Transcript Acids and Bases

Acids and Bases
Acids and Bases
• Acid: a substance that, when dissolved in
water, increases the concentration of H+
(protons)
H2O
• HCl
H+ + Cl• HCl + H2O
H3O+ + Cl-
• Strong acid: an acid that completely
ionizes/dissociates in water
• HNO3 + H2O
H3O+ + NO3• Strong electrolyte
• Strong Acids:
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•
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•
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Hydrochloric Acid
Nitric Acid
Sulfuric Acid
Perchloric Acid
Hydrobromic Acid
Hydroiodic Acid
HCl
HNO3
H2SO4
HClO4
HBr
HI
• Weak Acid: an acid that only partially
ionizes/dissociates in water
• CH3COOH(aq)
• Weak electrolytes
• Weak Acids:
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•
•
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Phosphoric acid
Acetic Acid
Carbonic Acid
Hydrocyanic Acid
Benzoic Acid
CH3COO-(aq) + H3O+(aq)
H3PO4
CH3COOH
H2CO3
HCN
C6H5COOH
Polyprotic Acids
• Polyprotic acids: acids that can release
more than one H+
• Sulfuric Acid
• H2SO4(aq)
• HSO4-(aq)
HSO4-(aq) + H3O+(aq)
SO42-(aq) + H3O+(aq)
Problem
• Write out the equation(s) for the ionization
of phosphoric acid, H3PO4, in water
Bases
• Base: a substance that, when put in water,
increases the concentration of OH- ions or
a substance that accepts H+ ions
• NaOH(aq)  Na+(aq) + OH-(aq)
• Strong Bases: bases that completely
ionize in water
• Weak Bases: bases that only partially
ionize in water
• Ammonia: NH3
• Pyridine: C5H5N
Congugate Acid-Base Pairs
Each acid has a conjugate base and every base
has a conjugate acid
conjugate acid-base pair 1
HA
+B
A− + BH+
conjugate acid-base pair 2
Problems
Identify the conjugate bases for the following:
1) HBr
2) H2S
3) H2CO3
Identify the conjugate acids for the following
1) NO22) NH3
3) OH-
In the following equations, identify the
conjugate acid and base pairs:
1) HCl + H2O  H3O+ + Cl-
2) C5H5N + H2O  C5H6N+ + OH-
Neutralization Reactions
• When strong acids and bases in aqueous
solution react with each other, they form
water and a salt
• HX(aq) + MOH(aq)  HOH(l) + MX(aq)
Water
Salt
• HCl(aq) + NaCl(aq)  H2O(l) + NaCl(aq)
Problems
Write out the molecular, total ionic, and net
ionic equations for the reactions of the
following acids and bases
1) HBr and KOH
2) HCl and Mg(OH)2
3) H2SO3 and NaOH
Acid Ionization Constant
• Acid Ionization Constant (Ka): the
equilibrium constant for the ionization
reaction of an acid with water
• HA + H2O
A- + H3O+
• Large Ka = Strong acid
• Small Ka = Weak acid
Questions
1) Write the equilibrium constant expression
for the dissociation of HCN (hydrocyanic
acid) in water
2) Write the equilibrium expression for the
dissociation of HF in water
3) If the Ka for HCN is 4.9 x 10-10 and the Ka
for HF is 7.2 x 10-4, which acid is
stronger?
4) Predict whether the equilibrium for each of
the following reactions favors the reactants
or products.
a) H3PO4(aq) + H2O(l)
H3O+(aq) + H2PO4(aq)
b) NH4+(aq) + H2O(l)
H3O+(aq) + NH3(aq)
Base Ionization Constant
• Base Ionization Constant (Kb): the
equilibrium constant for the ionization
reaction of a base with water
• B + H2O
OH- + BH+
• Large Kb = Strong base
• Small Kb = Weak base
Autoionization of Water
Autoionization of Water
• Also called “Self Ionization”
• About 1 out of every 10 million water
molecules form ions through self ionization
H2O  H+ + OH–
H2O + H2O  H3O+ + OH–
• All aqueous solutions contain both H3O+ and OH–
Ion Product Constant for Water
• Ion Product Constant for Water (Kw): the
numerical value obtained by multiplying
the molar concentrations for hydronium
and hydroxide ions present in pure water
at 25°C
• Kw = [H3O+][OH-] = 1.00 x 10-14 at 25 oC
• the concentration of H3O+ and OH– are equal
in pure water
• [H3O+] = [OH–] = 10-7M @ 25°C
Ion Product of Water
• the product of the H3O+ and OH–
concentrations is always the same
number
• Kw =[H3O+][OH–] = 1.00 x 10-14 @
25°C
– if you measure one of the concentrations,
you can calculate the other
• as [H3O+] increases the [OH–] must
decrease so the product stays constant
– inversely proportional
[H+] vs. [OH-]
