions between them.
Definitions:
Acid – proton (H+ ion) donor
Base – proton (H+ ion) acceptor
In an acid-base reaction, a proton is transferred from
an acid to a base:
HB + A- ⇌ HA + BConjugate base: the species formed when a proton is
removed from the acid.
Conjugate acid: the species formed when a proton is
added to a base.
Conjugate Acids & Bases

H2S + NH3 ⇌ HS− + NH4+



Example 1: For the equation above list the Bronsted-Lowry Acids
and Bases
 Ans: Acids - H2S, NH4+ Bases - HS-, NH3
Example 2: For the substances below list the B-L Bases
Conjugate Acid
HF
--
Ans: F-
HSO4-
--
NH4+
--
Ans: SO42Ans: NH3
Conjugate Base
Amphiprotic



A species that can either accept or donate
a proton.
For example: Water is amphiprotic:
H2O + H2O ⇌ H3O+ + OHExample 3:
What is the conjugate acid and base for

The bicarbonate ion


Ans: H2CO3, CO32-
Hydrogen sulfate ion

Ans: H2S, S2-
The Ion Product of Water:


The ionization of water can be simplified
to the ionization of one water molecule:
H2O ⇌ H+ (aq) + OH- (aq)
The ion product constant of water @
25⁰C:
Kw = [H+] [OH-] = 1.0 x 10-14
Ion Product Concentration of Water


The concentrations of H+ (H+ = H3O+) and
OH- ions are can be calculated from the Kw in
pure water @ 25⁰C
When the H+ ions equal the OH- ions the
solution is said to be a neutral solution:
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If [H+] is > 1.0 x 10-7, [OH-] < 1.0 x 10-7 solution is acidic
If [OH-] is > 1.0 x 10-7, [H+] < 1.0 x 10-7 solution is basic
These 2 quantities have an inversely
proportional relationship
pH


This is an alternative method for specifying the acidity
of a solution.
pH stands for the power of the hydrogen ion and can
be found using the following:
pH = -log[H+] = -log[H3O+]
or the same formula can be used to determine the [H+] concentration by:
[H+] = 10-pH

Determining acidic or basic:

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Acidic = pH < 7, the lower the # the more acidic the
solution
Basic = pH > 7, the higher the # the more basic the
solution
Neutral = pH = 7
pOH


A similar approach is used to determine
the hydroxide ion concentration:
pOH = -log[OH-]
Because [H+] [OH-] = 1.0 x 10-14 then the
pH and pOH are connected by the
following:
pH + pOH = 14.00
Calculating pH and pOH

Example 4: At 25 oC, calculate

the hydrogen ion concentration and pH of a
tap-water sample in which the hydroxide ion
concentration is 2.0 x 10-7.


The hydrogen and hydroxide ion
concentrations of human blood at a pH of
7.40.


Ans: 5.0 x 10-8M
Ans: 4.0 x 10-8M, 2.5 x 10-7M
The pOH of a solution in which the [H+] =
5.0[OH-].

Ans: 7.35
Strong Acids & Bases


The pH of strong acids and bases ionize
completely in water and because of this it
is relatively easy to determine the pH and
pOH.
Example 5: What is the pH of a solution
prepared by dissolving 1.00 g of barium
hydroxide per liter?

Ans: 12.07
Weak Acids & Their Equilibrium Constants
Weak acids react reversibly with water to form
H+ ions.
 Most weak acids fall into two categories:

1. Molecules containing an ionizable hydrogen
atom:
HNO2 (aq)
(aq)

+
H2O ⇌ H3O+ (aq)
+
NO2-
2. Cations. The ammonium ion acts as a weak
acid in water.
NH4+ (aq)
+
H2O ⇌ H3O+ (aq)
+
NH3
Ka
The equilibrium constant for a weak acid:
HX ⇌ H+ + XKa = [H+] [X-]
[HX]
*the smaller the Ka value the weaker the
acid.

pKa

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We sometimes use a pKa value for weak
acids: pKa = -logKa
The larger the pKa the weaker the acid.
Example 6: consider acetic acid (ka=1.8 x 10-5)
and the tetraaquazinc(II) ion (ka=3.3 x 10-10) .
 Write the equations to show why these
species are acidic.
 Which is the stronger acid?


