Instructor: Dr. Upali Siriwardanee-mail: upali@chem.latech.eduOffice: CTH 311 Phone 257-4941

Office Hours: M,W, 8:00-9:00 & 11:00-12:00 a.m.; Tu,Th,F 9:00 - 10:00 a.m.  

Test Dates: March 25, April 26, and May 18; Comprehensive Final Exam: May 20,2009 9:30-10:45 am, CTH 328.

March 30, 2009 (Test 1): Chapter 13 April 27, 2009 (Test 2): Chapters 14 & 15 May 18, 2009 (Test 3): Chapters 16, 17 & 18

Comprehensive Final Exam: May 20,2009 :Chapters 13, 14, 15, 16, 17 and 18

Chemistry 102(01) spring 2009

Chapter 16. Acids and Bases

16.1 The Brønsted-Lowry Concept of Acids and Bases

16.2 Types of acids/bases:Organic Acids and Amines

16.3 The Autoionization of Water

16.4 The pH Scale

16.5 Ionization Constants of Acids and Bases

16.6 Problem Solving Using Ka and Kb

16.7 Molecular Structure and Acid Strength

16.8 Acid-Base Reactions of Salts

16.9 Practical Acid-Base Chemistry

16.10 Lewis Acid and Bases

Types of Reactionsa) Precipitation Reactions. Reactions of ionic compounds or saltsb) Acid/base Reactions. Reactions of acids and basesc) Redox Reactions. reactions of oxidizing & reducing

agents

What are Acids &Bases?

Definition?

a) Arrhenius

b) Bronsted-Lowry

c) Lewis

Arrhenius, Svante August (1859-1927), Swedish chemist, 1903 Nobel Prize in chemistry

• AcidAcidAnything that produces hydrogen ions in a water solution.

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

• BaseBaseAnything that producs hydroxide ions in a water solution.

NaOH (aq) Na+ ( aq) + OH- ( aq) • Arrhenius definitions are limited proton acids

and hydroxide bases to aqueous solutions.

Arrhenius Definitions

Expands the Arrhenius definitions to include many bases other than hydroxides and gas phase reactions

AcidAcid Proton donor

BaseBase Proton acceptor

This definition explains how substances like ammonia can act as bases.

Eg. HCl(g) + NH3(g) ------> NH4Cl(s)

HCl (acid), NH3 (base).

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

Brønsted-Lowry definitions

Lewis DefinitionG.N. Lewis was successful in including acid and bases

without proton or hydroxyl ions.

Lewis Acid: A substance that accepts an electron pair.

Lewis base: A substance that donates an electron pair.

E.g. BF3(g) + :NH3(g) F3B:NH3(s)

the base donates a pair of electrons to the acid forming a

coordinate covalent bond common to coordination

compounds. Lewis acids/bases will be discussed later in

detail

Dissociation

Strong Acids:

HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)

H2SO4(aq) + H2O(l) H3+O(aq) + HSO4

-(aq)

Dissociation Equilibrium Weak Acid/base:

H2O(l) + H2O(l) H3+O(aq) + OH-(aq)

This dissociation is called autoionization of water.

HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2

-(aq)

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

Equilibrium constants: Ka, Kb and Kw

Conjugate acid-base pairs.Conjugate acid-base pairs.

Acids and bases that are related by loss or gain of H+ as H3O+ and H2O.

Examples.Examples. Acid Base

H3O+ H2O

HC2H3O2 C2H3O2-

NH4+ NH3

H2SO4 HSO4-

HSO4- SO42-

Brønsted-Lowry Definitions

Bronsted acid/conjugate base and base/conjugate acid pairs in

acid/base equilibria

HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)

HCl(aq): acid

H2O(l): base

H3+O(aq): conjugate acid

Cl-(aq): conjugate base

H2O/ H3+O: base/conjugate acid pair

HCl/Cl-: acid/conjugate base pair

Select acid, base, acid/conjugate base pair,base/conjugate acid pair

H2SO4(aq) + H2O(l) H 3+O(aq) + HSO4

-(aq)

