Transcript CHEMISTRY The Molecular Science

Chemistry 102(01) spring 2009
Instructor: Dr. Upali Siriwardane
e-mail: [email protected]
Office: 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
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 Reactions
a) Precipitation Reactions.
Reactions of ionic compounds or salts
b) Acid/base Reactions.
Reactions of acids and bases
c) Redox Reactions.
reactions of oxidizing & reducing
agents
What are Acids &Bases?
Definition?
a) Arrhenius
b) Bronsted-Lowry
c) Lewis
Arrhenius Definitions
Arrhenius, Svante August (1859-1927), Swedish
chemist, 1903 Nobel Prize in chemistry
• Acid Anything that produces hydrogen
ions in a water solution.
HCl (aq)
H+ ( aq) + Cl- ( aq)
• Base Anything 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.
Brønsted-Lowry definitions
Expands the Arrhenius definitions to include many
bases other than hydroxides and gas phase reactions
Acid
Proton donor
Base
Proton acceptor
This definition explains how substances like ammonia
can act as bases.
NH3(g) + H2O(l)
NH4+ + OH-
Eg. HCl(g) + NH3(g) ------> NH4Cl(s)
HCl (acid), NH3 (base).
Lewis Definition
G.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
Brønsted-Lowry Definitions
Conjugate acid-base pairs.
Acids and bases that are related by loss or
gain of H+ as H3O+ and H2O.
Examples.
Acid
Base
H3O+
H2O
HC2H3O2 C2H3O2NH4+
NH3
H2SO4
HSO4-
HSO4-
SO42-
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)
acid
base
conjugate acid
conjugate base
base/conjugate acid pair
acid/conjugate base pair
H 3+O(aq) + HSO4-(aq)
Types of Acids and Bases
Binary acids: HCl, HBr, HI, H2S
More than two elements: HCN
Oxyacid: HNO3, H2SO4, H3PO4
Polyprotic acids: H2SO4, H3PO4
Organic acids: R-COOH, R= CH3-, CH3CH2Acidic 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 Acids
Strong acid
completely ionized
Hydrioidic
Hydrobromic
Perchloric
Hyrdrochloric
Chloric
Sulfuric
Nitric
HI
HBr
HClO4
HCl
HClO3
H2SO4
HNO3
Ka ~ 1011
Ka ~ 109
Ka ~ 107
Ka ~ 107
Ka ~ 103
Ka ~ 102
Ka ~ 20
pKa = -11
pKa = -9
pKa = -7
pKa = -7
pKa = -3
pKa = -2
pKa = -1.3
Weak acid
partially ionized
Hydrofluoric acid HF
Formic acid HCOOH
Acetic acid CH3COOH
Nitrous acid HNO2
Acetyl Salicylic acid C9H8O4
Hydrocyanic acid HCN
Ka = 6.6x10-4
Ka = 1.77x10-4
Ka = 1.76x10-5
Ka = 4.6x10-4
Ka = 3x10-4
Ka = 6.17x10-10
pKa = 3.18
pKa = 3.75
pKa = 4.75
pKa = 3.34
pKa = 3.52
pKa = 9.21
Strong Base vs. Weak Base
Strong Base
completely ionized
Lithium hydroxide
Sodium hydroxide
Potasium hydroxide
Rubidium hydroxide
Cesium hydroxide
LiOH
NaOH
KOH
RbOH
CsOH
Boarder-line Bases
Magnesium hydroxide Mg(OH)2
Calcium hydroxide Ca(OH)2
Strotium hydroxide
Sr(OH)2
Barium hydroxide
Ba(OH)2
Weak Base
Kb~ 102-103
Kb~ 0.01 to0.1
partially ionized
Ammonia
NH3
Kb=1.79x10-5 pKb = 4.74
Ethyl amine CH3CH2NH2 Kb=5.6x10-4 pKb = 3.25
Acid and Base Strength
• Strong acids Ionize completely in water.
HCl, HBr, HI, HClO3,
HNO3, HClO4, H2SO4.
• Weak acids
Partially ionize in water.
Most acids are weak.
• Strong bases Ionize completely in water.
Strong bases are metal
hydroxides - NaOH, KOH
• Weak bases Partially ionize in water.
Common Acids and Bases
Acids
nitric
hydrochloric
sulfuric
acetic
Bases
ammonia
sodium hydroxide
*undiluted.
Formula Molarity*
HNO3
16
HCl
12
H2SO4
18
HC2H3O2 18
NH3(aq)
NaOH
15
solid
Autoionization of Water
Autoionization When water molecules react with one
another to form ions.
H2O(l) + H2O(l)
H3O+(aq) + OH-(aq)
(10-7M) (10-7M)
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
ion product
of water
pH and other “p” scales
Substance
pH
need to measure and use 0.0
acids and bases
1 MWe
HCl
over
a very large concentration
range.
Gastric
juices
1.0 - 3.0
Lemon
juice
2.2 -track
2.4 of
pH and
pOH are systems to keep
Classic
Coke
these
very large ranges. 2.5
CoffeepH
= -log[H3O+5.0
]
Pure Water
pOH
= -log[OH-]7.0
BloodpH + pOH
7.35 - 7.45
= 14
Milk of Magnesia
10.5
Household ammonia
12.0
1M NaOH
14.0
pH scale
A logarithmic scale used to keep track of the large
changes in [H+].
0
7
10-14 M
Very
acidic
Basic
10-7 M 10-14 M
Neutral
14
Very
When you add an acid to, the pH gets smaller.