Ans: acetic acid
What is the pKa of the tetraaquazinc ion?

Ans: 9.48
Calculating Ka using an equilibrium table…

Example 7: Aspirin is a weak organic acid
whose molecular formula maybe written
as HC9H7O4. A water solution of aspirin is
prepared by dissolving 3.60 g per liter.
The pH of the solution is found to be 2.60.
Calculate the Ka for aspirin.

Ans: 3.6 x 10-4
% Ionization


Percent Ionization allows us to measure
the percent of H+ ion that are ionized:
% ionization =
[H+] x 100
[HX]
Example 8: What percent of H+ ions from
the aspirin ionized in the previous
example?

Ans: 12%
Example 9

Nicotinic acid, HC6H4O2N (Ka = 1.4 x 10-5),
is another name for niacin, an important
member of the vitamin B group.
Determine the [H+] in a solution prepared
by dissolving 0.10 mol of nicotinic acid,
HNic, in water to form one liter of
solution.

[H+] = 1.2 x 10-3 M
Polyprotic Weak Acids
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These are acids that have more than one
ionizable hydrogen ion. These acids ionize in
steps, with a separate equilibrium constant
for each one:
H2C2O4 (aq) ⇌ H+ (aq)
+
HC2O4- (aq) ⇌ H+ (aq)
+
HC2O4C2O4-2
Ka1 = 5.9 x 10-2
Ka2 = 5.2 x 10-5
The anion formed in step one produces
another H+ ion in the next step.
The Ka becomes smaller with each successive
step:
Ka1
>
Ka2 >
Ka3
Calculating pH of a polyprotic acid

Example 10: The distilled water you use in
the laboratory is slightly acidic because of
dissolved carbon dioxide, which reacts to
form carbonic acid. Calculate the pH of a
0.0010 M solution of carbonic acid (Ka1 =
4.4 x 10-7, Ka2 = 4.7 x 10-11).

Ans: 4.68
Weak Bases & Their
Equilibrium Constants

Most weak bases, like acids, fall into two
categories:

1. Molecules, including organic compounds
known as amines:
NH3

F-
(aq)
+ H2O
⇌ NH4+ (aq)
+ OH-
(aq)
2. Anion. An anion derived from a weak acid
is itself a weak base:
(aq)
+ H2O
⇌ HF (aq)
+ OH- (aq)
Weak Bases

Example 11: Write an equation to explain
why each of the following produces a
basic water solution.



NO2Na2CO3
KHCO3
Calculation of [OH-]

Example 12: Calculate the pH of a 0.10 M
NaF (Kb = 1.4 x 10-11).

Ans: 8.08
Ka and Kb

The Ka of a weak acid and the Kb of a
weak base will equal Kw:

The stronger the
acid the weaker its
conjugate base and
vice versa for
bases.
Acid-Base Properties of Solutions

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A salt is an ionic solid containing a cation other than H+
and an anion other than OH-.
It is important to remember that anions of strong acids
and cations of strong bases are neutral in solution.
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All other cations will be acidic in nature.
All other anions will be basic in nature.
To predict whether a given salt solution will be acidic,
basic, or neutral, consider the three factors:


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1. Decide what effect, if any, the cation has on the pH of water.
2. Decide what effect, if any, the anion will have on the pH of
water.
3. Combine the 2 effects to decide the behavior of the salt.

If both cation and anion have a K value, the one with the larger
value decides acidic or basic.
Salts: Acidic, Basic, Neutral?

Example 13: Consider water solutions of the following six
salts and identify if the solution are acidic, basic, or
neutral:

Ammonium iodide


Zinc nitrate


Ans: basic
Ammonium fluoride


Ans: neutral
Sodium phosphate

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Ans: acidic
Potassium perchlorate


Ans: acidic
Ans: acidic
Ammonium hypochlorite

Ans: basic
MC #1

Which of the following ions is the
strongest Lewis acid?
(A) Na+
(B) Cl¯
(C) CH3COO¯
(D) Mg2+
(E) Al3+
MC #2

Each of the following can act as both a
Brönsted acid and a Brönsted base
EXCEPT
(A) HCO3¯
(B) H2PO4¯
(C) NH4+
(D) H2O
(E) HS¯
MC #3