acid

base

conjugate acid

conjugate base

base/conjugate acid pair

acid/conjugate base pair

Types of Acids and BasesBinary acids: HCl, HBr, HI, H2S

More than two elements: HCN

Oxyacid: HNO3, H2SO4, H3PO4

Polyprotic acids: H2SO4, H3PO4

Organic acids: R-COOH, R= CH3-, CH3CH2-

Acidic oxides: SO3, NO2, CO2,

Basic oxides: Na2O, CaO

Amine: NH3. R-NH2, R= CH3-, CH3CH2- : primary

R2-NH : secondary, R3-N: tertiary

Lewis acids & bases: BF3 and NH3

Strong Acid vs. Weak AcidsStrong acidcompletely ionized

Hydrioidic HI Ka ~ 1011 pKa = -11Hydrobromic HBr Ka ~ 109 pKa = -9Perchloric HClO4 Ka ~ 107 pKa = -7Hyrdrochloric HCl Ka ~ 107 pKa = -7Chloric HClO3 Ka ~ 103 pKa = -3Sulfuric H2SO4 Ka ~ 102 pKa = -2Nitric HNO3 Ka ~ 20 pKa = -1.3

Weak acidpartially ionized

Hydrofluoric acid HF Ka = 6.6x10-4 pKa = 3.18Formic acid HCOOH Ka = 1.77x10-4 pKa = 3.75Acetic acid CH3COOH Ka = 1.76x10-5 pKa = 4.75Nitrous acid HNO2 Ka = 4.6x10-4 pKa = 3.34Acetyl Salicylic acid C9H8O4 Ka = 3x10-4 pKa = 3.52Hydrocyanic acid HCN Ka = 6.17x10-10 pKa = 9.21

Strong Base vs. Weak BaseStrong Basecompletely ionizedLithium hydroxide LiOHSodium hydroxide NaOH

Potasium hydroxide KOH Kb~ 102-103

Rubidium hydroxide RbOHCesium hydroxide CsOHBoarder-line Bases

Magnesium hydroxide Mg(OH)2

Calcium hydroxide Ca(OH)2

Strotium hydroxide Sr(OH)2 Kb~ 0.01 to0.1

Barium hydroxide Ba(OH)2

Weak Base

partially ionized

Ammonia NH3 Kb=1.79x10-5 pKb = 4.74

Ethyl amine CH3CH2NH2 Kb=5.6x10-4 pKb = 3.25

• Strong acidsStrong acids Ionize completely in water. HCl, HBr, HI, HClO3,

HNO3, HClO4, H2SO4.

• Weak acidsWeak acids Partially ionize in water. Most acids are weak.

• Strong basesStrong bases Ionize completely in water. Strong bases are metal

hydroxides - NaOH, KOH

• Weak basesWeak bases Partially ionize in water.

Acid and Base StrengthAcid and Base Strength

Common Acids and BasesAcidsAcids Formula Molarity*

nitric HNO3 16

hydrochloric HCl 12

sulfuric H2SO4 18

acetic HC2H3O2 18

BasesBases

ammonia NH3(aq) 15

sodium hydroxide NaOH solid

*undiluted.

AutoionizationAutoionization When water molecules react with one another to form ions.

Acids and bases alter the dissociation equilibrium of water based on Le Chaterlier’s principle

Kw = [ H3O+ ] [ OH- ]

= 1.0 x 10-14 at 25oC

Note:Note: [H2O] is constant and is included in Kw.

ion productof water

ion productof water

H2O(l) + H2O(l) H3O+(aq) + OH-(aq) (10-7M) (10-7M)

Autoionization of Water

We need to measure and use acids and bases over a very large concentration range.

pH and pOH are systems to keep track of these very large ranges.pH = -log[H3O

+]

pOH = -log[OH-]

pH + pOH = 14

Substance pH

1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0

1M NaOH 14.0

Substance pH

1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0

1M NaOH 14.0

pH and other “p” scales

A logarithmic scale used to keep track of the large changes in [H+].

0 7 1410-14 M 10-7 M 10-14 M

Very Neutral Veryacidic BasicWhen you add an acid to, the pH gets smaller.

When you add a base to, the pH gets larger.

pH scale

Substance pH

1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0

1M NaOH 14.0

Substance pH

1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0

1M NaOH 14.0

pH of some common materials

pH of Aqueous Solutions

What is pH?

