CHEMISTRY
MRS. SINGH
CHAPTER 19 NOTES
SXN. 19.1
Acids
H+ ions
Taste sour
Sharp taste(lemons)
Bases
OH- ions
Taste bitter
Feel slippery (soap)
Definition of acids & bases:
1. Arrhenius
Acid = substance that contains hydrogen and ionizes in aqueous solution to
produce hydrogen ions Ex. HBr(g) ↔ H+(aq) + Br-(aq)
Base = substance that contains a hydroxide group and dissociates in aqueous
solution to produce hydroxide ions NaOH(s) ↔ Na+ (aq) + OH-(aq)
2. Bronsted-Lowry
Acid = hydrogen-ion donor; what remains is the conjugate base
Base = hydrogen-ion acceptor; what forms is the conjugate acid
Ex. NH3(aq) + H2O(l) ↔ NH4 (aq)+ +
OH(aq) base
acid
conjugate
conjugate
acid
base
Ex. Do # 2 on pg. 599
Monoprotic acid = an acid that can donate only one hydrogen ion
Ex. HCl(aq) + H2O(l) ↔ H3O+(aq) + Cl- (aq)
Polyprotic acid = acids that can donate more than one hydrogen ion
Ex. Write the steps in the complete ionization of boric acid H3BO3
1
Amphoteric = substances which can act as both acids and bases Ex. H2O
Anhydrides = oxides which become acids or bases by adding water
Ex. CO2(g) + H2O(l) → H2CO3 (aq)
Nonmetal oxide + water = acid
Ex. CaO(s) + H2O(l) → Ca+2(aq) + 2OH-(l)
Metal oxide + water = base
SXN. 19.2
Strong acid = an acid that ionizes completely in dilute aqueous solution
All strong acids have weak conjugate bases HX(aq)+ H2O(l) →H3O+(aq) + X-(aq)
Copy strong acids & ionization equations from table 19-1 on pg. 603
Weak acid = an acid that ionizes only partially in dilute aqueous solution
Any acid that is not a strong acid is a weak acid!
Ka = acid ionization constant
Ex. HCN(aq) + H2O (l) ↔ H3O + (aq) + CN- (aq)
Keq =
Do: #10 on pg. 605
2
SXN. 19.3
The Dissociation of Water
Since water is an amphoteric substance, it can act as an acid or a base
depending on what is dissolved in it.
CO32–(aq) + H2O → HCO3–(aq) + OH–(aq)
Water is acting as an acid since it is donating a proton to the
carbonate ion.
HCl(aq) + H2O → H3O+(aq) + Cl–(aq)
Water is acting as a base since it is accepting a proton from hydrochloric
acid.
H+ (aq) + OH– (aq)
H2O (l)
H2O (l)
+
Acid – Conjugate Base Pair
H3O+ (aq) + OH–(aq)
H2O (l)
Base – Conjugate Acid Pair
When water itself dissociates, it is acting as an acid and base at the same
time.
The K in the dissociation of water shown above is very small. It is also
always the same number if the temperature remains constant, say 25°C.
It is called the ION – PRODUCT CONSTANT.
Kw = [H3O+] [OH–] = [H+] [OH–] = 1.0 x 10–14
A neutral solution is where [H+] = [OH–]
An acidic solution is where [H+] > [OH–]
A basic solution is where [H+] < [OH–]
The pH(power of Hydrogen) scale measures the degree of acidity of a
solution and is based on common logarithms. The pH scale is a measurement
of the concentration of hydrogen ions [H+]. The numbers on the pH scale
are the negative logarithm of the [H+].
Since Kw = [H+] [OH–] = 1.0 x 10–14 (at 25°C), the pH scale measures
from 0 to 14.
[H+]
7
0
1
2
3
4
5
6
8
increasing acidity
NEUTRAL
9
10
11
12
increasing alkaline (base)
13
3
14
 Since the exponents are negative numbers, the greater the
concentration of H+, the smaller the pH number.
 The lower the pH number, the more acidic the solution.
 The higher the pH number, the more alkaline.
Because pH numbers are powers of ten, an acid of pH = 4 is ten times more
acidic than a solution with a pH = 5, and one-hundred times more acidic
than a solution with a ph = 6.
In a neutral solution, [H+] = [OH–] = 1.0 x 10–7
pH = – log [H+] = – log [10–7] = 7. So a pH of 7 is neutral.
pOH = -log [OH-]
pH + pOH = 14.00
Ex. The concentration of either the H+ ion or the OH- ion is given for three
aqueous solutions at 298 K. For each solution, calculate [H+] or [OH-].
State whether the solution is acidic, basic, or neutral.
a. [H+] = 1.0 x 10-4 M
b. [OH-] = 1.3 x 10-2 M
c. [H+] = 5.8 x 10-11 M
The rule for pH is: The number of significant figures the concentration has,
is the number of decimal places the pH should have.
Ex. Calculate the pH at 298 K of solutions having the following ion
concentrations.
a. [H+] = 1.0 x 10-6 M
b. [H+] = 3.8 x 10-11 M
Ex. Calculate the pOH and pH at 298 K of solutions having the following
ion concentrations.
a. [OH-] = 1.0 x 10-12 M
b. [OH-] = 1.3 x 10-2 M
4
Calculating ion concentrations from pH
a. What are [H+] and [OH-] in a healthy person’s blood that has a pH of
7.50? Assume temp is 298 K.
b. Find the [H+] and [OH-] if pOH = 2.95
Calculating the pH of solutions of strong acids and strong bases
For all strong monoprotic acids, the concentration of the acid is the
concentration of H+ ion.
Ex. HCl(aq) → H+(aq) + Cl-(aq) If you have a 0.1M HCl solution, it
contains 0.1 mole of H+ ions per liter and 0.1 mole of Cl- ions per liter
Most strong bases behave the same way: Ex. KOH(aq) → K+(aq) + OH-(aq)
A 0.5 M NaOH produces 0.5 mol of K+ ions per liter and 0.5 mol of OHions per liter.
Ex. What is the pH of a 0.0001 M solution of HCl?
Since HCl H+ + Cl–, a 1.0 x 10–4 concentration of HCl produces a
1.0 x 10–4 concentration of H+ ions. So, pH = – log [H+] = – log
[1.0 x 10–4] = 4.00
Ex: What is the pH of a 0.00432 M solution of HCl?
Ex: What is the pH of a 0.0329 M solution of NaOH?
5
Using pH to calculate Ka
Ex. Calculate the Ka for the following acids using the given
information:
a. 0.220 M solution of H3AsO4, pH = 1.50
b. 0.0400 M solution of HClO2, pH = 1.80
Measuring pH
The approximate pH of a solution can be obtained by wetting a piece of pH
paper with the solution and comparing the color of the wet paper with a set
of standard colors.
The pH meter provides a more accurate measurement in the form of a digital
display of the pH.
SXN. 19.4 NEUTRALIZATION
A neutralization rxn is a rxn in which an acid and a base react in aqueous
solution to produce a salt and water.
A salt is an ionic compound made up of a cation from a base and an anion
from an acid.
Neutralization is a double-replacement rxn
Ex. Mg(OH)2(aq) + 2 HCl(aq) → MgCl2(aq) + 2H2O(l)
Note that the cation from the base (Mg
the acid (Cl-) in the salt MgCl2
+2
) is combined with the anion from
Do #29 on pg. 617
6