Acids and Bases
Jan 14, 2016
Acids and Bases
Properties of Acids
Sour taste React w/ metals to form H2
Most contain hydrogen Are electrolytes Change color in the presence of
indicators (turns litmus red) Has a pH lower than 7
Two Types of Acids
Strong acids– Any acid that dissociates completely in
aqueous sol’n Weak acids
– Any acid that partially dissociates in aqueous sol’n
Properties of Bases
Bitter taste Slippery feel Are electrolytes Change color in the presence of
indicators (turns litmus blue) Has a pH higher than 7
Types of Bases
Strong Base– Any base that dissociates completely in
aqueous sol’n Weak Base
– Any base that partially dissociates in aqueous sol’n
Neutralization
Neutralization rxn: a rxn of an acid and a base in aqueous sol’n to produce a salt and water
Salt: compound formed from the positive ion of a base and a negative ion of an acid
Properties of the acid and base cancel each other
Arrhenius Model of Acids and Bases
Proposed the model in 1887 Acid: any compound that produces H+ ions in
aqueous (water) sol’n Base: any compound that produces OH-
(hydroxide) ion in aqueous sol’n Offers an explanation of why acids and bases
neutralize each other (H+ + OH- = H2O)
Problems with Model
Restricts acids and bases to water sol’ns (similar reactions occur in the gas phase)
Does not include certain compounds that have characteristics of bases (e.g., ammonia)
Brønsted-Lowry Model of Acids and Bases
Brønsted acid: a hydrogen ion donor (H+, or proton)
Brønsted base: a hydrogen ion acceptor Defines acids and bases independently of
how they behave in water Amphiprotic: having the property of behaving
as an acid and a base– Also called amphoteric, e.g., water
Conjugate Acid-Base Pairs
The rxn between Brønsted-Lowry acids and bases can proceed in the reverse direction (reversible reactions)
HX (aq) + H2O (l) H3O+ (aq) + X- (aq) The water molecule becomes a hydronium ion
(H3O+), and is an acid because it has an extra H+ to donate
The acid HX, after donating the H+, becomes a base X-
Conjugate Acids and Bases
HX (aq) + H2O (l) H3O+ (aq) + X- (aq)Acid
Base Conjugate Acid
Conjugate Base
Forward reaction: Acid and base
Reverse reaction: Conjugate acid and conjugate base
Conjugate Acid: species produced when a base accepts a hydrogen ion from an acid
Conjugate Base: species produced when an acid donates a hydrogen ion to a base
Conjugate Acid-Base Pair: two substances related to each other by the donating and accepting of a single hydrogen ion
Types of Acids
Monoprotic acids: acids that contain only 1 hydrogen; e.g., HCl
Diprotic acids: acids that contain 2 hydrogens; e.g. H2CO3
Triprotic acids: acids that contain 3 hydrogens; e.g. H3PO4
More Types of Acids
Binary acids: acids that contain only 2 elements; e.g. HF
Polyatomic acids: acids that contain more than 2 elements; e.g. H2SO4
– These acids contain polyatomic ions– Also called ternary or oxy- acids
Naming Binary Acids
Start with the prefix hydro- Put it in front of the root word of the anion
(- charged ion) Add –ic to the end Examples
– Hydrobromic (HBr)– Hydrofluoric (HF)– Hydroiodic (HI)– Hydrochloric (HCl)
Naming Polyatomic Acids
Start with the root word of the name of the polyatomic ion
Add –ous if name ends in –ite Add -ic if name ends in –ate Examples:
– Chlorous (from chlorite, ClO2-)
– Nitric (from nitrate, NO3-)
– Sulfurous (from sulfite, SO3-2)
pH and [H3O+]
pH: number that is derived from the concentration of hydronium ions ([H3O+]) in sol’n– pH = -log [H3O+]– As pH increases, [H3O+] decreases
Scale ranges from 0 – 14– pH = 7 is neutral– pH < 7 is acidic– pH > 7 is basic
p[OH]
pOH = - log [OH-] pH + pOH = 14.00 Calculating ion concentrations from pH
[H+] = antilog (-pH) [OH-] = antilog (-pOH)
Dissociation Constants
