Acids and Bases
Chapter 15
1
Properties of Acids
• Sour taste
• Change color of vegetable dyes
• React with “active” metals
– Like Al, Zn, Fe, but not Cu, Ag or Au
Zn + 2 HCl ZnCl2 + H2
– Corrosive
• React with carbonates, producing CO2
– Marble, baking soda, chalk
CaCO3 + 2 HCl CaCl2 + CO2 + H2O
• React with bases to form ionic salts
– And often water
2
Properties of Bases
•
•
•
•
Also Known As Alkalis
Taste bitter
Feel slippery
Change color of vegetable dyes
– Different color than acid
– Litmus = blue
• React with acids to form ionic salts
– And often water
– Neutralization
3
Arrhenius Theory
• Acids ionize in water to H+1 ions and anions
• Bases ionize in water to OH-1 ions and cations
• Neutralization reaction involves H+1 combining
with OH-1 to make water
• H+ ions are protons
• Definition only good in water solution
• Definition does not explain why ammonia
solutions turn litmus blue
– Basic without OH- ions
4
Brønsted-Lowery Theory
• H+1 transfer reaction
– Since H+1 is a proton, also known as proton transfer reactions
• Acid is H+ donor; Base is H+ acceptor
– Base must contain an unshared pair of electrons
• In the reaction, a proton from the acid molecule is
transferred to the base molecule
– H forms a bond to lone pair electrons on the base molecule
– We consider only 1 H transferred in each reaction
• Products are called the Conjugate Acid and Conjugate
Base
– After reaction, the original acid is the conjugate base and the
original base is changed to what is now called the conjugate
acid
5
Brønsted-Lowery Theory
H-A + :B  A-1 + H-B+1
A-1 is the conjugate base, H-B+1 is the conjugate acid
• Conjugate Acid-Base Pair is either the original
acid and its conjugate base or the original base and
its conjugate acid
– H-A and A-1 are a conjugate acid-base pair
– :B and H-B+1 are a conjugate acid-base pair
• The conjugate base is always more negative than
the original acid; and the conjugate acid is always
more positive than the original base
6
Example #1
Write the conjugate base for the acid H3PO4
• Determine what species you will get if you
remove 1 H+1 from the acid
– The Conjugate Base will have one more
negative charge than the original acid
H3PO4  H+1 + H2PO4-1
7
Brønsted-Lowery Theory
• In this theory, instead of the acid, HA, dissociating into
H+1(aq) and A-1(aq); The acid donates its H to a water
molecule
HA + H2O  A-1 + H3O+1
A-1 is the conjugate base, H3O+1 is the conjugate acid
• H3O+1 is called hydronium ion
• In this theory, substances that do not have OH-1 ions
can act as a base if they can accept a H+1 from water
H2O + :B  OH-1 + H-B+1
8
Strength of Acids & Bases
• The stronger the acid, the more willing it is to donate H
• Strong acids donate practically all their H’s
HCl + H2O  H3O+1 + Cl-1
• Strong bases will react completely with water to form
hydroxides
CO3-2 + H2O HCO3-1 + OH-1
• Weak acids donate a small fraction of their H’s
– The process is reversible, the conjugate acid and conjugate
base can react to form the original acid and base
HC2H3O2 + H2O  H3O+1 + C2H3O2-1
• Only small fraction of weak base molecules pull H off
water
HCO3-1 + H2O H2CO3 + OH-1
9
Figure 15.1: Graphical
representation of the
behavior of acids in
aqueous solution
10
Figure 15.2: The
relationship of acid
strength and
conjugate base
strength
11
Multiprotic Acids
• Monoprotic acids have 1 acid H,
diprotic 2, etc.