Kw = [H3+O][OH-] = 1 x 10-14

[H3+O][OH-] = 10-7 x 10-7

Extreme cases:

Basic medium

[H3+O][OH-] = 10-14 x 100

Acidic medium

[H3+O][OH-] = 100 x 10-14

pH value is -log[H+]

spans only 0-14 in water.

pH, pKw and pOHThe relation of pH, Kw and pOH Kw = [H+][OH-] log Kw = log [H+] + log [OH-] -log Kw= -log [H+] -log [OH-] ; previous equation multiplied by -1 pKw = pH + pOH; pKw = 14 since Kw =1 x 10-14

14 = pH + pOH pH = 14 - pOH pOH = 14 - pH

14 = pH + pOH pH = 14 - pOH pOH = 14 - pH

pH and pOH calculations of acid and base solutions

a) Strong acids/bases

dissociation is complete for strong acid such as HNO3 or base NaOH

[H+] is calculated from molarity (M) of the

solution

b) weak acids/bases

needs Ka , Kb or percent(%)dissociation

pH of Strong Acid/bases

HNO3(aq) + H2O(l) H3+O(aq) + NO3-(aq)

Therefore, the moles of H+ ions in the solution is equal to moles of HNO3 at the beginning.

[HNO3] = [H+] = 0.2 mole/L

pH = -log [H+]

= -log(0.2)

pH = 0.699

Substance pH

1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0

1M NaOH 14.0

Substance pH

1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0

1M NaOH 14.0

pH of 0.5 M H2SO4 Solution

H2SO4(aq) + H2O(l) H3+O(aq) + HSO4

-(aq)

HSO4-(aq) + H2O(l) H3

+O(aq) + SO42-(aq)

[H3+O][HSO4

-]

H2SO4 ; Ka1 = -------------------

[H2SO4]

[H3+O][SO4

2-]

H2SO4 ; Ka2 = ------------------- ; Ka2 ignored

[HSO4-]

H2SO4(aq) + H2O(l) H3+O(aq) + HSO4

-(aq)

the moles of H+ ions in the solution is equal to moles of H2SO4 at the beginning.

[H2SO4] = [H+] = 0.5 mole/L

pH = -log [H+]

pH = -log(0.5) pH = 0.30

pH of 0.5 M H2SO4 Solution

1.5 x 10-2 M NaOH.1.5 x 10-2 M NaOH.

NaOH is also a strong base dissociates completely in water.

[NaOH] = [HO- ] = 1.5 x 10-2 mole/L

pOH = -log[HO-]= -log(1.5 x 10-2)

pOH = 1.82

As defined and derived previously: pKw= pH + pOH; pKw= 14

pH = pKw + pOH

pH = 14 - pOH

pH = 14 - 1.82 ; pH = 12.18

Mixtures of Strong and Weak Acids

• the presence of the strong acid retards the dissociation of the weak acid

Measuring pH

Arnold Beckman

• inventor of the pH meter

• father of electronic instrumentation

Equilibrium, Constant, Ka & Kb

Ka: Acid dissociation constant for a equilibrium reaction.

Kb: Base dissociation constant for a equilibrium reaction.

Acid: HA + H2O H3+O + A-

Base: BOH + H2O B+ + OH-

[H3+O][ A-] [B+ ][OH-]

Ka = --------------- ; Kb = -----------------

[HA] [BOH]

HCl(aq) + H2O(l) H3

+O(aq) + Cl-(aq)

[H3+O][Cl-]

Ka= ----------------- [HCl]

[H+][Cl-] Ka= ----------------- [HCl]

Acid Dissociation Constant

Base Dissociation Constant

NH3 + H2O NH4+ + OH-

[NH4+][OH-]

K = [NH3]

Hydrated Metal Ions as Acids

[Fe(H2O)6]3+ (aq) + H2O () [Fe(H2O)5(OH)]2+ (aq) + H3O+ (aq)

Ka [Fe(H2O)5 (OH)2 ][H3O ]

Fe(H2O)63 6.310 3

Ionization Constants for Acids

Comparing Kw and Ka & Kb

• Any compound with a Ka value greater than Kw of water will be a an acid in water.

• Any compound with a Kb value greater than Kw of water will be a base in water.

WEAKER/STRONGER Acids and Bases & Ka and Kb values

• A larger value of Ka or Kb indicates an equilibrium favoring product side.

• Acidity and basicity increase with increasing Ka or Kb.

• pKa = - log Ka and pKb = - log Kb

• Acidity and basicity decrease with increasing pKa or pKb.

