Acid-Base Definitions
Type
Arrhenius
Brønsted-Lowry
Lewis
Acid
+
H producer
+
H donor
Base
–
OH producer
+
H acceptor
electron-pair acceptor electron-pair donor
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A chemist named Svante Arrhenius came up with a way to define acids and
bases in 1887
Arrhenius Acid - A substance that releases hydrogen ions (H+) in a solution.
Arrhenius Base - a substance that releases hydroxide ions (OH-) in a solution.
Chemists use the word "dissociated" to describe the breakup of a compound.
Brønsted-Lowry - These two chemists from Denmark and England looked at acids
as donors and bases as acceptors.
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What were they donating and accepting? Hydrogen ions.
Brønsted-Lowry acid: a hydrogen ion (H+) donor.
Brønsted-Lowry base: an hydrogen ion acceptor.
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Scientists use something called the pH scale to measure how acidic
or basic a liquid is.
Low pH = acidic = lots of H+
High pH = basic = lots of OH-
When ammonia dissolves in water, hydrogen ions are transferred from water to
ammonia to form ammonium ions and hydroxide ions.
Is there an acid or a base?
Important Note:
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In an aqueous solution, hydrogen ions are not actually present. Instead, the
hydrogen ions are joined to water molecules as hydronium ions.
A hydronium ion (H3O+) is the ion that forms when a water molecule gains a
hydrogen ion.
HCl + H--2O → H-+ + Cl● Acid
CH3COOH + H2O → CH3COO- + H● Acid
NaOH + H2O → Na+ + OH● Base
NH3 + H2O → NH4+ + OH● Base
A chemist named Lewis offered a third way to look at acids and bases. Instead of
looking at hydrogen ions, he looked at pairs of electrons.
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Lewis Acid: accept pairs of electrons.
Lewis Base: donate pairs of electrons.
A conjugate acid is the ion or molecule formed when a base
gains a hydrogen ion.
•NH4+ is the conjugate acid of the base NH3.
A conjugate base is the ion or molecule that remains after an
acid loses a hydrogen ion.
•OH– is the conjugate base of the acid H2O.
•The ammonia molecule and the ammonium ion are a
conjugate acid-base pair.
•The water molecule and the hydroxide ion are also a
conjugate acid-base pair.
Amphoteric - A substance that can act as either an acid or a base. Example: Water.
–In the reaction with hydrochloric acid, water accepts a proton and is therefore a
base.
–In the reaction with ammonia, water donates a proton and is therefore an acid.
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Scientists use something called the pH scale to measure how acidic
or basic a liquid is.
Low pH = acidic = lots of H+
High pH = basic = lots of OH-
For aqueous solutions, the product of the hydrogen-ion
concentration and the hydroxide-ion concentration equals 1.0 ×
10−14
[H+] + [OH-] = 1.0 × 10-14 =
0.00000000000001
Concentration of H+
Concentration of OH-
ion-product constant for water (Kw) - The product of the
concentrations of the hydrogen ions and the hydroxide ions.
+
−
Kw = [H ] × [OH ] = 1.0 × 10
Using this we know that
● Neutral solutions are half [H+] and half [OH-]
○ [H+] = (1.0 x 10-7) = 0.0000001
○ [OH-] = (1.0 x 10-7)
● Acidic solutions have more [H+].
● Basic solutions have more [OH-].
−14
If the [H+] in a solution is 1.0 × 10−5M, is the solution acidic,
basic, or neutral? What is the [OH−] of this solution?
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[H+] = 1.0 × 10−5M
Is this less than or greater than 1.0 × 10−7M?
1.0 × 10−5M = 0.00005 M
1.0 × 10−7M = 0.0000001 M
Greater! Thus the solution is acidic.
If the [H+] in a solution is 1.0 × 10−5M, is the solution acidic,
basic, or neutral? What is the [OH−] of this solution?
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Rearrange - Kw = [H+] × [OH−]
[OH-] = Kw / [H+]
Kw = 1.0 x 10-14 M
[H+]=1.0 x 10-5 M
[OH-] = (1.0 x 10-14 M) / (1.0 x 10-5 M) = 1.0 x 10-9 M
How is pH defined - the negative logarithm of the hydrogenion concentration.
+
pH = −log[H ]
In pure water OR a neutral solution, the [H+] = 1 × 10−7M,
and the pH is 7.
This means....
pH = −log(1 × 10−7) = −(log 1 + log 10−7)
pH = −(0.0 + (−7.0)) = 7.0
If the [H+] of a solution is greater than 1× 10−7M, the pH
is less than 7.0.
● If the [H+] is less than 1× 10−7M, the pH is greater than
7.0.
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If a solutions [H+] = 0.0010M or 1.0 × 10−3M, what is the
pH?
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pH = -log[H+]
pH = −log(1 × 10−3)
pH = -(log 1 + log 10-3)
pH = -(0 + (-3))
pH = 3
The pH of an unknown solution has a pH of 6.35. What is
the hydrogen-ion concentration of the solution?
● pH = -log[H+]
● 6.35 = -log[H+]
● [H+] = antilog( -6.35)
● [H+] = 10(-6.35) = 0.000000447 = 4.47 x 10-7 M
● [H+] = 0.000000447 = 4.47 x 10-7 M
What is the pH of a solution if [OH−] = 4.0 × 10−11M?
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First use Kw = [H+] × [OH−] to find the [H+].
Then use pH = -log[H+] to find the pH.
Rearrange - Kw = [H+] × [OH−]
[H+] = Kw / [OH-]
Kw = 1.0 x 10-14 M (always!)
[H+]= 1.0 x 10-14 M / 4.0 x 10-11 M
[H+] = 0.25 x 10-3 M = 2.5 x 10-4 M
What is the pH of a solution if [OH−] = 4.0 × 10−11M?
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[H+] = 0.25 x 10-3 M = 2.5 x 10-4 M
Now plug into pH = -log[H+]
pH = -log(2.5 x 10-4 M)
pH = 3.6
What is the pH of a solution if [OH−] = 4.0 × 10−11M?
● How can we check this???
● pH = 3.6
● A solution in which [OH−] is less than 1 × 10−7M is
acidic because [H+] is greater than 1 × 10−7M.