Chapter 8 Acids and Bases and Oxidation- Reduction Denniston Topping Caret 6 th Edition Copyright The McGraw-Hill Companies, Inc. Permission required

Chapter 8

Acids and Bases and Oxidation-Reduction

Denniston Topping Caret

6th 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

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

HCl is an acid:

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

Arrhenius Theory of Acids and Bases


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

NaOH is a base

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


• 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

Brønsted-Lowry Theory of Acids and Bases

• Acid - proton donor

• Base - proton acceptoroNotice that acid and base are not defined

using wateroWhen 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


baseacid

.

base acidNH3(aq) + H2O(l) NH4

+(aq) + OH-(aq)


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 propertiesoAmphiprotic – a substance possessing both acid

and base propertiesoWater is the most commonly used solvent for

both acids and baseso Solute-solvent interactions between water and

both acids and bases promote solubility and dissociation

Acid and Base Strength

• Acid and base strength – degree of dissociationoNot a measure of concentrationo Strong acids and bases – reaction with water is

virtually 100% (Strong electrolytes)

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)

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)

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+

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+

HA + B A- + HB+

Acid-Base Dissociation

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

oOne way arrow is used

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

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

• Weak acid - one which does not dissociate 100%

Conjugate Acid-Base Pairs

• Which acid is stronger:

HF or HCN? HF

• Which base is stronger:

CN- or H2O? CN -

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 dissociateoDissociation of acids and bases is often called

ionization

• Called autoionization

• Very weak electrolyte

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

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)

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?

o 1.0 x 10-14

o It is unitless

]OH][O[HK -3w

+=

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 solutiono 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 pHo if know one concentration, can calculate the

other

o if add an acid, [H3O+] and [OH-]

o if add a base, [OH-] and [H3O+]

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

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

A Definition of pH

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

known• [H3O+] • [OH-]

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

oMeasured using a pH meter whose sensor measures an electrical property of the solution that is proportional to pH

Measuring pH

• 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:

Calculating pH

pH = -log[H3O+]

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

Calculating pH from Acid Molarity

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

oHCl is a strong acid and dissociates in watero If 1 mol HCl is placed in 1 L of aqueous

solution it produces 1 mol [H3O+]o 1.0 x 10-4 M HCl solution has [H3O+]=1.0x10-4M

= -log [H3O+]

= -log [1.0x10-4]

= -[-4.00] = 4.00

pH = -log[H3O+]

Calculating [H3O+] from pH

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+]

Calculating the pH of a Base

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+]

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]

• = 3.6

pH = -log[H3O+]

Calculating [OH-] from pH

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

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

• [H3O+] = 10-4.95 • [H3O+] = 1.122 x 10-5

• Now solve for [OH-]• [OH-] = 1 x 10-14 / 1.122 x 10-5

= 8.91 x 10-10

pH = -log[H3O+]

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

The pH Scale

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

o Show only the changed componentso Omit any ions appearing the same on both sides of

equation = Spectator Ions

H+ + OH- H2O

• 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

Acid – Base Titration

Determine the Concentration of a Solution of Hydrochloric Acid• Place a known volume of acid whose concentration

is not known into a flask• Add an indicator, experience guides selection, here

phenol red is good• Known concentration of NaOH is placed in a buret• Drip NaOH into the flask until the indicator

changes color

Determine the Concentration of a Solution of Hydrochloric Acid

• Indicator changes color o equivalence point is reached

o mole OH- = mole H3O+ present in the unknown acid

• Volume dispensed from buret is determined• Calculate acid concentration:

o Volume of Hydrochloric Acid: 25.00 mLo Volume of NaOH added: 35.00 mLo Concentration of NaOH: 0.1000 Mo Balanced reaction shows that HCl and NaOH react 1:1

Determine the Concentration of a Solution of Hydrochloric Acid

• 35.00 mL NaOH x 1L NaOH x 0.1000 mol NaOH 103 mL NaOH 1L NaOH

= 3.500 x 10-3 mol NaOH• 3.500 x 10-3 mol NaOH x 1 mol HCl 1 mol NaOH

= 3.500 x 10-3 mol HCl– this amount of HCl is contained in 25.00 mL

• 3.500 x 10-3 mol HCl x 103 mL HCl 25.00 mL HCl 1 L HCl• = 1.400 x 10-1 mol HCl / L HCl = 0.1400 M HCl

• The previous examples have the acid and base at a 1:1 combining ratioo Not all acid-base pairs do this

• Polyprotic substance – donates or accepts more than one proton per formula unito Hydrochloric acid is monoprotic, producing one H+ ion for each unit of HCl

o Sulfuric acid is diprotic, each unit of H2SO4 produces 2 H+ ions

H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2 H2O(l)

Polyprotic Substances

Step 1.

