Acids And Bases
description
Transcript of Acids And Bases
ACIDS AND BASESChemistry Ms. Piela
Key Characteristics of Acids & BasesAcids
Taste sour
Reacts with alkali metals
(forms H2 gas)
Forms electrolyte solutions (conducts electricity)
pH paper color: Red
Neutralizes Bases
BasesTastes bitter
Slippery feel
Forms electrolyte solutions (conducts electricity)
pH paper color: Blue
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:
HCl H+ + Cl- Properties of bases are due to the
presence of OH- ions Example:
NaOH 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 H3O+ ion
Oftentimes H+ and H3O+ are used interchangeably
HCl H+ + Cl-HCl(g) + H2O(l) H3O+
(aq) + Cl-(aq)
Problems with the Arrhenius theory
Only deals with aqueous solutions (solutions in water)
Not all acids and bases contain H+ and OH- ionsExample: NH3 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+
ionsAcids are proton (H+) donors
Bases are substances that accept H+ ionsBases are proton (H+) acceptors
Example: HBr + H2O H3O+ + Br-
A B
Brønsted-Lowry Theory The behavior of NH3 can be
understood now:NH3 (aq) + H2O (l) ↔ NH4
+ (aq) + OH- (aq)
NH3 becomes NH4+, so NH3 is a
proton acceptor (or a Brønsted-Lowry base)
H2O becomes OH-, so H2O is a proton donor (or a Brønsted-Lowry acid)
Brønsted-Lowry Theory
Brønsted-Lowry TheoryConjugate Acid-Base Pairs
Definition: An acid and a base that differ only in the presence or absence of H+
Every acid has a conjugate base. Every base has a conjugate acid. These pairs only ever differ by
exactly one hydrogen ion
Brønsted-Lowry Theory Example Problems
Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base
NH3 + H2O NH4+ + OH-
AB CA CB
Brønsted-Lowry Theory Example
HCl (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
Amphoteric – substances that can act as both an acid and a base Examples: H2O, HCO3
-
Summary Of Theories
• Acids release H+
• Bases release OH-Arrheniu
s• Acids – proton donor• Bases – proton acceptor
Brønsted-Lowry
• Acids – electron acceptor• Bases – electron donorLewis
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 H3O+)
The pH scale
pH < 7Acidic
pH = 7Neutral
pH > 7Basic
pH scaleLow 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 (H3O+ and OH-) are equal.
[H3O+]=[OH-]= 1x10-7 This is because water will spontaneously dissociate
naturally:H2O (l) ↔ H3O+
(aq) + OH- (aq)
Writing the equilibrium expression for the self-ionization of water gives:
]][[ 3 OHOHKeq
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: Kw Unlike other equilibrium constants, the Kw
will always be the same value
Calculations of H3O+/OH-
Example #1 What is the H3O+ concentration in a solution
with [OH-] = 3.0 x 10-4 M? Kw = [H3O+][OH-]1 x 10-14 = [H3O+][3.0 x 10-4]
Mx 114-
14
103.310 x 3.010 x 1.0
_________________________3.0 x 10-4 3.0 x 10-4
Calculations of H3O+/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? Kw = [H3O+][OH-]
1 x 10-14 = [1 x 10-3][OH-]
MxxxOH 11
3
14
101100.1
101][
_________________________1.0 x 10-3 1.0 x 10-3
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?)00010.0log(pH4pH 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?
Kw = [H3O+][OH-] pH = -log [H3O+] ]0136.0][[101 3
14 MOHxKw
MxxOH 1314
3 10353.70136.0101][
1.12)10353.7log( 13 xpH
Base!
Calculating pH of a solution Practice #1
Practice #2
Calculating H3O+/OH- from pH Example #1
What is the hydronium ion concentration in fruit juice that has a pH of 3.3?
[H3O+] = 10-pH
MxOH 43.3 100.510][ 3
Calculating H3O+/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?
[H3O+] = 10-pH Kw = [H3O+][OH-]
MxOH . 60553 1091.810][
]][1091.8[101 614 OHxxKwMx
xxOH 9
6
14
1012.11091.8
101][
Calculating H3O+/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:
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%
ExampleHF H+ + F-
? M? M1 M
H+
F-
F-
F-H+
H+ HFHF
Naming Bases Bases are soluble metal hydroxides
Follow identical naming rules for ionic compounds Examples
NaOH
Ba(OH)2
NH3
NH4+
Sodium hydroxide
Barium hydroxide
Ammonia
Ammonium
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”
ExampleHCl Hydrochloric acid
Naming Acids Example
HBr
Practice HI
H2S
Hydrobromic acid
Hydroiodic acid
Hydrosulfuric acid
Naming Acids Polyatomic acids (aka oxoacids,
HxAyOz) 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 AcidsExamples
HClO4
HClO2
Sulfuric acid
Perchloric acid
Chlorous acid
H2SO4