Transcript Acids & Bases

CHAPTER 14: ACIDS & BASES
Chem
1212
Dr. Aimée Tomlinson
Section 14.1
Acid-Base Concepts:
The Brønsted-Lowry Theory
Three Theories for Acids & Bases
Arrhenius acids & bases
acid: an H + donor
base: an O H - donor
H A ( aq )
M O H ( aq )


H ( aq )  A ( aq )


M ( aq )  O H ( aq )
Brønsted-Lowry acids & bases


H ( aq )  A ( aq )
acid: an H + donor
H A ( aq )
base: H + acceptor
H ( aq )  B ( aq )


B H ( aq )
Lewis acids & bases
W e w ill see this type at the end of the chapter
Conjugate Acid-Base Pairs

conjugate acid: the acid that is created after the Brønsted-Lowry
base has accepted the proton, BH+

conjugate base: the base that is created after the Brønsted-Lowry
acid has donated the proton, A-

Examples basic : N H 3 ( g )  H 2 O ( l )
base
base

conj  acid
acid
acidic : H C l  H 2 O
acid

N H 4 ( aq )  O H ( aq )
Cl

conj  base
 H 3O
K b  1.76  10
conj  base

conj  acid
K a  1
5
Section 14.2
Acid & Base Strength
Strong Acids

A strong acid will completely dissociate/ionize:



All the reactant goes to product/single-headed arrow
The product is a very weak conjugate acid/base pair
List of Strong acids: HCl, HBr, H2SO4, HI, HClO4, HClO3, HNO3
Strong Bases

A strong base will completely dissociate/ionize:

List of Strong bases: MOH (M=alkali metal), NH2-, H-
Weak Acids
Only partially dissociate

H N O 2 ( aq )





H ( aq )  N O 2 ( aq )
K a  4.0  10
4
The eq constant is called Ka where “a” for acid
There is always some reactant still present at eq
unlike the strong acid case
The larger the Ka the stronger the acid


E.g. Ka >> 1 for HNO3
We will come back to this in a little bit
Section 14.3
Hydrated Protons &
Hydronium Ions
Meet Hydronium
 H3O+ is acidified water or what truly happens when
H+ is in H2O
 We call this ion hydronium
 We use H+ and H3O+ interchangeably
Amphoterism
Defn: A species that may act as both an acid and a base
Water as a base:
H A ( aq )  H 2 O ( l )


H 3 O ( aq )  A ( aq )
Water as an acid:
B ( aq )  H 2 O ( l )


B H ( aq )  O H ( aq )
Section 14.4
Dissociation of Water
What’s in Water & What it Means
2𝐻2 𝑂
𝑙
⇌ 𝑂𝐻−𝑎𝑞 + 𝐻3 𝑂+𝑎𝑞
𝐾𝑐 = [𝑂𝐻−𝑎𝑞 ][𝐻3 𝑂+𝑎𝑞 = 𝐾𝑤
NOTE H2O(l) as always is not in the equilibrium expression
Relationship between [OH-(aq)] and [H3O+(aq)]:






[ H 3 O ( aq ) ]  [ O H ( aq ) ]
[ H 3 O ( aq ) ]  [ O H ( aq ) ]
[ H 3 O ( aq ) ]  [ O H ( aq ) ]
acidic
neutral
basic
For both ions their concentrations at 298 K is 1.0 x 10-7M
making the Kw = 1.0 x 10-14
K w  (1.0  10
7
)(1.0  10
7
)  1.0  10
 14
Example
Determine the hydroxide concentration in a solution with
𝐻3 𝑂+𝑎𝑞 =1.89 x 10-4 M.
Section 14.5
The pH Scale
Power of Hydrogen aka pH

pH   log  H 3 O  aq  



pH < 7.0 acidic

pH = 7.0 neutral

pH > 7.0 basic
Power of Hydroxide aka pOH

pH   log  O H  aq  



pH > 7.0 acidic

pH = 7.0 neutral

pH < 7.0 basic
Relationship for pH, pOH & Kw
Section 14.6
Measuring pH
pH Indicators
More relevant in Chapter 15 so we will address it more fully there
Section 14.7
pH in Solutions of Strong
Acids & Strong Bases
The Strong Completely Dissociate
H3O+/OH- concentrations will become whatever those of the strong
acids or bases were
Example I
EXAMPLE: Write the balanced equation for each of the following
and determine the pH.
1.) 0.5000 M HClO4(aq)
2.) 0.0256 M LiOH(aq)
Example II
Determine the hydronium ion concentration for a 0.01500 M
Ca(OH)2 assuming complete dissociation.
Section 14.8
Equilibria of Weak Acids
Weak Acids & Equilibrium
Unlike the strong they only partially dissociate in water hence HA is
still present at eq:
H A( aq )  H 2 O ( l )