Which, if any, of the following species is in the
greatest concentration in a 0.100-molar solution
of H2SO4 in water?
(A) H2SO4 molecules
(B) H3O+ ions
(C) HSO4¯ ions
(D) SO42¯ ions
(E) All species are in equilibrium and therefore
have the same concentrations.
MC #4

Which of the following reactions does NOT
proceed significantly to the right in aqueous
solutions?
(A) H3O+ + OH¯ ---> 2 H2O
(B) HCN + OH¯ ---> H2O + CN¯
(C) Cu(H2O)42+ + 4 NH3 ---> Cu(NH3)42+ + 4H2O
(D) H2SO4 + H2O ---> H3O+ + HSO4¯
(E) H2O + HSO4¯ ---> H2SO4 + OH¯
MC #5

If the acid dissociation constant, Ka, for an acid
HA is 8 x 10¯4 at 25 °C, what percent of the acid
is dissociated in a 0.50-molar solution of HA at
25 °C?
(A) 0.08%
(B) 0.2%
(C) 1%
(D) 2%
(E) 4%
MC #6

All of the following species can function as
Brönsted-Lowry bases in solution EXCEPT
(A) H2O
(B) NH3
(C) S2¯
(D) NH4+
(E) HCO3¯
MC #7

As the number of oxygen atoms increases in any
series of oxygen acids, such as HXO, HXO2,
HXO3, ...., which of the following is generally
true?
(A) The acid strength varies unpredictably.
(B) The acid strength decreases only if X is a
nonmetal.
(C) The acid strength decreases only if X is a
metal.
(D) The acid strength decreases whether X is a
nonmetal or a metal.
(E) The acid strength increases.
MC #8

A 0.20-molar solution of a weak
monoprotic acid, HA, has a pH of 3.00.
The ionization constant of this acid is
(A) 5.0 x 10¯7
(B) 2.0 x 10¯7
(C) 5.0 x 10¯6
(D) 5.0 x 10¯3
(E) 2.0 x 10¯3
MC #9

HSO4¯ + H2O <===> H3O+ + SO42–
In the equilibrium represented above, the species that
act as bases include which of the following?
I. HSO4¯
II. H2O
III. SO42–
(A) II only
(B) III only
(C) I and II
(D) I and III
(E) II and III
MC #10

H2C2O4 + 2H2O <===> 2H3O+ + C2O42¯ Oxalic
acid, H2C2O4, is a diprotic acid with K1 = 5 x 10¯2
and K2 = 5 x 10¯5. Which of the following is
equal to the equilibrium constant for the
reaction represented above?
(A) 5 x 10¯2
(B) 5 x 10¯5
(C) 2.5 x 10¯6
(D) 5 x 10¯7
(E) 2.5 x 10¯8
MC #11

A 1-molar solution of which of the
following salts has the highest pH ?
(A) NaNO3
(B) Na2CO3
(C) NH4Cl
(D) NaHSO4
(E) Na2SO4
MC #12

What is the pH of a 1.0 x 10¯2 M solution
of HCN? (Ka = 4.0 x 10¯10.)
(A) 10
(B) Between 7 and 10
(C) 7
(D) Between 4 and 7
(E) 4
MC #13

The net ionic equation for the reaction that
occurs during the titration of nitrous aicd with
sodium hydroxide is
(A) HNO2 + Na+ + OH¯ ---> NaNO2 + H2O
(B) HNO2 + NaOH ---> Na+ + NO2¯ + H2O
(C) H+ + OH¯ --->H2O
(D) HNO2 + H2O ---> NO2¯ + H3O+
(E) HNO2 + OH¯ ---> NO2¯ + H2O
MC #14

What is the H+ (aq) concentration in 0.05 M
HCN (aq)? (The Ka for HCN is 5.0 x 10¯10)
(A) 2.5 x 10¯11
(B) 2.5 x 10¯10
(C) 5.0 x 10¯10
(D) 5.0 x 10¯6
(E) 5.0 x 10¯4
MC #15

A molecule or an ion is classified as a Lewis acid
if it
(A) accepts a proton from water
(B) accepts a pair of electrons to form a bond
(C) donates a pair of electrons to form a bond
(D) donates a proton to water
(E) has resonance Lewis electron-dot structures
FRQ #1
A pure 14.85 g sample of the weak base ethylamine, C2H5NH2 , is dissolved
in enough distilled water to make 500. mL of solution.




