Acids and Bases Chapter 15

Acids and Bases Chapter 15
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

Contents and Concepts Acid–Base Concepts Arrhenius Concept of Acids and Bases Brønsted–Lowry Concept of Acids and Bases Lewis Concept of Acids and Bases Acid and Base Strengths Relative Strengths of Acids and Bases Molecular Structure and Acid Strength

Self-Ionization of Water and pH Self-Ionization of Water
Solutions of a Strong Acid or Base The pH of a Solution I was told not to use my time to do this.

Learning Objectives Acid Base Concepts Arrhenius Concept of Acids and Base a. Define acid and base according to the Arrhenius concept. Brønsted–Lowry Concept of Acids and Bases a. Define acid and base according to the Brønsted–Lowry concept. b. Define the term conjugate acid–base pair. c. Identify acid and base species. d. Define amphiprotic species.

3. Lewis Concept of Acids and Bases
a. Define Lewis acid and Lewis base. b. Identify Lewis acid and Lewis base species. Acid and Base Strengths 4. Relative Strengths of Acids and Bases a. Understand the relationship between the strength of an acid and that of its conjugate base. b. Decide whether reactants or products are favored in an acid–base reaction.

5. Molecular Structure and Acid Strength
a. Note the two factors that determine relative acid strengths. b. Understand the periodic trends in the strengths of the binary acids HX. c. Understand the rules for determining the relative strengths of oxoacids. d. Understand the relative acid strengths of a polyprotic acid and its anions.

Self-Ionization of Water and pH
a. Define self-ionization (or autoionization). b. Define the ion-product constant for water. 7. Solutions of a Strong Acid or Base a. Calculate the concentrations of H3O+ and OH- in solutions of a strong acid or base.

8. The pH of a Solution a. Define pH. b. Calculate the pH from the hydronium-ion concentration. c. Calculate the hydronium-ion concentration from the pH. d. Describe the determination of pH by a pH meter and by acid–base indicators.

HCl(g) + NH3(g)  NH4Cl(s)
When gaseous hydrogen chloride meets gaseous ammonia, a smoke composed of ammonium chloride is formed. HCl(g) + NH3(g)  NH4Cl(s) This is an acid–base reaction.

Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases.

We will examine three ways to explain acid–base behavior:
Arrhenius Concept Brønsted–Lowry Concept Lewis Concept H+ and OH− donor acceptor H+ = proton donor acceptor electron pair acid base Note: H+ in water is H3O+

Arrhenius Concept of Acids and Bases
An Arrhenius acid is a substance that, when dissolved in water, increases the concentration of hydronium ion, H3O+(aq). An Arrhenius base is a substance that, when dissolved in water, increases the concentration of hydroxide ion, OH-(aq).

The Arrhenius concept limits bases to compounds that contain a hydroxide ion.
The Brønsted–Lowry concept expands the compounds that can be considered acids and bases.

Brønsted–Lowry Concept of Acids and Bases
An acid–base reaction is considered a proton (H+) transfer reaction. H+ H+ H+ H+

A Brønsted–Lowry acid is the species donating a proton in a proton-transfer reaction; it is a proton donor. A Brønsted–Lowry base is the species accepting a proton in a proton-transfer reaction; it is a proton acceptor.

Substances in the acid–base reaction that differ by the gain or loss of a proton, H+, are called a conjugate acid–base pair. The acid is called the conjugate acid; the base is called a conjugate base. Acid Base Conjugate base Conjugate acid

A Brønsted acid is a proton donor A Brønsted base is a proton acceptor
conjugate acid conjugate base base acid

What is the conjugate acid of H2O? What is the conjugate base of H2O?
The conjugate acid of H2O has gained a proton. It is H3O+. The conjugate base of H2O has lost a proton. It is OH-.

a. HCO3−(aq) + HF(aq) H2CO3(aq) + F−(aq)
Label each species as an acid or base. Identify the conjugate acid-base pairs. a. HCO3−(aq) + HF(aq) H2CO3(aq) + F−(aq) Base Acid Conjugate acid Conjugate base b. HCO3−(aq) + OH−(aq) CO32−(aq) + H2O(l) Acid Base Conjugate base Conjugate acid

Species that can act as both an acid and a base are called amphiprotic or amphoteric species.
Identify any amphiprotic species in the previous problem. HCO3− was a base in the first reaction and an acid in the second reaction. It is amphiprotic.

