Acid-Base Notes
Water has many unique chemical and physical properties. Possibly one of the most
important is its ability to dissolve other substances to form solutions. Solutions are
homogeneous mixtures of two or more substances. The solvent (usually the substance
present in the greatest quantity) causes the other substance(s), the solute(s), to dissolve
and enter into the solution.
In an aqueous solution, water is the solvent, or dissolving medium. The symbol “(aq)” in a
chemical equation indicates that the substance is dissolved in water.
CaCO3 (s) + H2O (l) + CO2 (aq) à Ca(HCO3)2 (aq)
Some substances dissociate (split apart) in solution to form ions. Since such solutions can
conduct an electric current, the dissolved substance is classified as an electrolyte. These
substances consist of ions that are held together by the electrostatic forces of attraction
between their opposite charges -- ionic bonding (see Chapter 8).
H 2O
NaCl (s) 
→ Na+(aq) + Cl–(aq)
Nonelectrolytes do not dissociate in aqueous solution and thus their solutions do not
conduct electricity. These substances consist of molecules where the atoms are covalently
bonded (see Chapter 8).
H 2O
C12H22O11 (s) 
→ C12H22O11 (aq)
Due to the strength of their ionic bonds, some ionic substances dissociate almost
completely in aqueous solution. These substances, such as hydrobromic acid, are
classified as strong electrolytes.
HBr (aq) à H+(aq) + Br–(aq)
Some molecular substances, such as acetic acid, partially dissociate in aqueous solution
and are classified as weak electrolytes.
+
–
CH3COOH (aq) ←
→ H (aq) + CH3COO (aq)
The double arrows indicate that the reaction occurs in both directions. The longer arrow
(pointing to the left) indicates that at chemical equilibrium, the concentration of the
undissociated acetic acid is greater than that of its dissociated ions.
NOTE:
Please do not confuse concentration (mole per liter) with strength (degree of
dissociation).
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Ionic Equations
During aqueous chemical reactions involving ionic substances, some ions are not involved
in the actual reaction. These uninvolved ions are known as spectator ions. When writing
the chemical equation to represent a reaction involving ions, we can choose to represent
the reaction in one of three manners.
The molecular equation shows all of the involved species and their phases:
AgNO3 (aq) + NaCl (aq) à AgCl (s) + NaNO3 (aq)
The complete ionic equation shows all of the involved species with the strong electrolytes
dissociated into aqueous ions. Solids, liquids, gases, and aqueous molecular species
retain their molecular formulas:
Ag+(aq) + NO3–(aq) + Na+(aq) + Cl–(aq) à AgCl (s) + Na+(aq) ) + NO3–(aq)
In the net ionic equation, the spectator ions are eliminated to show only the species that are
actually involved in the chemical reaction:
Ag+(aq) + Cl–(aq) à AgCl (s)
Acids and Bases
Simple properties such as the ability to change the color of certain dyes (indicators) and
taste (acids are sour and bases are bitter) were first used to identify acids and bases.
Svante Arrhenius (Sweden, 1859-1927) proposed the first successful acid-base definition in
1884. According to Arrhenius, an acid dissociates in water to form H3O+(aq) (often simplified
as H+(aq)) and a base dissociates in water to form OH–(aq). Strong acids dissociate
completely in an aqueous solution to produce H3O+(aq) and an anion. Strong bases
dissociate completely in an aqueous solution to produce OH–(aq) and a cation.
The neutralization reaction between the acid’s H+(aq) and the base’s OH–(aq) forms water:
: H+(aq) + OH–(aq) à H2O(l)
The neutralization of an acid and base has a ∆H o = – 55.90
kJ
.
mol
A salt is formed when the cation from the base combines with the anion from the acid.
Soluble salts can be recovered by evaporating the water from the solution.
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The reaction of nitric acid with potassium hydroxide (the base) produces water and
potassium nitrate (the salt):
HNO3 (aq) + KOH (aq) à H2O (l) + KNO3 (aq)
Some acid-base reactions produce salts that decompose to form water and a gas:
2 HNO3 (aq) + Na2CO3 (aq) à 2 NaNO3 (aq) + H2O (l) + CO2 (g)
Most acids and bases are weak… they are only partially dissociated in water and exists in
equilibrium with their respective ions:
HC2H3O2 (aq) ↔ H+(aq) + C2H3O2–(aq)
Strong acids and bases dissociate almost completely in water. Since there are so few
strong acids and bases, it is easier to remember which common acids and base are strong
rather than trying to learn rules to determine whether or not an acid or base is weak or
strong.
