mark s. cracolice edward i. peters mark s. cracolice the university of montana chapter 17...
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Mark S. CracoliceEdward I. Peters
Mark S. Cracolice • The University of Montana
www.cengage.com/chemistry/cracolice
Chapter 17Acid–Base (Proton Transfer) Reactions
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Arrhenius Acid–Base Theory
An acid ( HCl) is a substance which produces hydrogen ions in water solution. The properties of an acid is the properties of the hydrogen ions.
A base (NaOH) is a substance which produces hydroxide ions in water solution. The properties of a base are the properties of the hydroxide ions.
The net reaction between a strong acid and a strong base is :
H+ + OH- H2O
The Arrhenius concept of acids and base is limited because it applies only to aqueous solutions.
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Arrhenius Acid–Base Theory
An example of an Arrhenius acid:
Gaseous hydrogen chloride dissolved in water:
HCl(g) H+(aq) + Cl–(aq)
An example of an Arrhenius base:
Solid sodium hydroxide dissolved in water:
NaOH(s) Na+(aq) + OH–(aq)
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Arrhenius Acid–Base Theory
Properties of an acid must be due to properties of H+.
Properties of a base must be due to properties of OH–.
Thus the cause of the sour taste of acids is the H+ ion, the cause of the bitter taste of bases is the OH– ion.
Other characteristic properties of acids and bases arealso due to the H+ and OH– ions in water solutions.
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Arrhenius Acid–Base Theory
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Brønsted–Lowry Acid–Base Theory
An acid is a proton donor.
A base is a proton acceptor.
An acid-base reaction is a proton-transfer reaction in which the proton is transferred from the acid to a
base with formation of another acid and base.
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Brønsted–Lowry Acid–Base Theory
The base formed when the acid has donated a proton is called the conjugate base of the acid.
Acid A ↔ H+ + Conjugate base of acid A
The sign ↔ is used to show that the reaction is reversible.
The stronger the acid, the weaker the conjugate base, and the weaker the acid, the stronger the conjugate base.
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Brønsted–Lowry Acid–Base Theory
An acid base reaction is a proton transfer reaction in which a proton is transferred from a stronger acid to a stronger base with formation of a weaker acid and weaker base.
Stronger Acid1+ Stronger Base2↔Weaker Acid2+ Weaker Base1
HNO3 + NH3 ↔ NH4+ + NO3
-
HCl + CH3COO- ↔ HCH3COO + Cl-
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Brønsted–Lowry Acid–Base Theory
H3O+ is called hydronium ion. The conjugate base of acid HNO3 is NO3
-
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Brønsted–Lowry Acid–Base Theory
Water which can behave as a base in one case and an acid in another is said to be amphoteric.
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Brønsted–Lowry Acid–Base Theory
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Conjugate Acid–Base Pairs
B + HA HB+ + A–
base acid acid base
proton proton proton proton
remover source source remover
Conjugate Acid–Base Pair
Two species that transform into each other by
gain or loss of a proton, H+.
HB+ and B and HA and A– are conjugate acid–base pairs
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Conjugate Acid–Base Pairs
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Conjugate Acid–Base Pairs
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Lewis Acid–Base Theory
Lewis Theory of Acids and Bases
AcidElectron-pair acceptor.
BaseElectron-pair donor.
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Lewis Acid–Base Theory
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Relative Strengths of Acids & Bases
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Predicting Acid–Base Reactions
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The Water Equilibrium
Autoionization of water
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The Water Equilibrium
H2O(l) H+(aq) + OH–(aq)
Kw = [H+] [OH–] = 1.0 × 10–14
Kw is the water constant or equilibrium constant for water
If [H+] = [OH–] = x
Kw = [H+] [OH–] = 1.0 × 10–14 = x2
x = = 10–7 moles/liter
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The Water Equilibrium
For water or water solutions:
If [H+] = [OH–] = 10–7 M,
the solution is neutral.
If [H+] > [OH–],
the solution is acidic.
If [H+] < [OH–],
the solution is basic.
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The Water Equilibrium
Example:
What is the hydrogen ion concentration in a solution of 10–4 M sodium hydroxide in which the hydroxide ion concentration is 10–4 M? Is the solution acidic or basic?
Solution:
GIVEN: [OH–] = 10–4 M WANTED: [H+]
EQUATION: Kw = [H+] [OH–] = 10–14
Since [H+] = 10–10 M< [OH–] = 10–4 M, the solution is basic
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pH and pOH
By definition pH and pOH are given by
pH ≡ -log [H3O+]
pOH ≡ -log [OH-]
[H3O+] ≡ antilog(-pH) ≡ 10-pH
[OH-] ≡ antilog(-pOH) ≡ 10-pOH
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pH and pOH
What is the pH of a solution with [H+] = 10–5 M?
Solution:
pH = – log [H+] = – log 10–5 = 5
What is the [OH–] of a solution with pOH = 6?
Solution:
[OH–] = antilog (–pOH) = antilog (–6) = 10–6 M
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pH and pOH
Kw = [H+] [OH–] = 1.0 × 10–14
[H+] [OH–] = 1.0 × 10–14
– log ([H+] [OH–]) = – log (1.0 × 10–14)
– log ([H+] [OH–]) = 14
– log [H+] + (– log [OH–]) = 14
pH + pOH = 14
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pH and pOH
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pH and pOH
Example:
The hydrogen ion concentration of a 10–3 M HCl solution is
10–3 M. What are the pH, pOH, and [OH–] of the solution?
Solution:
pH = – log [H+] = – log 10–3 = 3
pH + pOH = 14
pOH = 14 – pH = 14 – 3 = 11
[OH–] = antilog (–pOH) = antilog (–11) = 10–11 M
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pH and pOH
A solution is neutral if [H+] = 10–7 M
A solution is acidic if [H+] > 10–7 M
A solution is basic if [H+] < 10–7 M
Using pH = – log [H+] and pOH = – log [OH–],
A solution is neutral if pH = 7
A solution is acidic if pH < 7
A solution is basic if pH > 7
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Significant Figures and Logarithms
In a logarithm, the digits to the left of the decimal are not counted as significant figures.
Counting significant figures in a logarithm begins at the decimal point.
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pH and pOH
Example:
The hydrogen ion concentration of a solution is 2.7 × 10–6 M. What are the pH, pOH, and hydroxide ion concentration?
Solution:
pH = – log [H+] = – log (2.7 × 10–6) = – log (10–6) – log (2.7) =
= 6 – log (2.7) = 5.57 (2 significant figures)
pH + pOH = 14.00
pOH = 14.00 – pH = 14.00 – 5.57 = 8.43
[OH–] = antilog(–pOH) = antilog(–8.43) = 10–8.43 M =3.7 × 10–9 M
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pH and pOH
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pH and pOH
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pH and pOH
Measurement of pH
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HOMEWORK
15, 17, 21, 23, 39, 41, 55, 59, 64.