– In oxyacids only the H on the O is
acidic
• In strong multiprotic acids, like
H2SO4, only the first H is strong;
transferring the second H is usually
weak
H2SO4 + H2O  H3O+1 + HSO4-1
HSO4-1 + H2O  H3O+1 + SO4-2
12
Water as an Acid and a Base
• Amphoteric substances can act as either an
acid or a base
– Water as an acid, NH3 + H2O  NH4+1 + OH-1
– Water as a base, HCl + H2O  H3O+1 + Cl-1
• Water can even react with itself
H2O + H2O  H3O +1 + OH-1
13
Autoionization of Water
• Water is an extremely weak electrolyte
– therefore there must be a few ions present
H2O + H2O  H3O+1 + OH-1
• all water solutions contain both H3O+1 and OH-1
– the concentration of H3O+1 and OH-1 are equal
– [H3O+1] = [OH-1] = 10-7M @ 25°C
• Kw = [H3O+1] x [OH-1] = 1 x 10-14 @ 25°C
– Kw is called the ion product constant for water
– as [H3O+1] increases, [OH-] decreases
14
Acidic and Basic Solutions
• acidic solutions have a larger [H+1] than [OH-1]
• basic solutions have a larger [OH-1] than [H+1]
• neutral solutions have [H+1]=[OH-1]= 1 x 10-7 M
[H+1]
-14
1
x
10
=
[OH-1]
[OH-1]
-14
1
x
10
=
[H+1]
15
Example #2
Determine the [H+1] and [OH-1] in a
10.0 M H+1 solution
 Determine the given information and the
information you need to find
Given [H+1] = 10.0 M
Find [OH-1]
 Solve the Equation for the Unknown
Amount
1
Kw  [H ] x [OH ]
Kw
-1
[OH ]  1
[H ]
-1
16
Example #2
Determine the [H+1] and [OH-1] in a
10.0 M H+1 solution
 Convert all the information to Scientific
Notation and Plug the given information
into the equation.
Given [H+1] = 10.0 M
= 1.00 x 101 M
Kw = 1.0 x 10-14
Kw
[OH ]  1
[H ]
-1
-14
1.0
x
10
-15
[OH -1 ] 

1.0
x
10
M
1
1.00 x 10
17
pH & pOH
• The acidity/basicity of a solution is often expressed as
pH or pOH
• pH = -log[H3O+1]
pOH = -log[OH-1]
– pHwater = -log[10-7] = 7 = pOHwater
• [H+1] = 10-pH
[OH-1] = 10-pOH
• pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral
• The lower the pH, the more acidic the solution; The higher
the pH, the more basic the solution
• 1 pH unit corresponds to a factor of 10 difference in acidity
• pOH = 14 - pH
18
Figure 15.3: The pH
scale and pH values
of some common
substances
19
Figure 15.4: A pH meter
20
Figure 15.5: Indicator
paper being used to
measure the pH of a
solution
21
Example #3
Calculate the pH of a solution with a
[OH-1] = 1.0 x 10-6 M
Find the concentration of [H+1]
Kw
[H ] 
1
[OH ]
1
-14
1.0 x 10
-8
[H ] 

1.0
x
10
M
-6
1.0 x 10
1
22
Example #3
Calculate the pH of a solution with a
[OH-1] = 1.0 x 10-6 M
Enter the [H+1] concentration into your
calculator and press the log key
log(1.0 x 10-8) = -8.0
Change the sign to get the pH
pH = -(-8.0) = 8.0
23
Example #4
Calculate the pH and pOH of a
solution with a [OH-1] = 1.0 x 10-3 M
Enter the [H+1] or [OH-1]concentration into
your calculator and press the log key
log(1.0 x 10-3) = -3.0
Change the sign to get the pH or pOH
pOH = -(-3) = 3.0
Subtract the calculated pH or pOH from
14.00 to get the other value
pH = 14.00 – 3.0 = 11.0
24
Example #5
Calculate the [OH-1] of a solution with a pH of 7.41
If you want to calculate [OH-1] use pOH, if you
want [H+1] use pH. It may be necessary to convert
one to the other using 14 = pH + pOH
pOH = 14.00 – 7.41 = 6.59
Enter the pH or pOH concentration into your
calculator
Change the sign of the pH or pOH
-pOH = -(6.59)
Press the button(s) on you calculator to take the
inverse log or 10x
[OH-1] = 10-6.59 = 2.6 x 10-7
25
Calculating the pH of a Strong,
Monoprotic Acid
• A strong acid will dissociate 100%
HA  H+1 + A-1
• Therefore the molarity of H+1 ions will be
the same as the molarity of the acid
• Once the H+1 molarity is determined, the pH
can be determined
pH = -log[H+1]
26
Example #6
Calculate the pH of a 0.10 M HNO3 solution
Determine the [H+1] from the acid concentration
HNO3  H+1 + NO3-1
0.10 M HNO3 = 0.10 M H+1
Enter the [H+1] concentration into your calculator
and press the log key
log(0.10) = -1.00
Change the sign to get the pH
pH = -(-1.00) = 1.00
27
Buffered Solutions
• Buffered Solutions resist change in pH when an acid
or base is added to it.
• Used when need to maintain a certain pH in the
system
– Blood
• A buffer solution contains a weak acid and its
conjugate base
• Buffers work by reacting with added H+1 or OH-1 ions
so they do not accumulate and change the pH
• Buffers will only work as long as there is sufficient
weak acid and conjugate base molecules present
28