Lewis Concept of Acids and Bases
A Lewis acid is a species that can form a covalent bond by accepting an electron pair from another species. A Lewis base is a species that can form a covalent bond by donating an electron pair to another species.

No protons donated or accepted!
Lewis Acids and Bases F B F N H • H F B F N H H + acid base No protons donated or accepted!

[ ] - Acid-Base Properties of Water autoionization of water
H2O (l) H+ (aq) + OH- (aq) autoionization of water O H + - [ ] conjugate acid base H2O + H2O H3O+ + OH- acid conjugate base

The Ion Product of Water
Kc = [H+][OH-] [H2O] H2O (l) H+ (aq) + OH- (aq) [H2O] = constant Kc[H2O] = Kw = [H+][OH-] The ion-product constant (Kw) is the product of the molar concentrations of H+ and OH- ions at a particular temperature. Solution Is [H+] = [OH-] neutral At 250C Kw = [H+][OH-] = 1.0 x 10-14 [H+] > [OH-] acidic [H+] < [OH-] basic

What is the concentration of OH- ions in a HCl solution whose hydrogen ion concentration is 1.3 M?
Kw = [H+][OH-] = 1.0 x 10-14 [H+] = 1.3 M [OH-] = Kw [H+] 1 x 10-14 1.3 = = 7.7 x M

Solutions of a Strong Acid or Strong Base
The concentration of hydronium or hydroxide in a solution of strong acid or base is related to the stoichiometry of the acid or base.

Calculate the hydronium and hydroxide ion concentration at 25°C in
a M HCl b. 1.4 × 10−4 M Mg(OH)2 When HCl ionizes, it gives H+ and Cl−. So [H+] = [Cl−] = [HCl] = 0.10 M. When Mg(OH)2 ionizes, it gives Mg2+ and 2 OH−. So [OH−] = 2[Mg2+] = 2[Mg(OH)2] = 2.8 × 10−4 M.

Solutions can be characterized as
Acidic: [H3O+] > 1.0 × 10−7 M Neutral: [H3O+] = 1.0 × 10−7 M Basic: [H3O+] < 1.0 × 10−7 M

pH – A Measure of Acidity
pH = -log [H+] Solution Is At 250C neutral [H+] = [OH-] [H+] = 1 x 10-7 pH = 7 acidic [H+] > [OH-] [H+] > 1 x 10-7 pH < 7 basic [H+] < [OH-] [H+] < 1 x 10-7 pH > 7 pH [H+]

Other important relationships
pOH = -log [OH-] [H+][OH-] = Kw = 1.0 x 10-14 -log [H+] – log [OH-] = 14.00 pH + pOH = 14.00 pH Meter

The pH of a Solution pH = –log[H3O+]
For a log, only the decimal part of the number has significant figures. The whole number part, called the characteristic, is not significant.

Calculate the pH of typical adult blood, which has a hydronium ion concentration of 4.0 × 10−8 M.
[H3O+] = 4.0 × 10−8 M pH = –log [H3O+] pH = – log (4.0 × 10−8) = – (– 7.40) pH = 7.40 Note: The two significant figures are the two decimal places.

pOH = –log[OH−] pH + pOH = (at 25°C)

[H3O+] = 10−pH [OH−] = 10−pOH

The pH of natural rain in 5. 60
The pH of natural rain in What is its hydronium ion concentration? pH = 5.60 [H3O+] = 10−pH = 10−5.60 [H3O+] = 2.5 × 10-6 M Because the pH has two decimal places, the concentration can have only two significant figures.