[H+] 100 10-1
+
H
OH-
Acid
10-3
10-5
+
H
OH-
[OH-]10-14 10-13 10-11
10-9
10-7
10-9
Base
10-11
H+
+
H
10-13 10-14
H+
OH OH OH
10-7
10-5
10-3
10-1 100
Even though it may look like it, neither H+ nor OH- will ever be 0
Acidic and Basic Solutions
• Neutral solutions have equal [H3O+] and [OH–]
• [H3O+] = [OH–] = 1 x 10-7
• acidic solutions have a larger [H3O+] than [OH–]
• [H3O+] > [OH–]
• [H3O+] > 1 x 10-7; [OH–] < 1 x 10-7
• basic solutions have a larger [OH–] than [H3O+]
• [H3O+] < [OH–]
• [H3O+] < 1 x 10-7; [OH–] > 1 x 10-7
Questions
1) Calculate the [OH] at 25°C when the [H3O+]
= 1.5 x 10-9 M, and determine if the solution is
acidic, basic, or neutral
2) Calculate the [H3O+] at 25°C in a solution that
has a [OH] 1.3 x 10-10 M and determine if the
solution is acidic, basic, or neutral
3) Calculate the [H3O+] at 25°C in a solution that
has a [OH] 1.0 x 10-7 M and determine if the
solution is acidic, basic, or neutral
• Acidic solutions
• pH < 7
pH and pOH
• Basic solutions
• pH > 7
• Neutral solutions
• pH = 7
• pH is a measure of the concentration of H+
or H3O+ in solution
• pH = -log [H+] OR pH = -log [H3O+]
• [H3O+] = 10-pH
• pOH is a measure of the concentration of
OH- in solution
• pOH = -log [OH-]
• [OH-] = 10-pOH
• pH + pOH = 14
Problems
1) What is the pH for an HCl solution with
an [H3O+] = 2.3 x 10-2?
2) What is the pH of an NaOH solution with
an [H3O+] = 3.39 x 10-10?
3) What is the pH of a KOH solution with a
[OH-] = 1.1 x 10-2?
4) What is the [H3O+] for a solution with a
pH of 8.45. What is the [OH-]?
Classification of Water Soluble
Substances
• Electrolytes: solutes that separate into
ions when dissolved in water (they’re
soluble)
– Have the ability to conduct electricity
– 2 types
• Strong electrolytes
• Weak electrolytes
• Strong electrolytes:
solutes that completely
dissociates into ions
when dissolved in
water
– Examples: NaCl,
MgBr2, HCl
– Strong electrical
conductors
– Strong electrolyte(aq or s) → Cation+(aq) + Anion-(aq)
– Example: NaCl(s) → Na+(aq) + Cl-(aq)
• Weak Electrolytes:
solutes that, when
dissolved in water,
only partially
dissociates into ions
– Examples: HF, NH3,
acetic acid
– Weak electrical
conductors
– Weak electrolyte(aq) ↔ Cation+(aq) + Anion-(aq)
– Example: HF(aq) ↔ H+(aq) + F-(aq)
• Nonelectrolytes: solutes
that dissolve in water
without separating into
ions
– Examples: sucrose,
ethanol
– Do not conduct electricity
Nonelectrolyte (s or l) → Nonelectrolyte(aq)
Example: C12H22O11(s) → C12H22O11(aq)
Problems
1) The following salts are strong electrolytes.
Write a balanced equation for their dissociation
in water
a) LiBr
b) FeCl3
2) HCN is a weak acid. Write a balanced
equation for its dissociation in water
3) Do you expect Fructose (C6H12O6) to be a
strong electrolyte, weak electrolyte, or
nonelectrolyte? Write a balanced equation for
its solvation in water
Titration
• Titration: a procedure for the
quantitative analysis of a
substance of unknown
concentration whereby a
measured quantity of another
substance, of know
concentration, is completely
reacted with the with the original
substance.
– Often used to determine the
concentration of acids and bases
• Equivalence point: the point in a titration at which
one reactant has been exactly consumed by the by
the addition of another reactant
– Midpoint of vertical rise
– Occurs at pH = 7 in a strong acid-strong base titration
– [H3O+] = [OH-]
Indicators
• Acid-Base Indicator: a chemical that
changes color with a change in pH
– Added to solutions in small amounts in order
to determine to solution’s pH visually
– Usually organic compounds
– Weak acid or base
– establishes an equilibrium with the H2O and
H3O+ in the solution
HInd(aq) + H2O(l)  Ind(aq) + H3O+(aq)
Phenolphthalein
53
Bromocresol Green
Yellow
Green
Blue
Methyl Red
Questions
1) You have 25.0 mL of a HCl solution of
unknown concentration. If you titrate your
acid with 17.3 mL of a 0.877 M NaOH
solution, what is the concentration of
your acid?
2) You also have 25.0 mL of a sulfuric acid
solution. If you titrate this solution with
32.1 mL of a 1.50 M KOH solution, what
is the concentration of your acid?