Chapter 10 Acids and Bases
Acids produce H+ ions in water
H2O
HCl(g)
H+(aq) + Cl(aq)
they are electrolytes
have a sour taste
turn litmus red
neutralize bases
Some acids like sulfuric and phosphoric release more than 1 H+ in water; other like acetic acid
(vinegar) release far less than 1 H+ per molecule
Bases produce OH− ions in water
are electrolytes
feel soapy and slippery
neutralize acids
NaOH
KOH
sodium hydroxide
potassium hydroxide
sodium and potassium hydroxide release 1 OH- /molecule
other bases such as ammonium hydroxide (NH4OH)
release far fewer OH-
fewer
vinegar C2H4O2
Cola H3PO4
Milk of Magnesia (Mg(OH)2
Tums Ca(OH)2
Strong acids completely ionize (100%) in aqueous solutions.
HCl(g) + H2O(l)
H3O+(aq) + Cl−(aq)
Amount of acid added
Weak acids dissociate only slightly in water to form a
solution of mostly molecules and a few ions.
H2CO3(aq) + H2O(l)
H3O+(aq) + HCO3−(aq)
NH3(g) + H2O(l)
H2CO3 + OHNaOH
NH4+(aq) + OH−(aq)
HCO3- + H2O
Na+ + OH-
CO3=
Windex
weak base
+ H3O+ Baking Soda weak base
Drano strong base
Water reacts with itself in the following manner:
H+ is transferred from one H2O molecule to another ;
one water molecule acts as an acid, while another acts as a base
H2O + H2O
H3O+ + OH−
..
..
..
..
H:O: + H:O:
H:O:H+ +
:O:H−
The concentration of
..
..
..
..
H3O+ = OH- = 10-7 mols/L
H
H
H
water
water
hydronium hydroxide
ion(+)
ion(-)
pH
The pH of a solution is used to indicate
the acidity of a solution;
it has values that usually range
from 0 to 14;
the solution is acidic when the
values are less than 7;
the solution is neutral with
a pH of 7;
the solution is basic when the
values are greater than 7
How is the numerical value of pH determined?
pH = - log[H3O+ concentration]; pOH = -log [OH- concentration]
when the H3O+ concentration is
expressed in mols/L
pH + pOH = 14
Reactions of acids and bases
Acid + Base = Salt + Water
Mg(OH)2 + HCl (gastric juice) = MgCl2 + H2O
Mg(OH)2 + 2HCl (gastric juice) = MgCl2 + 2 H2O
CaCO3 + HCl
=
CaCl2
+ H2CO3 = CaCl2 + H2O + CO2
CaCO3 + 2HCl = CaCl2 + 2H2CO3 = CaCl2 + 2H2O + 2CO2
+ burp
How does the pH vary if we add NaOH (0.1 mol/L) dropwise to a solution of HCl (0.1
mol/L)?
HCl + NaOH = H2O + NaCl
pH of resulting solution
7
6
5
pH
4
0.001
3
0.01
2
0.1
1
buffered in
this region
0
0
.
0
0 0
.
0
2 0
.
0
4 0
.
0
6 0
.
0
8 0
.
1
0 0
.
1
2
O
H
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)
How does the pH vary is we add NaOH (0.1 mol/L) dropwise to a solution of the weak
acid acetic acid (0.1 mol/L)?
HOAc
+ NaOH =
H2O + NaOAC
9
8
pH of resulting solution
7
pH
6
5
buffered in this
entire region
4
3
2
0
.
0
0 0
.
0
2 0
.
0
4 0
.
0
6 0
.
0
8 0
.
1
0
O
H
(
e
q
u
i
v
a
l
e
n
t
s
)
Suppose we have a liter of water and we either add a drop of water
containing 10-4 moles of HCl or 10-4 moles of NaOH;
What would be the resulting pH assuming no volume change with HCl addition?
H3O+ = 10-4 mol/L; pH= 4
What would be the resulting pH assuming no volume change with NaOH addition?
OH- = 10-4 mol/L; pOH = 4
pH = 14- pOH = 10
alternatively
[H+][OH-] = 1 *10-14;
[H+] = 10-10;
pH = 10
[H+] = 10-14/10-4;
How much Mg(OH)2 would be required to neutralize 100 mL of HCl that is 0.1 M?
Mg(OH)2 + HCl
=
Mg(OH)2 + 2HCl
=
MgCl2 + H2O
MgCl2 + 2 H2O
Balanced equation
How many moles of HCl are their in 100 mL of 0.1 M HCl ?
0.1 M HCl = 0.1 mol/L; 100 mL = 0.1 L
0.1 mol/L *0.1 L = 0.01 moles of HCl
0.5 Mg(OH)2 + HCl
= 0.5 MgCl2 + H2O
0.01 moles of HCl requires 0.005 moles of Mg(OH)2
What is the pH of a vinegar solution that is 0.1 M?
HOAc + H2O = H3O+ + OAcWhat is the equilibrium expression?
[H3O+][OAc-]/[HOAc] = K
K = 18*10-6
if we let x = [H3O+]; the X also = [OAc-]
x2/[0.1-x] = 18*10-6
lets assume that x is very small in comparison to 0.1
x2 = 1.8 * 10-6; x ≈ 1.3*10-3 pH = 2.74
•
•
H2CO3 is a very weak acid; however both hydrogens can be removed in the
presence of strong base; the pH of a solution of NaHCO3 is very close to
physiological pH; in the presence of an acid the HCO3- ion tends to pick up the
proton, thus buffering the solution and preventing the solution to become too acidic.
H+ + HCO3H2CO3
CO2 + H2O
In the presence of a base, the HCO3- ion can
lose its proton as H+ and thus neutralize the
strong base; thus the HCO3- ion can buffer the
solution in both directions
HCO3-
+ OH-
CO3-2
+ H2O
The pH in living systems is very important. For example the pH of blood is kept at 7.4
and must be maintained within ±0.5 pH units. How is this done?
At a pH of 7.4, most CO2 is in the form of HCO3HCO3- can react with either acid or base
HCO3- + H3O+
H2CO3
HCO3- + OH-
H2O
In this manner, HCO3stabilizes the pH and does not
allow it to become too acidic
or to basic; it acts as a buffer
+ CO3-2
CO2 + H2O
CO2 + H2O = H2CO3 = H+ + HCO3-