Mec chapter 8

Chapter 8

Acids and Bases and Oxidation-Reduction

Denniston Topping Caret

7th Edition

Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

8.1 Acids and Bases

• Acids: Taste sour, dissolve some metals, cause plant dye to change color

• Bases: Taste bitter, are slippery, are corrosive

• Two theories that help us to understand the chemistry of acids and bases1. Arrhenius Theory

2. Brønsted-Lowry Theory

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• Acid - a substance, when dissolved in water, dissociates to produce hydrogen ions– Hydrogen ion: H+ also called “protons”

HCl is an acid:

HCl(aq) H+(aq) + Cl-(aq)

Arrhenius Theory of Acids and Bases

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• Base - a substance, when dissolved in water, dissociates to produce hydroxide ions

NaOH is a base

NaOH(aq) Na+(aq) + OH-(aq)

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• Where does NH3 fit?

• When it dissolves in water it has basic properties but it does not have OH- ions in it

• The next acid-base theory gives us a broader view of acids and bases

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Brønsted-Lowry Theory of Acids and Bases

• Acid - proton donor

• Base - proton acceptor– Notice that acid and base are not defined

using water– When writing the reactions, both accepting

and donation are evident

HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)

What donated the proton? HClIs it an acid or base? Acid

What accepted the proton? H2OIs it an acid or base? Base


base

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acid

.

base acidNH3(aq) + H2O(l) NH4

+(aq) + OH-(aq)

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Now, let us look at NH3 and see why it is a

base.

Did NH3 donate or accept a proton? Accept

Is it an acid or base? Base

What is water in this reaction? Acid

Acid-Base Properties of Water

• Water possesses both acid and base properties– Amphiprotic - a substance possessing both acid

and base properties– Water is the most commonly used solvent for

both acids and bases– Solute-solvent interactions between water and

both acids and bases promote solubility and dissociation

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Acid and Base Strength

• Acid and base strength – degree of dissociation– Not a measure of concentration– Strong acids and bases – reaction with water is

virtually 100% (Strong electrolytes)

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Strong Acids and Bases

• Strong Acids:– HCl, HBr, HI Hydrochloric Acid,

etc.

– HNO3 Nitric Acid

– H2SO4 Sulfuric Acid

– HClO4 Perchloric Acid

• Strong Bases:– NaOH, KOH, Ba(OH)2

– All metal hydroxides

CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

H2CO3(aq) + H2O(l) HCO3-(aq) + H3O+(aq)

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Weak Acids

• Weak acids and bases – only a small percent dissociates (Weak electrolytes)

• Weak acid examples:– Acetic acid:

– Carbonic Acid:

• Weak base examples:– Ammonia:

– Pyridine:

– Aniline:C6H5NH2(aq) + H2O(l) C6H5NH3

+(aq) + OH-(aq)

C5H5NH2(aq) + H2O(l) C5H5NH3+(aq) + OH-(aq)

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

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Weak Bases

• The acid base reaction can be written in the general form:

• Notice the reversible arrows

• The products are also an acid and base called the conjugate acid and base

acid baseHA + B A– + HB+

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Conjugate Acids and Bases

acid base

• Conjugate acid - what the base becomes after it accepts a proton

• Conjugate base - what the acid becomes after it donates its proton

• Conjugate acid-base pair - the acid and base on the opposite sides of the equation

base acid

HA + B A- + HB+

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HA + B A– + HB+

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Acid-Base Dissociation

• The reversible arrow isn’t always written– Some acids or bases essentially dissociate 100%

– One way arrow is used

• HCl + H2O Cl- + H3O+ – All of the HCl is converted to Cl-

– HCl is called a strong acid – an acid that dissociates 100%

• Weak acid - one which does not dissociate 100%

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Conjugate Acid-Base Pairs

• Which acid is stronger:

HF or HCN? HF

• Which base is stronger:

