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Chapter 10 – Acids and Bases
Defining Acids and Bases
A single hydrogen atom can exist as any of three species,
depending on how many electrons it has:
+
H
hydrogen cation
(proton)
H
hydrogen radical
H
hydride
Hydrogen atoms are most stable when they have ____
electrons and a charge of ____. Since none of these three
species meet both requirements, they are all reactive.
H
+
+
H .. +
- . ..
.O
.. H
..
H O
.. H
..
H O
.. H
- . ..
H H + .O
.. H
Note that proton is being given an electron pair while
hydride is giving away an electron pair.
This is the primary difference between an acid and a base:
An acid is given an electron pair.
A base gives one away.
This is the broadest definition of acids and bases (Lewis
definition) and it is true of all acids and bases.
Thus, H+ is _____________ while H- is _____________.
Because H+ is an acid, we can reasonably conclude that
any compound that can make H+ is also an acid. This
leads to the Brønsted-Lowry definition of acids and bases:
An acid is a proton donor (gives up H+).
A base is a proton acceptor (receives H+).
Note that H- accepted H+ from the water in the reaction on
the previous page.
When a base accepts a proton, it is said to be
____________________.
When an acid gives up a proton, it is said to be
____________________.
The Arrhenius definition of acids and bases is almost
identical to the Brønsted-Lowry definition. The only
difference is that it looks at their behaviour in water:
An acid produces H+ when dissolved in water.
A base produces HO- when dissolved in water.
Note that when H- accepted H+ from the water, it
produced HO-. HO- is always the product when H+ is
accepted from H2O.
Naming Acids and Bases
Acids are named based on the ion produced when all
potential H+ have been removed. An H bonded to a C is
not seen as a potential H+. To be removed as H+, the
hydrogen atom has to start with a δ+ charge.
Ions that end in ‘ate’: replace the ate with ic acid.
Ions that end in ‘ite’: replace the ite with ous acid.
Ions that end in ‘ide’: add hydro to the front of the name
and replace the ide with ic acid.
e.g. chlorate (ClO3-) comes from chloric acid (HClO3)
chlorite (ClO2-) comes from chlorous acid (HClO2)
chloride (Cl-) comes from hydrochloric acid (HCl)
Bases are named according to the rules in Chapter 4.
Classify the following compounds as acids or bases, and
name them:
acid/base
_____________________
(a) H2CO3
acid/base
_____________________
(b) H3PO3
(c) LiOH
acid/base
_____________________
(d) HI
acid/base
_____________________
(e) NaCN
acid/base
_____________________
For each of compounds (c) to (e), write a reaction
equation showing its protonation or deprotonation.
Conjugate Acids and Bases
Since an acid donates H+ and a base accepts H+, it is
logical that acids and bases like to react with each other:
These reactions are often reversible:
Notice that in the reverse reaction, the compound (I-) that
we made from the acid (HI) acts as a base. We therefore
refer to I- as the conjugate base of HI.
Similarly, the compound (HOCl) that we made from the
base (-OCl) acts as an acid. We therefore refer to HOCl
as the conjugate acid of –OCl.
We can write a generalized reaction equation to show the
relationship between acids, bases, conjugate acids and
conjugate bases:
H A
acid
+
B
A
base
conjugate base
+
H B
conjugate acid
In each of the following reactions, identify the acid, the
base, the conjugate acid and the conjugate base:
NO3-(aq) + C5H5NH+(aq)
(a) HNO3(aq) + C5H5N(aq)
(b)
NH3(aq) + H2O(l)
NH4+(aq) + HO-(aq)
(c)
CO32-(aq) + H2O(l)
HCO3-(aq) + HO-(aq)
Strength of Acids and Bases
The strength of an acid is a measure of its ability to give
up H+. A strong acid makes a stable enough conjugate
base that, when it is dissolved in water, every molecule of
acid gives up a proton:
HClO4(aq)
+
H(aq) +
-
ClO4(aq)
A weak acid has a less stable conjugate base. While
some molecules of acid give up a proton, others do not:
HCN(aq)
There are six strong acids:
+
H(aq) +
-
CN(aq)
The strength of a base is a measure of its ability to accept
H+. A strong base makes a stable enough conjugate acid
that, when it is dissolved in water, every molecule of base
produces a hydroxide ion:
NaOH(aq)
+
Na(aq) +
-
OH(aq)
A weak base has a less stable conjugate acid. While
some molecules of base produce hydroxide, others do not:
NH3(aq)
+ H2O(aq)
+
NH4(aq) +
-
OH(aq)
For the purposes of this course, the hydroxides of the
alkali metals (Group 1) will be the only strong bases used:
The pH Scale
The acidity (or basicity) of a solution is reported as pH:
pH = - log[H+]
where [H+] is the concentration of H+ (or H3O+) in mol/L.
Since pH is a logarithmic scale, changing [H+] by a factor
of ten (×101) changes the pH by 1. Changing [H+] by a
factor of a hundred (×102) changes the pH by 2.
Note that the negative sign means that increasing [H+]
decreases pH. Thus, acidic solutions have lower pH
values than basic solutions:
• Solutions with pH less than 7 are acidic.
• Solutions with pH greater than 7 are basic.
Pure water is exactly neutral and has a pH of 7; however,
most water is not pure. Tap water, for example, often has
a pH of about 5 because it has absorbed CO2 from the air.
What does this tell us about CO2?
Because we know that every molecule of a strong acid
gives up one proton (H+), we can calculate the pH of a
solution of strong acid if we know its concentration.
Calculate the pH of:
(a) 0.04 mol/L HBr(aq)
(b) 0.00076 mol/L HNO3(aq)
(c) concentrated HCl(aq) (12 mol/L)
Calculate the concentration of a solution of HI(aq) with a
pH of 6.5.
***You will often see mol/L abbreviated as M.***
1 M = 1 mol/L
Reactions of Acids and Bases
A reaction between an acid and a base produces water and
a salt:
A strong acid will react completely with any base.
A strong base will react completely with any acid.
A reaction between a weak acid and a weak base will not
go to completion unless there is a driving force (e.g.
making a gas):
Write a balanced chemical equation for each of the
following acid-base reactions:
(a) KOH(aq)
+
H2SO4(aq) →
(b) H3PO4(aq) +
LiOH(aq)
→
(c) NH4OH(aq) +
HCl(aq)
→
(d) K2CO3(aq) +
HBr(aq)
→
Note that H2CO3(aq) tends to break into H2O(l) and CO2(g).
A strong acid will also react with a metal to give a salt
and hydrogen gas:
These reactions tend to be very exothermic!
Think back to Chapter 6. How would you classify these
acid-metal reactions?
Important Concepts from Chapter 10
• the difference between H+, H- and H2
• defining acids and bases:
o Lewis
o Brønsted-Lowry
o Arrhenius
• protonation vs. deprotonation
• conjugate acids and conjugate bases
• strong acids vs. weak acids
• strong bases vs. weak bases
• pH
o qualitative (low pH = acid; high pH = base)
o calculating pH from concentration
o calculating concentration from pH
• acid-base reactions