acids & bases. key characteristics of acids & bases acids taste sour reacts with alkali...
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Acids & Bases
Key Characteristics of Acids & Bases
AcidsTaste sour
Reacts with alkali metals (forms H2 gas)
Litmus paper: Red
Neutralizes Bases
BasesTastes bitter
Slippery feel
Litmus paper: Blue
Neutralizes Acids
Theories of Acids & Bases
Arrhenius Theory of Acids & Bases Properties of acids are due to the
presence of H+ ions Example:
HCl H+ + Cl- Properties of bases are due to the
presence of OH- ions Example:
NaOH Na+ + OH-
H+ ions in water
H+ ions are bare protons These H+ ions react strongly with the
nonbonding pair of electrons in a water molecule This forms the hydronium ion, H3O+
Oftentimes H+ and H3O+ are used interchangeably
HCl H+ + Cl-
HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)
Problems with Arrhenius
Arrhenius theory has limitations: Only deals with aqueous solutions
(solutions in water) Not all acids and bases produce H+ and
OH- ionsNH3 for example is a base
Brønsted and Lowry proposed a definition based on acid base reactions transferring H+ ion from one substance to another
Brønsted-Lowry Theory
Theories of Acids & Bases
Brønsted-Lowry Theory Acids are substances that donate H+ ions
Acids are proton donors Bases are substances that accept H+
ions Bases are proton acceptors
Example:
HBr + H2O H3O+ + Br-
A B
Brønsted-Lowry Theory
The behavior of NH3 can now be understood:
NH3 (aq) + H2O (l) ↔ NH4+
(aq) + OH-
(aq)
Since NH3 becomes NH4+, it is a
proton acceptor (or a Brønsted-Lowry base)
H2O becomes OH-, which means it is a proton donor (or a Brønsted-Lowry acid)
Brønsted-Lowry Theory
Conjugate Acid-Base Pairs An acid and a base that differ only in the
presence or absence of H+ are called a conjugate acid-base pair.
Every acid has a conjugate base.Every base has a conjugate acid.
HX is the conjugate acid of X-
H2O is the conjugate base of H3O+
Brønsted-Lowry Theory
These pairs differ by only one hydrogen ion Example
Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base
NH3 + H2O NH4+ + OH-
B A CA CB NH3 acts as a Brønsted base by accepting
a proton. Water acts as a Brønsted acid by donating
a proton.
Brønsted-Lowry Theory
ExampleHCl (g) + H2O (l) ↔ H3O+
(aq) + Cl- (aq)
HSO4- + HCO3
- ↔ SO4-2 + H2CO3
BA CA CB
A B CACB
Theories of Acids & Bases
Lewis Acids & Bases Acids are electron acceptors Bases are electron donors
Example: H2O + NH3 OH- + NH4
+
Is really: H2O + :NH3 OH- + H:NH3
+
Electron pair
donor(NH3)
Electron pair acceptor(H+)
Summary Of Theories
•Acids release H+
•Bases release OH-
•Defines acids & bases in H2
O
Arrhenius
•Acids – proton donor
•Bases – proton acceptor
•Can define acids & bases in solvents other than H2
O
Brønsted-Lowry
•Acids – electron acceptor
•Bases – electron donor
•Defines acids & bases without a solvent
Lewis
The Self-Ionization of Water
Even pure water contains a small number of ions:
H2O (l) ↔ H3O+ (aq) + OH- (aq)
In pure water, the concentrations of the ions (H3O+ and OH-) are equal.