H2SO4(aq) + H2O(l) HSO4-(aq) + H3O+(aq)

Step 2.

HSO4-(aq) + H2O(l) SO4

2-(aq) + H3O+(aq)

• In Step 1 H2SO4 behaves as a strong acid – dissociating completely

• In Step 2 HSO4-( behaves as a weak acid – reversibly dissociating,

note the double arrow

Dissociation of Polyprotic Substances

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

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

• Buffers consist of eithero a weak acid and its salt (conjugate base)o a weak base and its salt (conjugate acid)

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

• An equilibrium is established in solution between the acid and the salt anion• A buffer is LeChatelier’s Principle in action


The Buffer Process


Addition of Base (OH-) to a Buffer Solution

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

the H3O+ consumed by the added baseo 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+


Addition of Acid (H3O+) to a Buffer Solution

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

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

H3O+ to form more acido 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

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

• Buffering process is an equilibrium reaction described by an equilibrium-constant expression

In acids, this constant is Ka

• If you want to know the pH of the buffer, solve for [H3O+], then calculate pH

COOH]CH[

]COOCH][OH[K

3

-33

a

+

=

Preparation of a Buffer Solution


Henderson-Hasselbach Equation

• Solution of equilibrium-constant expression and pH can be combined into one operation

• Henderson-Hasselbach Equation is this combined expression

• Using these two equations:

o pKa = -log Ka just as pH = -log[H3O+]

o pKa = pH – log ( [CH3COO-] / [CH3COOH] )

o Henderson-Hasselbach –

o pH = pKa + log( [CH3COO-] / [CH3COOH] )


COOH]CH[

]COOCH][OH[K

3

-33

a

+

=

Henderson-Hasselbach Equation

• pH = pKa + log( [CH3COO-] / [CH3COOH] ) can be rewritten

pH = pKa + log ( [conjugate base] / [weak acid])

8.5 Oxidation-Reduction Processes

• Oxidation-reduction processes are responsible for many types of chemical change

• Oxidation - defined by one of the following o loss of electrons o loss of hydrogen atomso gain of oxygen atoms

• Example: NaNa+ + e-

o Oxidation half reaction

• Reduction - defined by one of the following:o gain of electronso gain of hydrogeno loss of oxygen

• Example: Cl + e- Cl-

o Reduction half reaction

• Cannot have oxidation without reduction.

Oxidation-Reduction Processes

Na + Cl Na+ + Cl-

Oxidizing Agent• Is reduced• Gains electrons• Causes oxidation

Reducing Agent• Is oxidized• Loses electrons• Causes reduction

Oxidation and Reduction as Complementary Processes

Na Na+ + e-

Cl + e- Cl-

Applications of Oxidation and Reduction

• Corrosion - the deterioration of metals caused by an oxidation-reduction process

Example: rust (oxidation of iron)

4Fe(s) + 3O2(g) 2Fe2O3(s)

• Combustion of Fossil Fuels

Example: natural gas furnaces

CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

• Bleaching

• Most bleaching agents are oxidizing agents

• The oxidation of the stains produces compounds that do not have color

Example: Chlorine bleach - sodium hypochlorite (NaOCl)

Applications of Oxidation and Reduction

Biological Processes

• Respirationo Electron-transport chain of aerobic

respiration uses reversible oxidation and reduction of iron atoms in cytochrome c

• Metabolismo Break down of molecules into smaller pieces

by enyzmes

• Is Zn oxidized or reduced?• Oxidized

• Copper is reduced

Voltaic Cells

• Voltaic cell – electrochemical cell that converts stored chemical energy into electrical energy

• Let’s consider the following reaction:

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Voltaic Cells

• If the two reactants are placed in the same flask they cannot produce electrical current

• A voltaic cell separates the two half reactions

• This makes the electrons flow through a wire to allow the oxidation and reduction to occur

Zn Zn2+ + 2e-

Oxidationanode – electrode

where oxidation occurs

Cu2+ + 2e- CuReduction

cathode – electrode where reduction occurs

Voltaic Cell Generating Electrical Current

Voltaic Cell Generating Electrical Current

Silver Battery• Batteries use the concept of the voltaic cell

Electrolysis

• Electrolysis reactions – uses electrical energy to cause nonspontaneous oxidation-reduction reactions to occur

• These reactions are the reverse of a voltaic cell

Rechargeable battery• When powering a device behaves as voltaic cell• With time the chemical reaction nears

completion• Battery appears to “run down”• Cell reaction is reversible when battery attached

to charger

Comparison of Voltaic and Electrolytic Cells

Documents

Chapter 8 Acids and Bases and Oxidation- Reduction Denniston Topping Caret 6 th Edition Copyright The McGraw-Hill Companies, Inc. Permission required