Transcript Chapter
ACIDS AND BASES
Chemistry Ms. Piela
Key Characteristics of Acids & Bases
Acids Bases
Taste sour Tastes bitter Reacts with alkali metals (forms H 2 gas) Forms electrolyte solutions (conducts electricity) pH paper color: Red Slippery feel Forms electrolyte solutions (conducts electricity) pH paper color: Blue Neutralizes Bases Neutralizes Acids
The 3 Main Theories of Acids/Bases This course will mainly deal with BL theory
Theories of Acids & Bases Arrhenius Theory of Acids & Bases: Properties of acids are due to the presence of H + ions Example: H Cl H + + Cl Properties of bases are due to the presence of OH ions Example: Na OH Na + + OH -
What is an H + ?
H + ions are bare protons These are so reactive that they do not exist naturally, but will bond with water to form a hydronium ion, or H 3 O + ion Oftentimes H + and H 3 O + HCl are used interchangeably H + + Cl HCl (g) + H 2 O (l) H 3 O + (aq) + Cl (aq)
Problems with the Arrhenius theory Only deals with aqueous (solutions in water) solutions Not all acids and bases contain H+ and OH- ions Example: NH 3 is a base Considered the most incomplete theory of acids and bases
Theories of Acids & Bases Brønsted-Lowry Theory of Acids & Bases Acids are substances that donate H + ions Acids are proton (H + ) donors Bases are substances that accept H + ions Bases are proton (H + ) acceptors Example: HBr + H 2 O A B H 3 O + + Br -
Brønsted-Lowry Theory The behavior of NH 3 NH 3 (aq) + H 2 O (l) can be understood now: ↔ NH 4 +
(aq)
+ OH -
(aq)
NH 3 becomes NH 4 + , so NH 3 acceptor is a proton (or a Brønsted-Lowry base) H 2 O becomes OH , so H 2 O is a proton donor (or a Brønsted-Lowry acid)
Brønsted-Lowry Theory
Brønsted-Lowry Theory
Brønsted-Lowry Theory Example Problems Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base NH 3 + H 2 O NH 4 + + OH -
Brønsted-Lowry Theory Example HCl
(g)
+ H 2 O
(l)
↔ H 3 O +
(aq)
+ Cl -
(aq)
HSO 4 + HCO 3 ↔ SO 4 -2 + H 2 CO 3
Theories of Acids & Bases Lewis Acids & Bases Acids are electron acceptors Bases are electron donors Amphoteric – substances that can act as both an acid and a base Examples: H 2 O, HCO 3 -
Summary Of Theories
Arrhenius
• Acids release H + • Bases release OH -
Brønsted Lowry
• Acids – proton donor • Bases – proton acceptor
Lewis
• Acids – electron acceptor • Bases – electron donor
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 for his workers The pH scale goes from 0 to 14 The acidity/basicity of the solutions depends on the concentration of H + (or H 3 O + )
The pH scale pH < 7 pH = 7 pH > 7
pH scale Low pH values means a high concentration of H + (acidic) High pH values means a low concentration of H + (basic)
Calculations of pH The Self Ionization of Water In pure water (pH = 7 ), the concentrations of the ions (H 3 O + and OH ) are equal .
[H 3 O + ]=[OH ]= 1x10 -7 This is because water will spontaneously dissociate naturally: H 2 O (l) ↔ H 3 O +
(aq)
+ OH -
(aq)
Writing the equilibrium expression for the self-ionization of water gives:
K eq
[
H
3
O
][
OH
]
The Self-ionization of Water 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 The ion product constant of water has its own symbol: K w Unlike other equilibrium constants, the K w the same value will always be
Calculations of H 3 O + /OH Example #1 What is the H 3 O + = 3.0 x 10 -4 M ?
concentration in a solution with [OH ] K w = [H 3 O + ][OH ] 1 x 10 -14 = [H 3 O + ][3.0 x 10 -4 ] 3.0 x 10 -4 3.0 x 10 -4 1.0
x 10 14 3.0
x 10 4 3 .
3
x
10 11
M
Calculations of H 3 O + /OH If the hydronium-ion concentration of an aqueous solution is 1.0 x 10 -3 M, what is the hydroxide ion concentration in the solution? K w = [H 3 O + ][OH ] 1 x 10 -14 = [1 x 10 -3 ][OH ] 1.0 x 10 -3 1.0 x 10 -3 [
OH
] 1
x
10 14 1 .
0
x
10 3 1
x
10 11
M
Calculations of pH pH can be expressed using the following equation: pH = -log [H 3 O + ] Example #1 or [H 3 O + ] = 10 -pH What is the pH of a solution with 0.00010 M H 3 O + ? Is this solution an acid or a base?
pH pH
4 log( 0 .
00010 ) Acid!
Calculating pH of a solution Example #2 What is the pH of a solution where the concentration of hydroxide ions is 0.0136 M? Is this an acid or a base?
K w = [H 3 O + ][OH ] pH = -log [H 3 O + ] [
H
3
O
]
pH
K w
1
x
10 14 1
x
10 0 .
0136 14 [
H
7 .
353 3
O x
10 ][ 0 .
0136
M
13
M
] Base!
log( 7 .
353
x
10 13 ) 12 .
1
Calculating pH of a solution Practice #1 Practice #2
Calculating H 3 O + /OH from pH Example #1 What is the hydronium ion concentration in fruit juice that has a pH of 3.3?
[H 3 O + ] = 10 -pH [
H
3
O
] 10 3 .
3 5 .
0
x
10 4
M
Calculating H 3 O + /OH from pH What are the concentrations of the hydronium and hydroxide ions in a sample of rain that has a pH of 5.05?
[H 3 O + ] = 10 -pH K w = [H 3 O + ][OH ] [
H
3
O
] 10 5
.
05
K w
8 .
91
x
10 6
M
1
x
10 14 [ 8 .
91
x
10 6 ][
OH
] [
OH
] 1
x
10 14 8 .
91
x
10 6 1 .
12
x
10 9
M
Calculating H 3 O + /OH from 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: HNO 3 (l) + H 2 O (l) ⇋ NO 3 -
(aq)
+ H 3 O +
(aq)
In a solution of nitric acid, no HNO 3 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 HNO 3 ) assume it is the same as the concentration of hydronium , H 3 O + , ions Given concentration of a strong base, assume it has the same concentration as the hydroxide , OH , ions
Strong Acids & Bases Ionize 100% Example 1 M NaOH 1 M Na + 1 M + OH Na + OH Na + OH Na OH +
Weak Acids & Bases Ionize X% Example 1 M HF ?
M H + ?
M + F H + F HF F H + HF F H +
Naming Bases Bases are soluble metal hydroxides Follow identical naming rules for ionic compounds Examples NaOH Ba(OH) 2 NH 3 NH 4 +
Naming Acids Binary Acids (HX) If the acid has an anion that ends in “-ide ” use the following basic format to name the acid: “ Hydro – root – ic acid ” Example HCl
Naming Acids Example HBr Practice HI H 2 S
Naming Acids Polyatomic acids (aka oxoacids, H x A y O z ) Name depends on the polyatomic used: If polyatomic ends in “ -ite ”, replace with “ ous acid ” If polyatomic ends in “ -ate ”, replace with “ ic acid ” Trick: “I ate something icky ”
Naming Acids Examples HClO 4 HClO 2 Sulfuric acid