The next slide summarizes the conversions involving H3O+, OH−, pH, and pOH.
Note that you can only go around the edges of the square; it takes two steps to go from one corner to the opposite corner.

[H3O+] [OH−] pH pOH

The pH of rainwater collected in a certain region of the northeastern United States on a particular day was What is the H+ ion concentration of the rainwater? pH = -log [H+] [H+] = 10-pH = = 1.5 x 10-5 M The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood? pH + pOH = 14.00 pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60 pH = – pOH = – 6.60 = 7.40

Strong Electrolyte – 100% dissociation
NaCl (s) Na+ (aq) + Cl- (aq) H2O Weak Electrolyte – not completely dissociated CH3COOH CH3COO- (aq) + H+ (aq) Strong Acids are strong electrolytes HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq) HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq) H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq)

Weak Acids are weak electrolytes
HF (aq) + H2O (l) H3O+ (aq) + F- (aq) HNO2 (aq) + H2O (l) H3O+ (aq) + NO2- (aq) HSO4- (aq) + H2O (l) H3O+ (aq) + SO42- (aq) H2O (l) + H2O (l) H3O+ (aq) + OH- (aq) Strong Bases are strong electrolytes NaOH (s) Na+ (aq) + OH- (aq) H2O KOH (s) K+ (aq) + OH- (aq) H2O Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq) H2O

Weak Bases are weak electrolytes
F- (aq) + H2O (l) OH- (aq) + HF (aq) NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq) Conjugate acid-base pairs: The conjugate base of a strong acid has no measurable strength. H3O+ is the strongest acid that can exist in aqueous solution. The OH- ion is the strongest base that can exist in aqeous solution.

Relative Strengths of Acids and Bases
The stronger an acid, the weaker its conjugate base. The weaker an acid, the stronger its conjugate base.

Formic acid, HCHO2, is a stronger acid than acetic acid, HC2H3O2
Formic acid, HCHO2, is a stronger acid than acetic acid, HC2H3O2. Which is the stronger base: formate ion, CHO2−, or acetate ion, C2H3O2−? Because formic acid is stronger than acetic acid, formate ion (which is the conjugate base of formic acid) will be a weaker base than acetate ion (which is the conjugate base of acetic acid). The acetate ion is a stronger base than the formate ion.

Molecular Structure and Acid Strength
H X H+ + X- The stronger the bond The weaker the acid HF << HCl < HBr < HI acidity increases

The strength of an acid depends on how easily the proton, H+, is lost or removed. The more polarized the bond between H and the atom to which it is bonded, the more easily the H+ is lost or donated. We will look now at factors that affect how easily the hydrogen can be lost and, therefore, acid strength.

For a binary acid, as the size of X in HX increases, going down a group, acid strength increases.

For a binary acid, going across a period, as the electronegativity increases, acid strength increases.

Which is a stronger acid: HF or HCl?
Which is a stronger acid: H2O or H2S? Which is a stronger acid: HCl or H2S? HF and HCl These are binary acids from the same group, so we compare the size of F and Cl. Because Cl is larger, HCl is the stronger acid.

H2O and H2S These are binary acids from the same group, so we compare the size of O and S. Because S is larger, H2S is the stronger acid. HCl and H2S These are binary acids from the same period, but different groups, so we compare the electronegativity of Cl and S. Because Cl is more electronegative, HCl is the stronger acid.

