Acids and Bases. Properties of Acids Sour taste React w/ metals to form H 2 Most contain hydrogen...
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Transcript of Acids and Bases. Properties of Acids Sour taste React w/ metals to form H 2 Most contain hydrogen...
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Acids and Bases
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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
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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
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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
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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
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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
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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)
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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)
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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
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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-
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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
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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
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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
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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
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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)
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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)
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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
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p[OH]
pOH = - log [OH-] pH + pOH = 14.00 Calculating ion concentrations from pH
[H+] = antilog (-pH) [OH-] = antilog (-pOH)
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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
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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
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Examples
HMnO4 (aq) + H2O (l)
H2S (aq) + H2O (l)
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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
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Examples
H2NOH (aq) + H2O (l)
NH3 (aq) + H2O (l)
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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
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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
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Examples
What is the pH of a 0.001 M sol’n of HCl, a strong acid?
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Examples
What is the pH of a sol’n if [H3O+] = 3.4 x 10-5 M?
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Examples
The pH of a sol’n is measured with a pH meter and determined to be 9.00. What is the [H3O+]?
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Examples
The pH o f a sol’n is measured with a pH meter and determined to be 7.52. What is [H3O+]?
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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
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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.
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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
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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
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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
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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
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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
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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?
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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?