Which is weaker?

• a. HNO2    ;  Ka= 4.0 x 10-4.

• b. HOCl2    ;   Ka= 1.2 x 10-2.

• c. HOCl     ;  Ka= 3.5 x 10-8.

• d. HCN      ;  Ka= 4.9 x 10-10.

What is Ka1 and Ka2?

• H2SO4(aq) + H2O(l) H3+O(aq) + HSO4

-(aq)

• HSO4-(aq) + H2O(l) H3

+O(aq) + SO42-(aq)

H2SO4(aq) + H2O(l) H3+O(aq) + HSO4

-(aq)

HSO4-(aq) + H2O(l) H3

+O(aq) + SO42-(aq)

[H3+O][HSO4

-]

H2SO4 ; Ka1 = -------------------

[H2SO4]

[H3+O][SO4

2-]

H2SO4 ; Ka2 = -------------------

[HSO4-]

Ka Examples

HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2

-(aq)

[H+][C2H3O2-]

H C2H3O2; Ka= ------------------

[H C2H3O2]

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

[NH4+][OH-]

NH3; Kb= --------------

[ NH3]

Ka Examples

How do you calculate pH of How do you calculate pH of weak acids/basesweak acids/bases

From % dissociation

From Ka or Kb

What is % dissociation

Amount dissociated

% Dissoc. = ------------------------- x 100

Initial amount

How do you calculate % How do you calculate % dissociation from Kdissociation from Kaa or K or Kbb

1.00 M solution of HCN; Ka = 4.9 x 10-10

What is the % dissociation for the acid?

1.00 M solution of HCN; Ka = 4.9 x 10-10

First write the dissociation equilibrium equation:HCN(aq) + H 2O(l) <===> H 3

+O(aq) + CN-(aq)

[HCN] [H+ ] [CN- ]

Ini. Con. 1.00 M 0.0 M 0.00 M

Cha. Con -x x x

Eq. Con. 1.0 - x x x

[H 3+O ][CN-] x2

Ka = ----------- = ----------------

[HCN] 1.0 - x

1.00 M solution of HCN; Ka = 4.9 x 10-10

1.0 - x ~ 1.00 since x is small

x2

Ka = -----------; Ka = 4.9 x 10-10 = x2

1.0

x = 4.9 x 10-10 = 2.21 x 10 -5

Amount disso. 2.21 x 10 -5

----------------- x 100 =- ------------- x 100 Ini. amount 1.00

% Diss. =2.21 x 10 -5 x 100 = 0.00221 %

1.00 M solution of HCN; Ka = 4.9 x 10-10

% Dissociation gives x (amount dissociated) need for pH calculation

Amount dissociated

% Dissoc. = ------------------------- x 100

Initial amount/con.

x

% Dissoc. = --------------------------- x 100

concentration

1 M HF, 2.7% dissociated

Notice the conversion of % dissociation to a fraction (x): 2.7/100=0.027) x=0.027

Calculate the pH of a weak acid from % dissociation

• HF(aq) + H 2O(l) <===> H 3+O(aq) + F-(aq)

• [H+][F-] • Ka = -----------• [HF]• [HF] [H+ ] [F- ]• Ini. Con. 1.00 M 0.0 M 0.00 M• Chg. Con -x x x• Eq.Con. 1.0-0.027 0.0270.027• pH = -log [H+] • pH = -log(0.027) • pH = 1.57

Calculate the pH of a weak acid from % dissociation

Weak acid EquilibriaExampleExample

Determine the pH of a 0.10 M benzoic acid solution at 25 oC if Ka = 6.28 x 10-5

HBz(aq) + H2O(l) H3O+(aq) + Bz-(aq)

The first step is to write the equilibrium expression

Ka = [H3O+][Bz-]

[HBz]

Weak acid Equilibria HBz H3O+ Bz-

Initial conc., M 0.10 0.00 0.00

Change, M -x x x

Eq. Conc., M 0.10 - x x x

[H3O+] = [Bz-] = x

We’ll assume that [Bz-] is negligible compared to [HBz]. The contribution of H3O+ from water is also negligible.