For oxoacids, several factors are relevant: the number and bonding of oxygens, the central element, and the charge on the species. For a series of oxoacids, (OH)mYOn, acid strength increases as n increases. (OH)ClO n = 1 (OH)ClO2 n = 2 (OH)ClO3 n = 3 (OH)Cl n = 0 Weakest Strongest

Molecular Structure and Oxoacid Strength
Z O H O- + H+ d- d+ The O-H bond will be more polar and easier to break if: Z is very electronegative or Z is in a high oxidation state

For a series of oxoacids differing only in the central atom Y, the acid strength increases with the electronegativity of Y. Stronger Weaker

Molecular Structure and Oxoacid Strength
1. Oxoacids having different central atoms (Z) that are from the same group and that have the same oxidation number. Acid strength increases with increasing electronegativity of Z H O Cl O O • H O Br O O • Cl is more electronegative than Br acidity increases HClO3 > HBrO3

2. Oxoacids having the same central atom (Z) but different numbers of attached groups. Acid strength increases as the oxidation number of Z increases. HClO4 > HClO3 > HClO2 > HClO

The acid strength of a polyprotic acid and its anions decreases with increasing negative charge.
H2CO3 is a stronger acid than HCO3−. H2SO4 is a stronger acid than HSO4−. H3PO4 is a stronger acid than H2PO4−. H2PO4- is a stronger acid than HPO42−.

A reaction will always go in the direction from stronger acid to weaker acid, and from stronger base to weaker base.

Decide which species are favored at the completion of the following reaction:
HCN(aq) + HSO3−(aq)  CN−(aq) + H2SO3(aq) We first identify the acid on each side of the reaction: HCN and H2SO3. Next, we compare their acid strength: H2SO3 is stronger. This reaction will go from right to left (), and the reactants are favored.

Strong Acid (HCl) Weak Acid (HF)

What is the pH of a 2 x 10-3 M HNO3 solution?
HNO3 is a strong acid – 100% dissociation. Start 0.002 M 0.0 M 0.0 M HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) End 0.0 M 0.002 M 0.002 M pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7 What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution? Ba(OH)2 is a strong base – 100% dissociation. Start 0.018 M 0.0 M 0.0 M Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq) End 0.0 M 0.018 M 0.036 M pH = – pOH = log(0.036) = 12.6

Weak Acids (HA) and Acid Ionization Constants
HA (aq) + H2O (l) H3O+ (aq) + A- (aq) HA (aq) H+ (aq) + A- (aq) Ka = [H+][A-] [HA] Ka is the acid ionization constant weak acid strength Ka

What is the pH of a 0.5 M HF solution (at 250C)?
Ka = [H+][F-] [HF] = 7.1 x 10-4 HF (aq) H+ (aq) + F- (aq) HF (aq) H+ (aq) + F- (aq) Initial (M) 0.50 0.00 0.00 Change (M) -x +x +x Equilibrium (M) x x x Ka = x2 x = 7.1 x 10-4 Ka << 1 0.50 – x  0.50 Ka  x2 0.50 = 7.1 x 10-4 x2 = 3.55 x 10-4 x = M [H+] = [F-] = M pH = -log [H+] = 1.72 [HF] = 0.50 – x = 0.48 M

When can I use the approximation?
Ka << 1 0.50 – x  0.50 When x is less than 5% of the value from which it is subtracted. 0.019 M 0.50 M x 100% = 3.8% Less than 5% Approximation ok. x = 0.019 What is the pH of a 0.05 M HF solution (at 250C)? Ka  x2 0.05 = 7.1 x 10-4 x = M 0.006 M 0.05 M x 100% = 12% More than 5% Approximation not ok. Must solve for x exactly using quadratic equation or method of successive approximations.

Solving weak acid ionization problems:
Identify the major species that can affect the pH. In most cases, you can ignore the autoionization of water. Ignore [OH-] because it is determined by [H+]. Use ICE to express the equilibrium concentrations in terms of single unknown x. Write Ka in terms of equilibrium concentrations. Solve for x by the approximation method. If approximation is not valid, solve for x exactly. Calculate concentrations of all species and/or pH of the solution.

