Chemistry B2A
Chapter 16
Acids and bases
Arrhenius definitions: an acid is a substance that produces H3O+ ions in aqueous solution. A
base is a substance that produces OH- ions in aqueous solution.
CH3COOH(aq) + H2O(l)
NH3(aq) + H2O(l)
CH3COO-(aq) + H3O+(aq)
CH3COOH is an acid.
NH4+(aq) + OH-(aq)
NH3 is a base.
H3O+ (hydronium ion): an H+ ion in water immediately combines with an H2O molecule (H+
cannot exist in water. Because an H+ ion is a bare proton, and a charge of +1 is too
concentrated to exist on such a tiny particle).
H+(aq) + H2O(l) → H3O+(aq)
Bronsted and Lowry definitions: an acid donates H+ (proton). A base accepts H+ (proton).
When an acid transfers a proton to a base, the acid is converted to its conjugate base. When a
base accepts a proton, it is converted to its conjugate acid. Any pair of molecules or ions that
can be interconverted by transfer of a proton is called a conjugate acid-base pair.
Note: If water is not involved in a chemical equation and H3O+ and/or OH- ions are not
produced, we should use the Bronsted and Lowry definitions.
CH3COO- + NH4+
CH3COOH + NH3
base
acid
Conjugate Conjugate
base
acid
Conjugate acid-base pair
Conjugate acid-base
Weak acid and base: an acid or a base that is only partially ionized in aqueous solution.
CH3COOH(aq) + H2O(l)
NH3(aq) + H2O(l)
CH3COO-(aq) + H3O+(aq)
NH4+(aq) + OH-(aq)
Strong acid and base: an acid or a base that ionizes completely in aqueous solution.
HCl(aq) + H2O(l)
NaOH(aq) + H2O(l)
Cl-(aq) + H3O+(aq)
Na+(aq) + OH-(aq)
Note: In general, the strength of an acid (a base) is defined by the position of its ionization
(dissociation) reaction. A strong acid is one for which the forward reaction predominates
(only one direction for the arrow). This means that almost all the original acid is dissociated
Dr. Behrang Madani
Chemistry B2A
Bakersfield college
(ionized). For a weak acid (a weak base), almost all of the molecules remain undissociated
and the reverse reaction predominates (thus the arrow pointing left is longer).
Note: A strong acid (base) is a good electrolyte and it can conduct the electricity current very
well. On the other hand, a weak acid (base) is a weak electrolyte, which means that only a few
ions are present.
Note: It is important to understand that the strength of an acid or a base is not related to its
concentration. HCl is a strong acid, whether it is concentrated or dilute, because it dissociates
completely in water to chloride ions and hydronium ions.
Note: A strong acid contains a relatively weak conjugate base.
Note: Acids are classified as Monoprotic (HCl), Diprotic (H2SO4 and H2CO3), or Triprotic
(H3PO4) depending on the number of protons (H+) each may give up. But not all hydrogen
atoms can be given up. For example, acetic acid, CH3COOH, has four hydrogens but is
Dr. Behrang Madani
Chemistry B2A
Bakersfield college
monoprotic; it gives up only one of them. This is because a hydrogen atom must be bonded to
a strongly electronegative atom, such as oxygen or a halogen, to be acidic.
Amphiprotic: a substance that can act as either an acid or a base (such as water).
Cl-(aq) + H3O+(aq)
HCl(aq) + H2O(l)
NaOH(aq) + H2O(l)
Na+(aq) + OH-(aq)
H2O is a base.
H2O is an acid.
Oxyacids: most acids are oxyacids, in which the acidic hydrogen is attached to an oxygen
atom.
Organic acids: those with a carbon atom backbone, commonly contain the carboxyl group
(-COOH). Acids of this type are usually weak. An example is acetic acid, CH3COOH.
Naming binary acids: we use the name of the anion that they produce when they dissociate.
We replace the suffix “-ic acid” instead of “-ide ion”. Then, we add it after the prefix “hydro”.
HF
HCl
anion: fluoride ion
anion: chloride ion
acid name: hydrofluoric acid
acid name: hydrochloric acid
Naming ternary acids: we use the name of the polyatomic anion that they produce when they
dissociate. For the smaller charge, we replace the suffix “-ous acid” instead of “-ite ion” and
for the larger charge, we replace the suffix “-ic acid” instead of “-ate ion”.
NO2-: nitrite ion
NO3-: nitrate ion
HNO2: nitrous acid
HNO3: nitric acid
Acid ionization constant (Ka): it shows us how strong any weak acid is. Because the
ionisation of weak acids in water are equilibria, we can only use equilibrium constants and
acid ionisation constants for these acids. We do not have the Ka for a strong acid.