Acid dissociation constant: (Ka): the equilibrium constant for the rxn of an aqueous weak acid and water
Base dissociation constant: (Kb): the equilibrium constant for the rxn of an aqueous weak base w/ water
Both are derived from the ratio of the concentration of the products and reactants at equilibrium
Acid Dissociation Constant
Ka = [H3O+] [A-]
[HA] Ka is a measure of the strength of an acid
Ka values for weak acids are always less than one
Used mostly w/ weak acids because the Ka values for strong acids approach infinity
Examples
HMnO4 (aq) + H2O (l)
H2S (aq) + H2O (l)
Base Dissociation Constant
Kb = [HB+] [OH-][B]
Kb is a measure of the strength of a base
Kb values for weak bases are always less than 1
Kb values for strong bases approach infinity
Examples
H2NOH (aq) + H2O (l)
NH3 (aq) + H2O (l)
Water
Water can dissociate into its component ions, H+ and OH-
– 2H2O (l) H3O+ (aq) + OH- (aq) One water molecule acts as a weak acid, and
the other acts as a weak base The ions are present in such small amounts
they can’t be detected by a conductivity apparatus
In pure water, [H3O+] =1.0 x 10 –7 M and [OH-] = 1.0 x 10-7 M
Dissociation Constant for Water
It is defined as Kw: the ion product constant for water
Kw = [H3O+] [OH-] Kw = (1.0 x 10-7)(1.0 x 10-7) Kw = 1.0 x 10-14
The value of Kw can always be used to find the concentration of either H3O+ or OH- given the concentration of the other
Examples
What is the pH of a 0.001 M sol’n of HCl, a strong acid?
Examples
What is the pH of a sol’n if [H3O+] = 3.4 x 10-5 M?
Examples
The pH of a sol’n is measured with a pH meter and determined to be 9.00. What is the [H3O+]?
Examples
The pH o f a sol’n is measured with a pH meter and determined to be 7.52. What is [H3O+]?
Calculating Ka
In these problems, remember that the concentration of the [H3O+] ions will equal the concentration of the conjugate base ions.– This is because for every molecule of
weak acid that dissociates, there will be an equal number of H3O+ ions and base ions
Example
Assume that enough lactic acid is dissolved in sour milk to give a solution concentration of 0.100 M lactic acid. A pH meter shows that the pH of the sour milk is 2.43. Calculate Ka for the lactic acid equilibrium system.
Titrations
An analytical procedure used to determine the concentration of a sample by reacting it with a standard sol’n
In a titration, an indicator is used to determine the end point
Standard sol’n: a sol’n of precisely known concentration
Indicator: any substance in sol’n that changes color as it reacts with either an acid or a base
Titrations
Each indicator changes its color over a particular range of pH values (transition interval)
An unknown acid sol’n will be titrated with a standard sol’n that is a strong base
An unknown base sol’n will be titrated with a standard sol’n that is a strong acid
Titrations
Equivalence point: point at which the concentration of H3O+ ions is the same as the concentration of OH- ions; [H3O+ ] = [OH-]
Endpoint: the point at which the indicator changes color
Titration curve: graph that shows how pH changes in a titration
Titrations
The equivalence point is at the center of the steep, vertical region of the titration curve
At the equivalence point, pH increases greatly w/ only a few drops
Example Problem 1
What is the molarity of a CsOH solution if 20.0 mL of the solution is neutralized by 26.4 mL of 0.250M HBr solution?
HBr + CsOH → H2O + CsBr
Example Problem 2
What is the molarity of a nitric acid solution if 43.33 mL 0.200M KOH solution is needed to neutralize 20.00 mL of unknown solution?
Example Problem 3
What is the concentration of a household ammonia cleaning solution (NH4OH) if 49.90 mL of 0.5900M H2SO4 is required to neutralize 25.00 mL solution?