Weak Acid Equilibria

Solve the equilibrium equation in terms of Solve the equilibrium equation in terms of xx

Ka = 6.28 x 10-5 =

x = (6.28 x 10-5 )(0.10)

H3O+ = 0.0025 M

pH = 2.60

x2

0.10

pH from Ka or Kb

1.00 M solution of HCN; Ka = 4.9 x 10-10

First write the dissociation equilibrium equation:

HCN(aq) + H 2O(l) H 3+O(aq) + CN-(aq)

[HCN] [H+ ] [CN- ]

Ini. Con. 1.00 M 0.0 M 0.00 M

Chg. Con -x x x

Eq. Con. 1.0 - x x x

[H 3+O ][CN-] x2

Ka = --------------- = ----------------

[HCN] 1.0 - x

1.0 - x ~ 1.00 since x is small x2 Ka = -----------; Ka = 4.9 x 10-10 = x2

1.0

x = 4.9 x 10-10 = 2.21 x 10 -5

pH = -log [H+]

pH = -log(2.21 x 10-5)

pH = 4.65

Weak Acid Equilibria

The Conjugate Partners of Strong Acids and Bases

The conjugate acid/base of a strong base/acid has no net effect on the pH of a solution

The conjugate base of a weak acid hydrolyze in water and basic or

pH of a solution > 7.00 E.g. Na+C2H3O2- sodium acetate

The conjugate acid of a weak base hydrolyze in water and acidic or

pH of a solution < 7.00 E.g NH4Cl

Reaction of a basic anion or acidic cation with water is an ordinary Brønsted-Lowry acid-base reaction.

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

NH4+(aq) + H2O(l) NH3 (aq) + H3O+(aq)

This type of reaction is given a special name.

HydrolysisHydrolysis

The reaction of an anion with water to produce the conjugate acid and OH-.

The reaction of a cation with water to produce the conjugate base and H3O+.

Hydrolysis

Acid-Base Properties of Typical Ions

What salt solutions would be acidic, basic and

neutral?1) strong acid + strong base = neutral 2) weak acid + strong base = basic 3) strong acid + weak base = acidic 4) weak acid + weak base = neutral, basic or an acidic solution

depending on the relative strengths of the acid and the base.

What pH? Neutral, basic or acidic?

•a)NaCl • neutral•b) NaC2H3O2

• basic•c) NaHSO4 • acidic•d) NH4Cl• acidic

How do you calculate pH of a salt solution?

• Find out the pH, acidic or basic?

• If acidic it should be a salt of weak base

• If basic it should be a salt of weak acid

• if acidic calculate Ka from Ka= Kw/Kb

• if basic calculate Kb from Kb= Kw/Ka

• Do a calculation similar to pH of a weak acid or base

What is the pH of 0.5 M NH4Cl salt solution? (NH 3; Kb = 1.8 x 10-5)

• Find out the pH, acidic

• if acidic calculate Ka from Ka= Kw/Kb

• Ka= Kw/Kb = 1 x 10-14 /1.8 x 10-5)

• Ka= 5.56. X 10-10

• Do a calculation similar to pH of a weak acid

Continued

NH4+ + H2O H 3

+O + NH3

[NH4+] [H3

+O ] [NH3 ]Ini. Con. 0.5 M 0.0 M 0.00 MChange -x x xEq. Con. 0.5 - x x x

[H 3+O ] [NH3 ]

Ka(NH4+) = -------------------- =

[NH 4+] x2

---------------- ; appro.:0.5 - x . 0.5 (0.5 - x)

x2 Ka(NH4

+) = ----------- = 5.56 x 10 -10

0. 5 x2

= 5.56 x 10 -10 x 0.5 = 2.78 x 10 -10

x= 2.78 x 10 -10 = 1.66 x 10-5

[H+ ] = x = 1.66 x 10-5 MpH = -log [H+ ] = - log 1.66 x 10-5

pH = 4.77pH of 0.5 M NH4Cl solution is 4.77 (acidic)

Continued

Types of Acids and Bases

• Binary acids

• Oxyacid

• Organic acids

• Acidic oxides

• Basic oxides

• Amine

• Polyprotic acids

Influence of Molecular Structure on Acid Strength

Binary Hydrides– hydrogen & one other element

• Bond Strengths– weaker the bond, the stronger the acid

• Stability of Anion– higher the electronegativity, stronger the acid

Binary Acids

Compounds containing acidic protons bonded to a more electronegative atom.

e.g. HF, HCl, HBr, HI, H2S

The acidity of the haloacid (HX; X = Cl, Br, I, F)Series increase in the following order: HF < HCl < HBr < HI

Oxyacids

Compounds containing acidic - OH groups in the molecule.

