ap chapter 16 acid - base equilibria
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Acid - Base Equilibria
AP Chapter 16
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Acids and Bases
Arrhenius acids have properties that are due to the presence of the hydronium ion (H+ (aq)) They turn litmus red.
Arrhenius bases have properties that are due to the presence of the hydroxide ion (OH- (aq)). They turn litmus blue.
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Brønsted-Lowery Acids and Bases
Bronsted-Lowery definitions state that acid-base reactions involve the transfer of hydronium ions (H+) from one substance to another.
A hydronium ion is simply a proton with no surrounding valence electrons.
Remember – acids donate and bases accept!
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Hydronium ion
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Proton transfer reactions
The polar water molecule promotes the ionization of acids in water solution by accepting a proton to form H3O+.
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Brønsted-Lowery definitions
ACID
BASEAcids donate protons
Bases accept protons
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The Hydronium ion
The hydronium ion (H3O+) is a hydrated proton. When water accepts a proton from an acid, the product is a hydronium ion.
Hydronium ions are represented by either the H3O+(aq) or H+(aq)
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Conjugate Acid-Base Pairs
Conjugate acid-base pairs are two substances in an aqueous solution whose formulas differ by an H+.
The acid is the more positive species having an extra H.
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Amphoteric
An amphoteric substance is a substance that can act as either an acid or a base.
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Relative Strengths of Acids and Bases
A strong acid completely transfers its protons to water, leaving no undissociated molecules in water. It totally dissociates in water.
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Relative Strengths, continued
A weak acid only partially dissociates in water, and exists as a mixture of acid molecules and their constituent ions.
The conjugate base of a weak acid is a weak base.
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Relative Strengths, continued
A substance with negligible acidity, such as CH4, contains hydrogen, but does not demonstrate any acidic behavior in water.
It’s conjugate base is a strong base, reacting completely with water.
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Acid-Base Equilibrium
In every acid-base reaction the position of equilibrium favors transfer of the proton from the stronger acid to the stronger base to form the weaker acid and the weaker base.
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Autoionization of Water
Water has the ability to act as either an acid or a base.
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The ion product of water
Because the autoionization of water is an equilibrium process, there is an equilibrium-constant expression:
Kc = [H3O+][OH-]
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Ion-Product Constant for H2O
Kw = [H3O+][OH-] = 1.0 x 10 -14
can also be written
= [H+][OH-] = 1.0 x 10 -14
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Acid and Base Ionization Constants
The acid ionization constant (Ka) is the equilibrium constant for the ionization of a weak acid in water.
The base ionization constant (Kb) is the equilibrium constant for a weak base.
For any conjugate acid-base pair,
Kw = Ka x Kb.
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pH Scale
pH = -log[H+] Neutral solution:
pH = -log(1.0 x 10-7) = -(-7.00) = 7.00
The pH decreases as the [H+] increases.
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The pH Scale
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The pH Scale
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Calculating the pH of a Basic Solution
Calculate the pH of a basic solution, where the [OH-] > 1.0 x 10-7 M. Suppose [OH-] = 2.0 x10-3 M. Calculate the H+ value for this solution.
[H+] = = = 5.0 x 10-12 M Kw 1.0 x 10-14
[OH-] 2.0 x 10-3
pH = -log(5.0 x 10-12) = 11.30
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pH and pOH
pH and pOH = 14.00
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Measuring pH
pH meter Acid-base indicators (less precise)
Methyl orangeLitmusphenolphthaleinEtc.
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Strong Acids and Bases
The seven most common strong acids include 6 monoprotic acids (HCl, HBr, HI, HNO3, HClO3, and HClO4) and one diprotic acid, H2SO4.
HNO3(aq) + HOH(l) → H3O+(aq) + NO3-(aq)
HNO3(aq) → H+(aq) + NO3-(aq)
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Calculating the pH of a strong acid
What is the pH of a 0.040 M solution of HClO4?
pH = -log(0.040) = 1.40
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Strong Bases
Common strong bases are the ionic hydroxides of alkali metals and the heavy alkaline earth metals.
The cations of these metals have negligible acidity.
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Weak Acids
Weak acids are only partially ionized (or dissociated.)
They are weak electrolytes. HA(aq) + HOH(l) ↔ H3O+(aq) + A-(aq)
Ka =
The larger the value of Ka, the stronger the acid.
[H3O+][A-]
[HA]
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Polyprotic Acids
Polyprotic acids have more than one ionizable proton, such as H2SO3.
These acids have acid-dissociation constants that decrease in magnitude in the order Ka1>Ka2>Ka3.
Because nearly all the H+(aq) in a polyprotic solution comes from the first dissociation, the pH can usually be estimated using only Ka1.
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Weak Bases
Weak bases include NH3, amines and the anions of weak acids.
Kb = the dissociation constant for the base.
The relationship between the strength of an acid and the strength of its conjugate base is expressed by the equation Ka x Kb = Kw
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Using Kb to Calculate OH-
NH3(aq) + HOH(l) ↔ NH4+(aq) + OH-(aq)
Kb = [NH4
+][OH-]
[NH3]= 1.8 x 10-5
Reference the problem example on page 691.
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Hydrolysis
Acid-base properties of salts can be attributed to the behavior of their respective cations and anions.
The reaction with water, with a resulting change in pH, is called hydrolysis.
Cations of alkali metals and alkaline earth metals and anions of strong acids don’t hydrolyze.
Salt + water = acid + base
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Acid-Base Behavior and Chemical Structure
A molecule containing H will transfer a proton only of the H-X bond is polarized:
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Bond Strength
Strong bonds do not dissociate as easily as weaker bonds, so they are less likely to form acidic ions in solution.
Since HF has such a strong bond due to the electronegativities, it does not dissociate readily and is therefore a weak acid.
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Oxyacids
Oxyacids are acids in which OH groups and possible additional oxygen atoms are present.
What determines whether it is an acid or a base?
Generally, as the electronegativity of the attached element increases, so will the acidity of the substance.
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Oxyacids
The strength of an acid will increase as additional electronegative atoms bond to the central atom.
Electronegativity
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Acid strength increases as the number of oxygen atoms attached to the central atom increases.
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Carboxylic Acids Acids that contain carboxyl groups are
called carboxylic acids.
These form the largest category of
organic acids.
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Lewis Acids and Bases The Lewis concept of acids emphasizes
the shared electron pair rather than the proton.
A Lewis acid is an electron-pair acceptor.
A Lewis base is an electron-pair donor. This concept is more general than the
Brønsted-Lowery definition – it explains why many hydrated metal cations can form acidic aqueous solutions.
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Acidity of Metal Cations The acidity of a hydrated metal cation
depends on the cation charge and size.