HA + H2O
K=
A- + H3O+
[A-][H3O + ]
[HA] [H2O]
Ka = K[H2O] =
[A-] [H3O + ]
[HA]
- log Ka = pKa
Note: The “p” of anything is just the negative logarithm of that thing. the weaker the acid,
the smaller its Ka, but larger its pKa.
Note: Ka is always smaller than 1 and the pKa is a positive number.
Dr. Behrang Madani
Chemistry B2A
Bakersfield college
pH and pOH: pH is a scale to measure the strength of an acidic and a basic solution. This
scale is between 0 and 14. A solution is acidic if its pH is less than 7.0. A solution is basic if
its pH is greater than 7.0. A solution is neutral if its pH is equal to 7.0.
Note: there is a relationship between pH and hydronium ion concentration [H3O+] (also,
between pOH and hydroxide ion concentration [OH-]).
pH = - log [H3O+] or [H3O+] = 10-pH
pOH = - log [OH-] or [OH-] = 10-pOH
Note: we can use the following formulas as well:
pH + pOH = 14
[H3O+] × [OH-] = 1×10-14
pH meter: an instrument that can rapidly measure the pH of an aqueous solution precisely.
We dip the electrode of the pH meter into the solution whose pH is to be measured, and then
read the pH on a display. The accuracy of a pH meter depends on correct calibration.
pH paper: is made by soaking plain paper with a mixture of pH indicators. We place a drop
of solution on this paper, the paper turns a certain color. We compare the color of the paper
with the colors on a chart supplied with the paper.
pH indicator: is a substance that changes color at a certain pH. We use a pH indicator to
measure the pH of an aqueous solution (for example: Methyl orange and Litmus).
pH of strong acids: a strong acid dissociates (ionizes) completely in aqueous solution. Each
molecule of an acid dissociates into ions. Therefore, the concentration of ions including H+ is
the same as the concentration of acid.
[H+] = [HA] concentration of a strong acid
Reactions of acids and bases:
1. Reaction with metals: strong acids react with certain metals (called active metals) to
produce hydrogen gas and a salt.
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
2. Reaction with metal hydroxides: acids react with metal hydroxides to produce a salt and
water.
HCl(aq) + KOH(aq) → KCl(aq) + H2O(l)
3. Neutralization: acids and bases react with each other to produce a salt and water.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Dr. Behrang Madani
Chemistry B2A
Bakersfield college
Note: when a strong acid reacts with a strong base, the product has neither acidic nor basic
properties. We call such a solution neutral.
Titration: an analytical procedure to determine the unknown concentration of an acidic or a
basic solution. In a titration, we react a known volume of a solution of known concentration
with a know volume of a solution of unknown concentration.
B
A
MA × VA = MB × VB
Equivalence point: the point at which there is an equal amount of acid and base in a
neutralization reaction. In this situation, the solution becomes neutral ([H3O+] = [OH-] pH =
7.0).
End Point: the point at which an indicator changes color during a titration. It is convenient if
the end point and the equivalence point are the same.
Normality (N): is another unit of concentration (when dealing with acids and bases).
Normality is defined as the number of equivalents of solute per liter of solution.
Number of equivalents
Normality (N) =
Volume of solution (L)
N × V = Number of equivalents
Equivalent of an acid: is the amount of that acid that can furnish 1 mole of H+.
Equivalent of a base: is the amount of that base that can furnish 1 mole of OH-.
Equivalent weight: the equivalent weight of an acid or a base is the mass in grams of 1
equivalent of that acid or base.
Dr. Behrang Madani
Chemistry B2A
Bakersfield college
Note: during a neutralization reaction, the number of equivalents of an acid becomes equal to
the number of equivalents of a base ([H+] = [OH-]).
We use this formula if the
coefficients of an acid and a base
are different in a neutralization
reaction.
Buffer: a solution that resists change in pH when limited amounts of an acid or a base are
added to it. The most common buffers consist of approximately equal molar amounts of a
weak acid and a salt of the weak acid (or its conjugate base).
Note: The body manages to keep the pH of blood remarkably constant (when we eat the
limited amounts of an acidic or a basic food). Our body uses three buffer systems: carbonate
(HCO3-/H2CO3), phosphate (HPO42-/H2PO4-), and proteins.
How do buffers work in our body? We use the carbonate system (HCO3-/H2CO3) as an
example:
HCO3- + H3O+
H2CO3 + OH-
H2CO3 + H2O
If we eat an acidic food.
H2O + HCO3-
If we eat a basic food.
Henderson-Hasselbalch equation: it gives us a way to calculate the pH of a buffer when the
concentration of the weak acid and its conjugate are not equal [HA] ≠ [A-]:
HA + H2O
A- + H3O+
pH = pKa + log
Dr. Behrang Madani
Chemistry B2A
[A-]
[HA]
Bakersfield college