CN- or H2O? CN -

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Acid-Base Practice• Write the chemical reaction for the following

acids or bases in water• Identify the conjugate acid-base pairs

1. HF (a weak acid)

2. H2S (a weak acid)

3. HNO3 (a strong acid)

4. CH3NH2 (a weak base)

Note: The degree of dissociation also defines weak and strong bases

• Pure water is virtually 100% molecular

• Very small number of molecules dissociate– Dissociation of acids and bases is often called

ionization

• Called autoionization

• Very weak electrolyte

H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

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The Dissociation of Water

• H3O+ is called the hydronium ion• In pure water at room temperature:

– [H3O+] = 1 x 10-7 M– [OH-] = 1 x 10-7 M

• What is the equilibrium expression for:

Remember, liquids are not included in equilibrium expressions

]OH][O[HK -3eq

H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

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Hydronium Ion

• This constant is called the ion product for water and has the symbol Kw

• Since [H3O+] = [OH-] = 1.0 x 10-7 M, what is the value for Kw?

– 1.0 x 10-14

– It is unitless

]OH][O[HK -3w

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Ion Product of Water

8.2 pH: A Measurement Scale for Acids and Bases

• pH scale - a scale that indicates the acidity or basicity of a solution– Ranges from 0 (very acidic) to 14 (very basic)

• The pH scale is rather similar to the temperature scale assigning relative values of hot and cold

• The pH of a solution is defined as:

pH = -log[H3O+]

• Use these observations to develop a concept of pH– if know one concentration, can calculate the

other

– if add an acid, [H3O+] and [OH-]

– if add a base, [OH-] and [H3O+]

– [H3O+] = [OH-] when equal amounts of acid and base are present

• In each of these cases 1 x 10-14 = [H3O+][OH-]8.2

pH: A

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• pH of a solution can be:– Calculated if the concentration of either is

known• [H3O+] • [OH-]

– Approximated using indicator / pH paper that develops a color related to the solution pH

– Measured using a pH meter whose sensor measures an electrical property of the solution that is proportional to pH8.

2 pH

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Bas

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• How do we calculate the pH of a solution when either the hydronium or hydroxide ion concentration is known?

• How do we calculate the hydronium or hydroxide ion concentration when the pH is known?

• Use two facts:

8.2

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pH = -log[H3O+]

1 x 10-14 = [H3O+][OH-]

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Calculating pH from Acid Molarity

What is the pH of a 1.0 x 10-4 M HCl solution?

– HCl is a strong acid and dissociates in water

– If 1 mol HCl is placed in 1 L of aqueous solution it produces 1 mol [H3O+]

– 1.0 x 10-4 M HCl solution has [H3O+]=1.0x10-4M

= -log [H3O+]

= -log [1.0 x 10-4]

= -[-4.00] = 4.00

pH = -log[H3O+]

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pH: A

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What is the [H3O+] of a solution with pH = 6.00?

• 4.00 = -log [H3O+]

• Multiply both sides of equation by –1

• -4.00 = log [H3O+]

• Take the antilog of both sides

• Antilog -4.00 = [H3O+]

• Antilog is the exponent of 10

• 1.0 x 10-4 M = [H3O+]

pH = -log[H3O+]

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pH: A

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What is the pH of a 1.0 x 10-3 M KOH solution?• KOH is a strong base (as are any metal hydroxides)• 1 mol KOH dissolved and dissociated in aqueous

solution produces 1 mol OH-

• 1.0 x 10-3 M KOH solution has [OH-] = 1.0 x 10-3 M

• Solve equation for [H3O+] = 1 x 10-14 / [OH-]• [H3O+] = 1 x 10-14 / 1.0 x 10-3 = 1 x 10-11

• pH = -log [1 x 10-11]

= 11.00

1 x 10-14 = [H3O+][OH-]

pH = -log[H3O+]

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Calculating pH from Acid Molarity

What is the pH of a 2.5 x 10-4 M HNO3 solution?