H 3 O ( aq )  A( aq )
or
H A( aq )


H ( aq )  A( aq )
 H 3 O (aq )   A(aq ) 
 H (aq )   A(aq ) 






Ka 

 HA
HA
Ka & Acid Strength
The larger the Ka:

More strongly the eq will lie
toward product

More likely the acid is to
dissociate

The larger the [H3O+]

The lower the pH

The stronger the acid

Ka is large for strong acid HCl
but very small for weak acid
CH3OH
Section 14.9
Calculating the Equilibria
of Weak Acids
Weak Acids & Equilibrium
Calculate [H+] and the pOH of 0.050M of benzoic acid. Ka = 6.5 x 10-5
Weak Acid Flowchart
Section 14.10
Percent Dissociation
of Weak Acids
Percent Dissocation
Degree of ionization/dissociation: percentage that an acid ionizes

H A( aq )
H

( aq )
 A

( aq )
[H ]
 100%
[ H A]
Example: Determine the percent dissociation of 0.050M of benzoic
acid.
Section 14.11
Polyprotic Acids
Polyprotic Acids
Acids which possess more than one proton
Polyprotic Acid Example
Calculate the [H+] of 0.050M of sulfuric acid.


H 2 SO 4 ( aq )  H ( aq )  H SO 4 ( aq )

H SO 4 ( aq )

2
H ( aq )  SO 4 ( aq )
K a  1
K a  1.2  10
2
Polyprotic Acid Flowchart
Why Ka1 > Ka2


H 2 SO 4 ( aq )  H ( aq )  H SO 4 ( aq )

H SO 4 ( aq )

2
H ( aq )  SO 4 ( aq )
K a  1
K a  1.2  10
2

Electrostatically it is more difficult to remove H+ from SO42- than
from HSO4-

Hence Ka2 is always smaller than Ka1 and so on
Section 14.12
Equilibria of Weak Bases
Weak Base Equilibria
Calculate pH of 0.050M of ammonia. Kb = 1.8 x 10-5
Weak Base Flowchart
Section 14.13
The Relationship Between
K a & Kb
The Link Between Ka & Kb is Kw
HCN  H 2O
CN

CN
 H 2O
2 H 2O
 H 3O
HCN  OH
H 3O



 OH

Ka  Kb  Kw


Ka
Kb
Kw
Example
Determine the Kb of HCN if Ka = 4.9 x 10-10.
Section 14.14
Acid/Base Properties of
Salts
Stronger Partner Dominates



Strong acid + weak base = acidic solution
Weak acid + strong base = basic solution
Strong acid + strong base = neutral solution
Example: Classify each of the following as acidic, basic, or neutral.
1.) KBr
2.) NaNO2
3.) NH4Cl
What if both are weak?
Example II: Classify NH4CN as acidic, basic, or neutral.
Finding pH/pOH of a Salt Solution
Calculate the pH of a 0.25M NaC2H3O2, Ka = 1.76x10-5
Salt Flowchart
Section 14.15
Factors that Affect Acid
Strength
Recall Electronegativity Trend
EN Trend I
Increasing acid strength going down the table:




As we go down a column we decrease EN
We thereby weaken the H-X bond
Allows H+ to more readily go into solution
Acid strength: HF < HCl < HBr < HI
EN Trend II
Increasing acid strength from left to right in the table:




As we go across we increase EN
We make the H-X bond polar
This eventually gives an EN difference which leads
to H+
Acid strength: CH4 < NH3 < H2O < HF
Oxoacids Trend I – more EN
An oxoacid is any acid with acidic proton connected
to an O-atom – they have the form HnXOm




As increase the EN of the halogen X we weaken
the O-H bond
This is done by pulling electron density from the Oatom
This will allow the H+ to break-away more eqsily
and go into solution
Acid strength: HOI < HOBr < HOCl < HOF
Oxoacids Trend II
Increasing the number of O-atoms increases acid
strength




As increase the number of O-atoms weakens the
O-H bond
Again this is done by pulling electron density from
the O-atom
This will allow the H+ to break-away more easily
and go into solution
Acid strength: HOCl < HO2Cl < HO3Cl < HO4Cl
Acid
Oxidiation State of Cl
Ka
HClO
HClO2
HClO3
HClO4
+1
+3
+5
+7
2.9 x 10-8
1.1 x 10-2
1
1 x 108
Amine Base Trends
Increasing the number of electro-donating groups will
increase base strength
N H 3  N H 2 C H 3  N H  C H 3  2  N  C H 3 3
Increasing the number of electron-withdrawing/EN
groups will decrease base strength
N H 3  N H 2 C l  N H C l 2  N C l3
Section 14.16
Lewis Acids and Bases
Definitions
Lewis Acid
Electron-pair acceptor
N H 3  B F3  H 3 N B F 3
Lewis Base
Electron-pair donor