When you add a base to, the pH gets larger.
pH of some common materials
Substance
pH
1 M HCl
Gastric juices
Lemon juice
Classic Coke
Coffee
Pure Water
Blood
Milk of Magnesia
Household ammonia
1M NaOH
0.0
1.0 - 3.0
2.2 - 2.4
2.5
5.0
7.0
7.35 - 7.45
10.5
12.0
14.0
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 pOH
The 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
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
Substance
pH
HNO
(aq) + H2O(l)
H3+O(aq)
1 M 3HCl
0.0+ NO3-(aq)
Gastric juices
- 3.0
Therefore,
the moles of H+ ions 1.0
in the
solution is
Lemon
- 2.4
equal tojuice
moles of HNO3 at the 2.2
beginning.
Classic
[HNO3Coke
] = [H+] = 0.2 mole/L 2.5
Coffee
5.0
pH = -log [H+]
Pure Water
7.0
= -log(0.2)
Blood
7.35 - 7.45
pH = 0.699
Milk of Magnesia
10.5
Household ammonia
12.0
1M NaOH
14.0
pH of 0.5 M H2SO4 Solution
H2SO4(aq) + H2O(l)
(aq)
H3+O(aq) + HSO4-
HSO4-(aq) + H2O(l)
H3+O(aq) + SO42-(aq)
[H3+O][HSO4-]
H2SO4 ; Ka1 = ------------------[H2SO4]
[H3+O][SO42-]
H2SO4 ; Ka2 = ------------------- ; Ka2 ignored
pH of 0.5 M H2SO4 Solution
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
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 + ABase: BOH + H2O
B+ + OH[H3+O][ A-]
[B+ ][OH-]
Ka = --------------- ; Kb = ----------------[HA]
[BOH]
Acid Dissociation Constant
HCl(aq) + H2O(l)
Ka=
[H3+O][Cl-]
----------------[HCl]
+
[H ][Cl-]
Ka=
----------------[HCl]
H3+O(aq) + Cl-(aq)
Base Dissociation Constant
NH3 + H2O
K =
NH4+ + OH-
[NH4+][OH-]
[NH3]
Hydrated Metal Ions as Acids
[Fe(H2O)6]3+ (aq) + H2O ( )
[Fe(H2O)5(OH)]2+ (aq) + H3O+ (aq)
[Fe(H2 O) 5 (OH) 2 ][H3 O ]
3
Ka 

6.3
10
Fe(H2 O) 3
6
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
b. HOCl2
c. HOCl
d. HCN
; Ka= 4.0 x 10-4.
; Ka= 1.2 x 10-2.
; Ka= 3.5 x 10-8.
; 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)
Ka Examples
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][SO42-]
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]
How do you calculate pH of
weak acids/bases
From % dissociation
From Ka or Kb
What is % dissociation
Amount dissociated
% Dissoc. = ------------------------- x 100
Initial amount
How do you calculate %
dissociation from Ka or Kb
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
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 %
% Dissociation gives x (amount
dissociated) need for pH calculation
Amount dissociated
% Dissoc. = ------------------------- x 100
Initial amount/con.
x
% Dissoc. = --------------------------- x 100
concentration
Calculate the pH of a weak acid from %
dissociation
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
Weak acid Equilibria
Example
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
[H3O+][Bz-]
Ka =
[HBz]
Weak acid Equilibria
HBz
Initial conc., M
0.10
Change, DM
-x
H3O+
0.00
x
Bz-
0.00
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 x
Ka = 6.28 x 10-5 =
x
x2
0.10
= (6.28 x 10-5 )(0.10)
H3O+ = 0.0025 M
pH = 2.60
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
Weak Acid Equilibria
[H 3+O ][CN-]
Ka = --------------[HCN]
=
x2
---------------1.0 - x
1.0 - x ~ 1.00 since x is small
x2
Ka = -----------; Ka = 4.9 x 10-10 =
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
x2
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 NH Cl
Hydrolysis
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.
Hydrolysis
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+.
Acid-Base Properties of Typical Ions
What salt solutions would be
acidic, basic and neutral?
1)
2)
3)
4)
strong acid + strong base = neutral
weak acid + strong base = basic
strong acid + weak base = acidic
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 M
Change
-x
x
x
Eq. Con. 0.5 - x
x
x
[H 3+O ] [NH3 ]
Ka(NH4+) = -------------------=
[NH 4+]
x2
---------------- ;
appro.:0.5 - x . 0.5
(0.5 - x)
Continued
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 M
pH = -log [H+ ] = - log 1.66 x 10-5
pH = 4.77
pH of 0.5 M NH4Cl solution is 4.77 (acidic)
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
H
H
H
O
C
C
C
C
H
H
H
O
nonacidic hydrogens
H
acidic hydrogen
butanoic acid
O
H
C
3
electron-attracting
oxygen atom
C
O
H
C
OH
acetic acid
acidic hydrogen
3
O
C
H
C
O
3
-
-
C
OH
acetate ion
Organic or Carboxylic Acids
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
Amines
Class of organic bases derived
from ammonia NH3 by replacing
hydrogen by organic groups.
They are defined as bases
similar to NH3 by BronstedLowery or Lewis acid/base
definitions.
Amines
Acid-Base Chemistry
of 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
A Typical Synthetic Detergent Molecule
H
C
H 2CH2C
H 2CH2C
H 2C
H 2C
H 2C
H 2CH2C
H 2CH2C
H 2
3CH2C
SO3-Na+
Water-soluble part
(hydrophilic)
Oil-soluble part
(hydrophobic)
A nonionic detergent
H
C
H
3(C
H
(C
O
2)4CO
hydrocarbon
chain
(hydrophobic)
(
2)2O
CH2C
H
) 2C
H 2C
H 2O
H
2O
alcohol group
(hydrophilic)
ester
link
(hydrophilic)
ether
link
ether
link
(hydrophilic)
Dishwashing Detergent
Lewis Definition
G.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  NH4
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)