[H3O+]=[OH-]= 1x10-7 M
The Self-ionization of Water
Writing the equilibrium expression for the self-ionization of water gives:
Plugging in the concentrations in pure water, this gives an equilibrium constant of 1x10-14 this is referred to as the ion product constant
of water This ion product constant of water is given
the symbol Kw
]][[ 3 OHOHKeq
The Self-ionization of water
Example #1 What is the H3O+ concentration in a solution
with [OH-] = 3.0 x 10-4 M?Kw = [H3O+][OH-]
1x10-14 = [H3O+][3.0x10-4]
114-
14
103.310 x 3.0
10 x 1.0
x
Example #2
If the hydroxide-ion concentration of an aqueous solution is 1.0 x 10-3 M, what is the [H3O+] in the solution? Kw = [H3O+][OH-]
1x10-14 = [H3O+][1.0x10-3]
113
14
3 101100.1
101][
x
x
xOH
The pH scale
Developed by Søren Sørensen in order to determine the acidity of ales
Used in order to simplify the concept of acids and bases
The pH scale goes from 1 to 14 A change in one pH unit corresponds to a
power of ten change in the concentration of hydronium (H3O+) ions A pH = 2.0 has 10 times the concentration of
H3O+ than a pH = 3.0, and 100 times greater than pH = 4
The pH scale
pH < 7• Acid
pH = 7• Neutral
pH > 7• Base
Calculations of pH
pH can be expressed using the following equation:
pH = -log [H3O+] or [H3O+] = 10-pH
Example #1 What is the pH of a solution with 0.00010 M
H3O+? Is this solution an acid or a base?
Acid
)00010.0log(pH
4
Calculations of pH
Example #2 What is the pH of a solution with the
concentration of hydroxide ions 0.0136 M? Is this an acid or a base?pH = -log [H3O+] Kw = [H3O+]
[OH-]
Base
]0136.0][[101 314 MOHxKw
13
14
3 10353.70136.0
101][
x
xOH
1.12)10353.7log( 13 xpH
Calculations of pH
Practice #1
Practice #2
Calculations of pH
Example #1 What is the hydronium ion concentration in
fruit juice that has a pH of 3.3?[H3O+] = 10-pH
43.3 100.510][ 3 xOH
Calculations of pH
What are the concentrations of the hydronium and hydroxide ions in a sample of rain that has a pH of 5.05?
[H3O+] = 10-pH Kw = [H3O+][OH-]
60553 1091.810][ xOH .
]][1091.8[101 614 OHxxKw9
6
14
1012.11091.8
101][
x
x
xOH
Calculation of pH
Practice #1
Practice #2
Strength of Acids & Bases
When a solution is considered strong, it will completely ionize in a solution Nitric acid is an example of strong acid:
HNO3 (l) + H2O (l) NO3- (aq) + H3O+ (aq)
In a solution of nitric acid, no HNO3 molecules are present
Strength is NOT equivalent to concentration!
Strength of Acids & Bases
Knowing the strength of an acid is important for calculating pH If given concentration of strong acid (such
as HNO3) assume it is the same as the concentration of hydronium, H3O+, ions
Given concentration of a strong base, assume it has the same concentration as the hydroxide, OH-, ions
Strong Acids & Bases Ionize 100%
ExampleNaOH Na+ + OH-
1 M1 M1 M
Na+
Na+
Na+
OH-OH-
OH-
Weak Acids & Bases Ionize X%
Example
HF H+ + F-
? M? M1 M
H+
F-
F-
F-
H+
H+ HF
HF
Strength of Acids & Bases
Stronger the acid
Weaker the conj. base
Stronger the base
Weaker the conj. acid
Strength of Acids & Bases
Strong Acids
•Perchloric acid, HClO4
•Chloric acid, HClO3
•Hydrochloric acid, HCl•Hydrobromic acid, HBr•Hydroiodic acid, HI•Nitric acid, HNO3
•Sulfuric acid, H2SO4
Strong
Acids
Must be memorized!
Strong Acids
6 of 7 strong acids are monoprotic (HX) Exists only as H ions and X ions
HI(aq) H+(aq) + I-(aq)
2M HI = [H+]= [I-] = 2M Determining pH of Strong Acids
For Strong Acids: pH = -log [H+] For monoprotic strong acids: [H+] = [X]