What is the pH of a 0.122 M monoprotic acid whose
Ka is 5.7 x 10-4? HA (aq) H+ (aq) + A- (aq) Initial (M) 0.122 0.00 0.00 Change (M) -x +x +x Equilibrium (M) x x x Ka = x2 x = 5.7 x 10-4 Ka << 1 0.122 – x  0.122 Ka  x2 0.122 = 5.7 x 10-4 x2 = 6.95 x 10-5 x = M M 0.122 M x 100% = 6.8% More than 5% Approximation not ok.

Ka = x2 x = 5.7 x 10-4 x x – 6.95 x 10-5 = 0 -b ± b2 – 4ac  2a x = ax2 + bx + c =0 x = x = HA (aq) H+ (aq) + A- (aq) Initial (M) Change (M) Equilibrium (M) 0.122 0.00 -x +x x x [H+] = x = M pH = -log[H+] = 2.09

For a monoprotic acid HA
Ionized acid concentration at equilibrium Initial concentration of acid x 100% percent ionization = For a monoprotic acid HA Percent ionization = [H+] [HA]0 x 100% [HA]0 = initial concentration

Weak Bases and Base Ionization Constants
NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) Kb = [NH4+][OH-] [NH3] Kb is the base ionization constant weak base strength Kb Solve weak base problems like weak acids except solve for [OH-] instead of [H+].

Ionization Constants of Conjugate Acid-Base Pairs
HA (aq) H+ (aq) + A- (aq) Ka A- (aq) + H2O (l) OH- (aq) + HA (aq) Kb H2O (l) H+ (aq) + OH- (aq) Kw KaKb = Kw Weak Acid and Its Conjugate Base Ka = Kw Kb Kb = Kw Ka

Diprotic and Triprotic Acids
May yield more than one hydrogen ion per molecule. Ionize in a stepwise manner; that is, they lose one proton at a time. An ionization constant expression can be written for each ionization stage. Consequently, two or more equilibrium constant expressions must often be used to calculate the concentrations of species in the acid solution.

Acid-Base Properties of Salts
Neutral Solutions: Salts containing an alkali metal or alkaline earth metal ion (except Be2+) and the conjugate base of a strong acid (e.g. Cl-, Br-, and NO3-). NaCl (s) Na+ (aq) + Cl- (aq) H2O Basic Solutions: Salts derived from a strong base and a weak acid. NaCH3COOH (s) Na+ (aq) + CH3COO- (aq) H2O CH3COO- (aq) + H2O (l) CH3COOH (aq) + OH- (aq)

Acid Solutions: Salts derived from a strong acid and a weak base. NH4Cl (s) NH4+ (aq) + Cl- (aq) H2O NH4+ (aq) NH3 (aq) + H+ (aq) Salts with small, highly charged metal cations (e.g. Al3+, Cr3+, and Be2+) and the conjugate base of a strong acid. Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq) 3+ 2+

Acid Hydrolysis of Al3+

Solutions in which both the cation and the anion hydrolyze: Kb for the anion > Ka for the cation, solution will be basic Kb for the anion < Ka for the cation, solution will be acidic Kb for the anion  Ka for the cation, solution will be neutral

Oxides of the Representative Elements
In Their Highest Oxidation States Na2O (s) + H2O (l) NaOH (aq) CO2 (g) + H2O (l) H2CO3 (aq) N2O5 (g) + H2O (l) HNO3 (aq)

Chemistry In Action: Antacids and the Stomach pH Balance
NaHCO3 (aq) + HCl (aq) NaCl (aq) + H2O (l) + CO2 (g) Mg(OH)2 (s) + 2HCl (aq) MgCl2 (aq) + 2H2O (l)

Acids and Bases Chapter 15

Similar presentations

Presentation on theme: "Acids and Bases Chapter 15"— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

Acids and Bases Chapter 15

Similar presentations

Presentation on theme: "Acids and Bases Chapter 15"— Presentation transcript:

Similar presentations

About project

Feedback