• We know that as a strong acid HNO3

dissociates to produce 2.5 x 10-4 M [H3O+]

• pH = -log [2.5 x 10-4]

• = 3.60

pH = -log[H3O+]

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pH: A

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What is the [OH-] of a solution with pH = 4.95?

• First find [H3O+] • 4.95 = -log [H3O+]

• [H3O+] = 10-4.95 • [H3O+] = 1.12 x 10-5

• Now solve for [OH-]• [OH-] = 1 x 10-14 / 1.12 x 10-5

= 1.0 x 10-9

pH = -log[H3O+]

1 x 10-14 = [H3O+][OH-]

The pH Scale8.

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1.0 x 100 0.001.0 x 10-1 1.001.0 x 10-2 2.001.0 x 10-3 3.001.0 x 10-4 4.001.0 x 10-5 5.001.0 x 10-6 6.001.0 x 10-7 7.00

For a strong acidHCl molarity pH

Mor

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1.0 x 100 14.001.0 x 10-1 13.001.0 x 10-2 12.001.0 x 10-3 11.001.0 x 10-4 10.001.0 x 10-5 9.001.0 x 10-6 8.001.0 x 10-7 7.00

For a strong baseNaOH molarity pH

Mor

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Each 10 fold change in concentration changes the pH by one unit

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8.3 Reactions Between Acids and Bases

• Neutralization reaction - the reaction of an acid with a base to produce a salt and water

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) Acid Base Salt Water• Break apart into ions:

H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O• Net ionic equation

– Show only the changed components– Omit any ions appearing the same on both sides of

equation = Spectator ions

H+ + OH- H2O

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• The net ionic neutralization reaction is more accurately written:

H3O+(aq) + OH-(aq) 2H2O(l)• This equation applies to any strong acid / strong

base neutralization reaction• An analytical technique to determine the

concentration of an acid or base is titration• Titration involves the addition of measured

amount of a standard solution to neutralize the second, unknown solution

• Standard solution - solution of known concentration

Net Ionic Neutralization Reaction

Buret – long glass tube calibrated in mL which contains the standard solution

Flask contains a solution of unknown concentration plus indicator

Indicator – a substance which changes color as pH changes

Standard solution is slowly added until the color changes

The equivalence point is when the moles of H3O+ and OH- are equal

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Acid – Base Titration

8.4 Acid-Base Buffers

• Buffer solution - solution which resists large changes in pH when either acids or bases are added

• These solutions are frequently prepared in laboratories to maintain optimum conditions for chemical reactions

• Buffers are also used routinely in commercial products to maintain optimum conditions for product behavior

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• Buffers act to establish an equilibrium between a conjugate acid – base pair

• Buffers consist of either– a weak acid and its salt (conjugate base)– a weak base and its salt (conjugate acid)

– Acetic acid (CH3COOH) with sodium acetate (CH3COONa)

• An equilibrium is established in solution between the acid and the salt anion• A buffer is Le Chatelier’s principle in action

CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

The Buffer Process

CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)8.4

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Addition of Base (OH-) to a Buffer Solution

• Adding a basic substance to a buffer causes changes– The OH- will react with the H3O+ producing water– Acid in the buffer system dissociates to replace

the H3O+ consumed by the added base– Net result is to maintain the pH close to the initial

level

• The loss of H3O+ (the stress) is compensated by the dissociation of the acid to produce more H3O+

CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)8.4

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Addition of Acid (H3O+) to a Buffer Solution

• Adding an acidic substance to a buffer causes changes– The H3O+ from the acid will increase the overall

H3O+ – Conjugate base in the buffer system reacts with the

H3O+ to form more acid– Net result is to maintain the H3O+ concentration and

the pH close to the initial level

• The gain of H3O+ (the stress) is compensated by the reaction of the conjugate base to produce more acid

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Buffer Capacity

• Buffer capacity - a measure of the ability of a solution to resist large changes in pH when a strong acid or strong base is added

• Also described as the amount of strong acid or strong base that a buffer can neutralize without significantly changing pH

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Mec chapter 8