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8
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Vinegar is simply an aqueous solution of acetric acid
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Until relatively recent times, the
principal source of bases or alkalis has been wood ashes
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The properties that are most commonly associated with acids are: their sour taste, the
prickling sensation they produce on the skin, their ability to react with most metals to
form hydrogen gas (H2), and their ability to dissolve limestone and other carbonate
minerals form carbon dioxide (CO2)
...

All bases also share several common properties: They have a bitter taste and their
solutions feel slippery like soapy water (basic properties commonly found in soap and
other household cleaners)
...

Both acids and bases have the ability to affect the colour of certain naturally occurring
plant constituents
...

ARRHENIUS DEFINITION OF ACIDS AND BASES
Arrhenius proposed that:
+

Acids are substances that dissociate to produce hydrogen ions (H )
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Chemical and electrical activity of the acids and bases are directly related to the extent
to which these H+ and OH– are present in the solution
...
A strong
electrolyte is completely dissociated into ions in a water solution; a weak electrolyte is
only partially ionized
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(2)

H+ + Cl– ,

→

HCl

Kc very large

When sulfuric acid dissolves in water, it is almost completely ionized (hardly any

reverse reaction)
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Concentration of H+ is very small
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8x10 -5
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[CH 3 COOH]

NaOH, KOH and barium hydroxide Ba(OH)2 are strong bases, almost 100%

ionized in solution
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Most of ammonia remains as NH3
molecule dissolved in water
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8x10 –5
[NH 3 ]

Neutralization
This refers to the reaction of an acid and a base yielding water and a salt
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For reactions between strong acid and strong base: For example,
HCl + NaOH → H2O + NaCl
Since HCl

and NaOH are almost 100% ionized, the above can be written in hydrated

ionic form,

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H+ (aq) + Cl– (aq) + Na+(aq) + OH–(aq) → H2O + Na+(aq) + Cl– (aq) …(i)
Equation (i) is more simply written as net ionic equation, as
H+ (aq) + OH–(aq) → H2O

…(ii)

For reactions between strong acids and bases the net ionic equation is always given by
Eqn (ii)
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1 kJ/mol H2O
formed)
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Limitations of Arrhenius’s Theory
One of the most glaring limitations of this theory is in its treatment of the weak base
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ammonia, NH3
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Where is the OH in NH3?

BRONSTED- LOWRY THEORY OF ACIDS AND BASES
According to Bronsted- Lowry, an acid is a proton (H+) donor and a base is a proton
(H+) acceptor
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(1)

base

H2O acts as an acid
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As a result
of this transfer, the ions NH4+ and OH- are formed
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(2)

base

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Here, NH4+ donates a proton (H+) so it acts as an acid, and OH– accepts a proton acting
like a base
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HCl/ Cl– , and H3O+ / H2O are

known as conjugate acid-base pair
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SELF IONIZATION OF WATER
Even in pure water, there is a very low concentration of ions, detectable by precise
measurements of electrical conductivity
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According to Arrhenius theory, we
would write the dissociation reaction as:
H2O

H+ + OH–

Kc is very small
...
That is, one water molecule acts as an acid and the
+

other as a base
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The reaction is reversible, and in the reverse
+

-

reaction, H3O donates a proton to OH
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K = [H+ ][ OH– ]
At 25oC in pure water: [H+ ] = [ OH– ] = 1
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At 25oC
14

Kw = [H+ ][ OH– ] = 1
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At
14

13

60oC, Kw = 9
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5 x10 –
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Calculating ion concentrations in an aqueous solution of a strong acid
In acid solutions (eg
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Most of H+ coming from acid and only a negligible amount from H2O
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Example 1:

Estimation of [ OH– ] in a 0
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Solution:

A 0
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10 M H+ (100% dissociation)
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0 x10 –7 M, which is negligible compared to H+ from HCl
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10 M + 1
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10 M
Kw = [H+ ][ OH– ] = 1
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0 x10 –14 / 0
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0 x10 –13 M

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Calculating ion concentrations in an aqueous solution of a strong base
In basic solutions, the concentration of OH– >> [H+ ]
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But water dissociation is the only source of
H+
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01 M NaOH
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[ OH– ] ≈ 0
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[H+ ] = Kw / [ OH– ] = 1
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010 = 1
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Definition of pH:

pH = – log [H+]
[H+ ] = 10 –pH

Or,
Examples,

for water, [H+ ] = 1
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10 M HCl, [H+ ] = 0
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010 M NaOH, [H ] = 1
...

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 As a solution becomes more basic, its [OH ] increases, its pOH decreases and
its pH increases
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00 is neutral; with pH < 7
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00, a solution is basic
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Human blood maintains a pH between 7
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45, although it
transports chemical and waste products, many of which by themselves are acids and
bases
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A buffer
(buffered) solution is one whose pH changes only very slightly upon the addition of small
amounts of either an acid or a base
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Common buffer solutions are mixtures containing a
 Weak acid and its conjugate base (one of its salts) or a
 Weak base and its conjugate acid (one of its salts)

Example
CH3COOH and CH3COO–, a weak acid and a salt of that weak acid, such as
CH3COONa, that provides the conjugate base
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8x10 -5
[CH 3 COOH]
Or,

[H+] = K c

[CH 3 COOH ]
= 1
...
In this case,
[H+] = 1
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7
If small amounts of H+ are added to this buffer solution, the reaction (1) is driven to the
left to restore equilibrium, consuming most of the added H+
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The ratio
[CH3COOH] / [CH3COO –] changes only slightly, thus the [H+] , and therefore the pH
remains unchanged
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Other examples of buffer are NH3 / NH4+ (coming from NH4Cl) system
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8x10 –5
[NH 3 ]
Or,

[OH-] = Kc [NH3]/[NH4+] = Kc , when [NH3] and [NH4+] are large and equal
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8x 10-5 = 5x 10 –10 , and pH = 9
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