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Chapter 17 Acid–Base (Proton Transfer) Reactions

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Presentation on theme: "Chapter 17 Acid–Base (Proton Transfer) Reactions"— Presentation transcript:

1 Chapter 17 Acid–Base (Proton Transfer) Reactions

2 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.

3 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)

4 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 are also due to the H+ and OH– ions in water solutions.

5 Arrhenius Acid–Base Theory

6 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.

7 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.

8 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 HNO NH ↔ NH NO3- HCl CH3COO- ↔ HCH3COO Cl-

9 Brønsted–Lowry Acid–Base Theory
H3O+ is called hydronium ion. The conjugate base of acid HNO3 is NO3-

10 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.

11 Brønsted–Lowry Acid–Base Theory

12 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

13 Conjugate Acid–Base Pairs

14 Conjugate Acid–Base Pairs

15 Lewis Acid–Base Theory
Lewis Theory of Acids and Bases Acid Electron-pair acceptor. Base Electron-pair donor.

16 Lewis Acid–Base Theory

17 Relative Strengths of Acids & Bases

18 Predicting Acid–Base Reactions

19 Autoionization of water
The Water Equilibrium Autoionization of water

20 Kw is the water constant or equilibrium constant for water
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

21 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.

22 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

23 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

24 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? [OH–] = antilog (–pOH) = antilog (–6) = 10–6 M

25 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

26 pH and pOH

27 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

28 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

29 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.

30 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 = – pH = – 5.57 = 8.43 [OH–] = antilog(–pOH) = antilog(–8.43) = 10–8.43 M =3.7 × 10–9 M

31 pH and pOH

32 pH and pOH

33 pH and pOH Measurement of pH

34 Homework 15, 17, 21, 23, 39, 41, 55, 59, 64.


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