Common Strong Acids
Hydrochloric
Hydrobromic
Hydroiodic
Chloric
Perchloric
Nitric
Sulfuric
HCl
HBr
HI
HClO3
HClO4
HNO3
H2SO4**
Common Strong Bases
Lithium hydroxide
Sodium hydroxide
Potassium hydroxide
Rubidium hydroxide
Cesium hydroxide
Calcium hydroxide
Strontium hydroxide
Barium hydroxide
LiOH
NaOH
KOH
RbOH
CsOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
**
Note: H2SO4’s first ionization is strong while its second ionization is weak
H2SO4 (aq) à H+(aq) + HSO4–(aq)
+
2–
HSO4–(aq) ←
→ H (aq) + SO4 (aq)
Most acids are monoprotic – they have one acidic proton. Sulfuric acid, H2SO4(aq), is
diprotic – it has two acidic protons. Phosphoric acid, H3PO4(aq), is triprotic – it has three
acidic protons.
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Most acids have the acidic proton attached to an oxygen atom – these acids are known as
oxyacids. Organic acids contain the carboxyl group:
O
C
O H
The Arrhenius acid-base definition does not deal with nonaqueous solutions or with
substances that exhibit acid-base behavior that do not fit into his original definitions.
Independently, Johannes BrØnsted (Denmark, 1879-1947) and Thomas Lowry (England,
1874-1936) proposed that an acid is any substance that acts as a proton-donor. A
BrØnsted-Lowry base is any substance that acts as a proton-acceptor.
When dissolved in benzene, hydrogen chloride and ammonia do not dissociate into ions.
HCl(benzene) + NH3 (benzene) à NH4Cl(s).
Hydrogen chloride gas reacts with ammonia gas to form solid ammonium chloride:
HCl(g) + NH3(g) à NH4Cl(s).
In both instances, the HCl acts as the proton-donor (acid) and NH3 acts as the protonacceptor (base).
In 1923, G. N. Lewis (United States, 1875-1946) introduced an expanded acid-base theory
involving electron pairs. A Lewis acid is an electron-pair acceptor. A Lewis base is an
electron-pair donor.
In the reaction between BF3 and NH3, the BF3 is the acid and NH3 is the base:
F
F
B
F
F
H
+
N
H
F
H
H
B
F
N
H
H
The BF3 uses the lone-pair of electrons on the N in NH3 to complete the octet for the B in
BF3. Since both electrons in the covalent bond between the N and B atoms came from the
same atom (N), this is an example of a coordinate covalent bond.
Amphiprotic (a.k.a. amphoteric) substances act as either an acid or as a base depending
on the other reactant.
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The hydrogen carbonate (bicarbonate) ion, HCO3– is an example of an amphiprotic
substance:
In the reaction: HCO3–(aq) + HF(aq) ↔ H2CO3 (aq) + F–(aq), the HCO3–(aq) is the base since
it accepts an H+ from the HF to form H2CO3.
In the reaction: HCO3–(aq) + OH–(aq) ↔ CO32–(aq) + H2O(l), the HCO3–(aq) is the acid since
it donates an H+ to the OH– to form H2O.
In the first example: HCO3–(aq) + HF(aq) ↔ H2CO3 (aq) + F–(aq),
The HF is the acid and F– is its conjugate-base (formed when the acid molecule donates
an H+).
The H2CO3 is the conjugate-acid (formed when the base accepts an H+) of the base,
HCO3–.
If the base has a much greater affinity for the proton (H+) than does the conjugate-base,
the equilibrium will shift to the right (most of the acid is dissolved). If the conjugate-base
has a greater affinity for the proton (H+) than does the base, the equilibrium will shift to the
left (most of the acid would be undissociated).
For the first example the equilibrium expression would be:
[H 2CO3 ][F – ]
where [ ] represents the molar concentration and Ka is the acid
[HCO3 – ]
dissociation constant for that the acid HCO3–. Strong acids have large Ka values while
weak acids have very small Ka values.
Ka =
Water, H2O, is the most common amphoteric substance. Water undergoes
autoionization to produce hydroxide and hydronium ions:
+
H
H
O
H
+
O
H
↔
[
O
H
H
H
] +[
_
0 H
]
At 25 oC, the concentrations of the H+ and OH– ions in “pure” neutral water are 10–7 M
each. The [H+] and [OH–] concentrations change with temperature changes, but chemists
typically use the concentrations for 25 oC.
Acidic solutions have [H+] > 10–7 M and basic solutions have [H+] < 10–7 M.
The ion-product constant for water, Kw = [H+][OH–]. At 25 oC, the Kw = 10–14.
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Due to the extremely small H+ concentrations associated with acidic and basic solutions, a
special measurement scale is used. The pH scale is logarithmic.
pH = – log [ H + ]
pOH = – log [OH – ]
[H + ] = 10 – pH
pH + pOH = 14
[OH – ] = 10 – pOH
For aqueous solutions: pH < 7 (acidic), pH = 7 (neutral), and pH > 7 (basic).
Please note that concentrated strong acids may have a negative pH and concentrated
strong bases may have a pH value greater than 14.