(a) Calculate the molar concentration of the C2H5NH2 in the solution.
The aqueous ethylamine reacts with water according to the
equation below.
C2H5NH2(aq) + H2O(l)  C2H5NH3+(aq) + OH-(aq)
(b) Write the equilibrium-constant expression for the reaction
between C2H5NH2(aq) and water.
(c) Of C2H5NH2(aq) and C2H5NH3+(aq), which is present in the
solution at the higher concentration at equilibrium? Justify your
answer.
*(d) A different solution is made by mixing 500. mL of 0.500 M
C2H5NH2 with 500. mL of 0.200 M HCl. Assume that volumes are additive.
The pH of the resulting solution is found to be 10.93.




(i) Calculate the concentration of OH-(aq) in the solution.
(ii) Write the net-ionic equation that represents the reaction that occurs when
the C2H5NH2 solution is mixed with the HCl solution.
(iii) Calculate the molar concentration of the C2H5NH3+(aq) that is formed in the
reaction.
(iv) Calculate the value of Kb for C2H5NH2.
FRQ #2





HF(aq) + H2O(l)  H3O+(aq) + F–(aq)
Ka = 7.2×10–4
Hydrofluoric acid, HF, dissociates in water as represented by
the equation above.
(a) Write the equilibrium-constant expression for the
dissociation of HF in water.
(b) Calculate the molar concentration of H3O+ in a 0.40 M HF
solution.
HF(aq) reacts with NaOH(aq) according to the reaction
represented below.
HF(aq) + OH–(aq)  H2O(l) + F–(aq)
A volume of 15 mL of 0.40 M NaOH(aq) is added to 25 mL of
0.40 M HF(aq) solution. Assume that volumes are additive.
*(c) Calculate the number of moles of HF(aq) remaining in the
solution.
(d) Calculate the molar concentration of F–(aq) in the solution.
(e) Calculate the pH of the solution.
FRQ #3
HC3H5O2(aq) ↔ C3H5O2– (aq) + H+(aq) Ka = 1.34 x 10–5
Propanoic acid, HC3H5O2, ionizes in water according to the equation above.



(a) Write the equilibrium constant expression for the reaction.
(b) Calculate the pH of a 0.265 M solution of propanoic acid.
(c) A 0.496 g sample of sodium propanoate, NaC3H5O2, is added to a 50.0
mL sample of a 0.265 M solution of propanoic acid. Assuming that no
change in the volume of the solution occurs, calculate each of the following.


(i) The concentration of the propanoate ion, C3H5O2–(aq) in the solution
(ii) The concentration of the H+(aq) ion in the solution.
The methanoate ion, HCO2–(aq) reacts with water to form methanoic acid
and hydroxide ion, as shown in the following equation.
HCO2–(aq) + H2O (l) ↔ H2CO2(aq) + OH–(aq)

*(d) Given that [OH–] is 4.18 x 10–6 M in a 0.309 M solution of sodium
methanoate, calculate each of the following.



(i) The value of Kb for the methanoate ion, HCO2–(aq)
(ii) The value of Ka for methanoic acid, HCO2H
*(e) Which acid is stronger, propanoic acid or methanoic acid? Justify your
answer.
FRQ #4



Hypochlorous acid, HOCl, is a weak acid commonly
used as a bleaching agent. The acid–dissociation
constant, Ka, for the reaction represented above is
3.2×10–8.
(a) Calculate the [H+] of a 0.14–molar solution of
HOCl.
(b) Write the correctly balanced net ionic equation for
the reaction that occurs when NaOCl is dissolved in
water and calculate the numerical value of the
equilibrium constant for the reaction.
*(c) Calculate the pH of a solution made by combining
40.0 milliliters of 0.14–molar HOCl and 10.0 milliliters
of 0.56–molar NaOH.
FRQ #5




Give a brief explanation for each of the
following.
(a) For the diprotic acid H2S, the first
dissociation constant is larger than the second
dissociation constant by about 105 (K1 ~ 105 K2).
(b) In water, NaOH is a base but HOCl is an acid.
(c) HCl and HI are equally strong acids in water
but, in pure acetic acid, HI is a stronger acid
than HCl.
(d) When each is dissolved in water, HCl is a
much stronger acid than HF.