Acidity of H2SO4 is greater than H2SO3 because of the extra O (oxygens)

The order of acidity of oxyacids from the a halogen (Cl, Br, or I) shows a similar trend.

HClO4 > HClO3 > HClO2 > HClO

perchloric chloric chlorus hyphochlorus

Influence of Molecular Structure on Acid Strength

Oxyacids– hydrogen, oxygen, & one other element

H-O-E– higher the electronegativity on E, stronger the

acid as this weakens the bond between the O and H

< <<

<

Oxo Acid

Acidic Oxides

These are usually oxides of non-metallic elements such as P, S and N.

E.g. NO2, SO2, SO3, CO2

They produce oxyacids when dissolved in water

SO3 + H2O ---> H2SO4

CO2 + H2O ---> H2CO3

NO2 + H2O ---> HNO3

Basic Oxides

Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water.

e.g. CaO + H2O ---> Ca(OH)2

Na2O + H2O ---> 2 NaOH

Protic Acids

Monoprotic Acids: The form protic refers to acidity due to protons. Monoprotic acids have only one acidic proton. e.g. HCl.

Polyprotic Acids: They have more than one acidic proton.

e.g. H2SO4 - diprotic acid

H3PO4 - triprotic acid.

Polyprotic Acids

• acids where more than one hydrogen per molecule is released

Polyprotic Acids

Organic or Carboxylic Acids

H C

H

H

C

H

H

C

H

H

C

O

O H

nonacidic hydrogens

butanoic acid

acidic hydrogen

CH 3 C

O

acetic acid

OH

electron-attractingoxygen atom

acidic hydrogen

CH 3 C

O

OH

-

CH 3 C

O

O-

acetate ion

FCH2CO2H (strongest acid) > ClCH2CO2H > BrCH2CO2H (weakest acid).

Acid Ka pKa

HCOOH (formic acid) 1.78 X 10-43 0.75

CH3COOH (acetic acid) 1.74 X 10-54 0.76

CH3CH2COOH (propanoic acid)1.38 x 10-5 4.86

Organic or Carboxylic Acids

Amines

Class of organic bases derived from ammonia NH3 by replacing hydrogen by organic groups. They are defined as bases similar to NH3 by Bronsted-Lowery or Lewis acid/base definitions.

Amines

Acid-Base Chemistryof Some Antacids

Acid-Base in the Kitchen

vinegar - acetic acid

lemon juice (citrus juice) - citric acid

baking soda - NaHCO3

milk - lactic acid

baking powder - H2PO4- & HCO3

-

Household Cleaners

CH 3CH2CH 2CH2CH 2CH2CH 2CH 2CH 2CH 2CH2CH 2CH2CH 2SO3

-Na+

Oil-soluble part(hydrophobic)

Water-soluble part(hydrophilic)

A Typical Synthetic Detergent Molecule

CH 3(CH 2)4COO(CH 2)2O( CH2CH 2O) 2CH 2CH 2OH

esterlink

(hydro-philic)

etherlink

etherlink

(hydrophilic)

hydrocarbonchain

(hydro-phobic)

alcohol group(hydrophilic)

A nonionic detergent

Dishwashing Detergent

Lewis DefinitionG.N. Lewis was successful in including acid and bases

without proton or hydroxyl ions.

Lewis Acid: A substance that accepts an electron pair.

Lewis base: A substance that donates an electron pair.

E.g. BF3(g) + :NH3(g) F3B:NH3(s)

the base donates a pair of electrons to the acid forming a

coordinate covalent bond common to coordination

compounds. Lewis acids/bases will be discussed later in

detail

Lewis Acids and Bases Reactions

H+ + NH3

acid base

Cu+2 + 4 NH3 [Cu(NH3)4+2]

acid base

What acid base concepts (Arrhenius/Bronsted/Lewis) would best

describe the following reactions:

•a) HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)

•b)HCl(g) + NH3(g) ---> NH4Cl(s)

•c)BF3(g) + NH3(g) ---> F3B:NH3(s)

•d)Zn(OH)2(s) + 2OH-(aq) ---> [Zn(OH)4]2- (aq)