Example:
What is the pH for a solution this has an H+ concentration of 0.005?
pH = – log [ H + ] = – log [0.005] = – (–2.3) = 2.3
pH can be measured directly using meters (instruments) or indirectly using color changes
for selected indicators such as phenolphthalein, litmus, methyl red, etc., during titrations.
Salt solutions are formed by the reaction of an acid and a base. The salt is formed by the
cation from the base and the anion from the acid. The pH of the resulting solutions
depends on the pH values for the starting acid and base solutions.
Acid
strong
strong
weak
Base
strong
weak
strong
weak
weak
pH of Salt Solution
7 (neutral)
< 7 (acidic)
> 7 (basic)
depends on the relative
Ka and Kb values
Common-ion effect – the addition of a solute that produces aqueous ions that take part in
the equilibrium. The added solute ions shift the equilibrium to relieve the stress on the
system.
Buffers are solutions that resist changes in pH. The buffered solution contains a
“common-ion” that, while present, neutralizes the added acid or base.
For a buffered solution, pH = pKa + log
[base]
[acid]
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Acid-Base Titrations
The acid (or base) concentration for an unknown solution can be determined by titration.
During a titration, a standardized base (or acid) is added until the equivalence point is
reached – a stoichiometric amount of reactant. The equivalence point is indicated by the
color change for a weak-organic acid (indicator). As an alternative, a digit pH meter that
measures electrical conductivity may be used to determine the equivalence point.
For example: due to its high purity, the weak acid primary standard potassium hydrogen
phthalate (KHC8H4O4, MM = 204.2212) can be used to standardize a sodium hydroxide
solution. Phenolphthalein would be used as the indicator. A known mass of potassium
hydrogen phthalate would be dissolved in distilled water and a few drops of
phenolphthalein would be added.
The choice of indicator is selected based on the pH range in which the equivalence point
occurs for the particular acid and base being used. The equivalence point for a strong-acid
reacting with a strong-base would have a pH around 7 where litmus or bromthymol blue.
Methyl orange or methyl red might be used for the reaction of a strong-acid and a weakbase where the equivalence point would be around 4 or 5. Phenolphthalein would be the
indicator of choice for the reaction of a strong-base with a weak-acid where the equivalence
point would be above 7.
The concentration unit molarity (M) represents the number of moles of solute per liter of
solution. The product of molarity (M) and volume (expressed in liters) is moles (mol).
During a titration the number of moles of H+ should equal the moles of OH–. Thus, for
monoprotic acids and monobasic bases the relationship: MaVa = MbVb. A more specific
variation for acids and bases that contain multiple H+ or OH– ions would be:
 number of ionizible H + ions 
 number of ionizible OH – ions 

 M acid Vacid = 
 M baseVbase
acid molecule
base molecule




Solution Preparation
The solutions one uses in a chemistry laboratory typically are prepared by diluting the
“stock” reagent to the desired concentration by adding distilled water. If the reagent is a
solid, it will need to be dissolved first.
Example 1: What mass of sodium hydroxide is needed to prepare 1.000 L of a 0.1000 M
solution?
1.000 L x 0.1000
mol NaOH
39.998 g NaOH
x
= 4.000 g NaOH
L
1 mol NaOH
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Example 2: What volume of concentrated (11.7 M) hydrochloric acid is needed to prepare
1.000 L of 0.1000 M
1.000 L x 0.1000
mol HCl
1 L
x
= 0.008547 L
L
11.7 mol HCl
As you can see, a very small volume of concentrated HCl is needed. Since it is extremely
difficult (if not impossible) to measure exactly 4.000 g of NaOH pellets or 8.547 mL of
concentrated acid, we prepare the solutions as closely as possible and then standardize
the solutions.
Standardization
Once the solutions have been prepared, their concentrations can be compared
experimentally to a known number of moles of acid/base. As mentioned previously, the
weak acid potassium hydrogen phthalate has a very high purity so a known mass can be
measured out, dissolved in distilled water, and then titrated with the NaOH.
Example 3: What is the OH– concentration for a solution if 37.58 mL of NaOH solution
was required to titrate 0.675 g of potassium hydrogen phthlate?
0.675 g KHC8 H 4O4 x
M =
1 mol KHC8 H 4O4
1 mol NaOH
x
204.2212 g KHC8 H 4O4 1 mol KHC8 H 4O4
1L
37.58 mL x
1000 mL
M = 8.80 x 10 –2 M NaOH
Example 4: What is the H+ concentration for a hydrochloric acid solution if the titration of a
25.00 mL sample required 41.04 mL of the standardized NaOH solution?
8.80 x 10 –2 M NaOH x 41.04 mL NaOH x
M HCl =
25.00 mL HCl
1 mol HCl
1 mol NaOH
=
1.44 x 10 –1 M HCl
NOTE:
Since volumes were given for both the HCl and the NaOH, it was not
necessary to convert the mL to L … they cancel out either way.
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