ACIDS, BASES, & SALTS
H ow to get a weak acid out of its salt? E.g. hydrofluoric
acid, HF, out of KF.
Answer:
T r e a t t h e s a l t wi t h a s t r o n g a c i d , e . g . wi t h
H2SO4: KF(aq) + H2SO4(aq)

HF (aq)+KHSO4
+
Net ionic: F + H HF
Same approach to get a weak , or low-soluble base, out of
the salt:
Treat the salt solution
with a strong base (alk ali):
CuCl2(aq) + KOH(aq)

Cu(OH)2 + 2KCl(aq)
2+
Net Ionic: Cu (aq) +2OH (aq)

Cu(OH)2
OXYACIDS & BASES:
A Unifying View
Oxyacids & Bases are all element hydroxides:
E-O-H
i.e. produced by combination of an oxide with
water:
K2O+H2OK-OHK++OH-
Cl2O+H2OHO-ClH++ClO2 ways of E-O-H ionization:
E-O-HEO + H
+
Base: E-O-HE + OH
Acid:
-
+
Which way prevails? depends on the relative polarity of
two bonds:
+
EO- vs .-OH+
The more polar will be ionized with water:
Cl & O are close in electronegativity, Cl – O is a low-polar
bond, while H – Cl is highly polar bond, therefore:
H – O – Cl → H+ + ClO-
In KOH, H – O bond is less polar than ionic K O bond,
therefore: KOH → K+ + OH-
PERIODIC TRENDS
in BASICITY - ACIDITY
The most active metals form basic oxides
& metal hydroxides (alkali).
Non-metals form acidic oxides & acids
Transition metals may form basic, or acidic oxides & hydroxides.
Basicity-acidity depends on the oxidation number of the
element:
Higher oxidation number  higher acidity
(more oxygen  higher acidity)
MnO, basic oxide
Mn(OH)2 a base
Mn2O7 acidic oxide
HMnO4 permanganic acid & its salts, as KMnO4
CrO basic oxide,
Cr(OH)2 a base
CrO3 acidic oxide,
H2CrO4 Chromic acid & its salts, as Na2CrO4
Arrhenius definitions of acids & bases were
generalized to include non-aqueous
solutions & dry chemistry reactions.
BRÖNSTED-LOWRY definition:
Acid is any compound that can serve as a
proton donor:
HCl(g)+H2OH3O+(aq)+Cl-(aq)
HCl(g) is proton (H+) donor to water
HCl(g)+NH3(g)NH4Cl(s)
HCl(g) is proton donor to ammonia.
No n e e d fo r w a te r o r o th e r s o l v e n t!
Brönsted BASE
is a proton acceptor
:NH3(g) accepts a proton,
it is a base
Water is an acid when reacts with
ammonia:
H2O + NH3(g)NH4+(aq) + OH-(aq)
In the reverse reaction:
NH4+(aq) is acid, OH-(aq) is base.
B-L acid-base reaction is a proton
transfer.
In a neutralization reaction:
+
H3O
acid
+ OH  2H2O
OH
base
Anions of weak acids are Brönsted bases:
they produce OH- with water:
CO32- + H2OHCO3- + OHthis is
hydrolysis reaction.
salts of weak acids
& strong bases, when dissolved, produce
It is the reason why
alkaline solutions:
Na2CO3  2Na+ + CO32CO32- +H2O HCO3- + OH-
BRÖNSTED Acid-Base PAIRS
A CONJUGATE base is the species that
remains after a Brönsted acid has given up H+:
HCN(aq)+H2OH3O+(aq)+CN-(aq)
ACID
CONJUGATE BASE
A CONJUGATE acid is the species formed when a
Brönsted base gains a proton:
NH3(aq)+H2ONH4+(aq)+OH-(aq)
BASE
H2CO3+H2O 
acid
1
base
2
CONJUGATE
ACID
H3O+ + HCO3-
conjugate
conjugate
acid
2
base
1
Reactions between acids & bases trend to form
weaker bases & acids, from stronger bases &
acids:
HCl+NaHCO3H2CO3+NaCl
or, in net ionic form: H++HCO3-H2CO3
pH
logarithm reminder:
a
a = log10N means 10 = N
these are two ways to present the same statement.
(They are as similar as: 3=6/2 & 2=6/3)
+
pH  -log10[H ]
or:
if [H+] = 10-3 M, pH=3, or pH 3
if [H+] = 2×10-3 M, pH 2.7
if pH 2, [H+] = 10-2 M
if pH -1, [H+] = 10 M
+
- pH
[H ] = 10
pH of common liquids and body fluids:
Battery electrolyte
pH -1
Human gastric juice
1-2
Lemon juice
2.2-2.4
Vinaigre
2.4-3.4
Carbonated drinks
2-4
Black coffee
3.7-4.1
Tomato juice
4.0-4.4
Urine
5.5-7.0
Cow’s milk
6.3-6.6
Saliva
6.5-7.5
Human blood
7.3-7.5
Sea water
7.8-8.3
Bile
7.8-8.8
Mono Lake
10-10.5
0.1 M Sodium carbonate
1 1 .7
1 M NaOH
14
pH
&
SELF-IONIZATION of WATER
Water is a very weak electrolyte. It is self-ionized by
proton transfer:
H2O + H2O  H3O+ + OHIn pure water:
[H3O+] = [OH-] = 10-7 mol/L
[H2O] = (1000g)/(18g/mol) =55.5mol/L
10-7<<<55.5, & is always the same in diluted aqueous solutions.
When an acid is added to water, the concentration of protons it
releases is so much higher than what water would produce on its
own, that the latter can be ignored.
In a [HCl] = 0.1 mol/L, HCl  H+ + ClHCl, as a strong electrolyte, gives
[H+] = [HCl] = 0.1 mol/L, pH = -log0.1 = 1
In a 0.01 M HNO3, HNO3  H+ + NO3[H+] = [HNO3] = 0.01 mol/L, pH = -log0.01 = 2
In 1.0 M HCl: HCl  H+ + Cl-
[H+]= [HCl] = 1.0 mol/L,
p H = -l o g 1 = 0 (1 0 0 = 1 )
Water self-ionization, as a source of protons, may be neglected compared to
HCl ionization.
When acid is added, [H3O+] (or [H+]) increases, while [OH-]
goes down.
However:
the product of both [H+] & [OH-] concentrations is a
constant independent of each of these concentrations.
At room temperature:
Kw = [H+].[OH-] = 10-14 M2
H2O  H+ + OHbackward rxn will occur when H+ & OH- collide.
T h e c o l l i s i o n s f re q u e n c y i s p ro p o rt i o n a l t o t h e i r c o n c e n t ra t i o n s , i . e .
the rate of the backward rxn:
R = k [H+][OH-]
The rate of the forward rxn:
R = k [H2O]
At equilibrium:
R = R
k [H2O] = k [H+][OH-]
[H+][OH-]
is equilibrium
K=
constant
[H2O]
or
Kw = K [H2O] = [H+][OH-]
Since [H2O] = 55.5 always, we put it into the constant.
Kw = [H+][OH-] = 10-14
Ionic product for water
In pure water: [H+] = [OH-] = 10-7 M
In a 1.0 M HCl, [H+] = 1.0 M
In a 1.0 M KOH
[OH-] = 1.0 M
Since Kw = [H+] [OH-] = 10-14
[H+] = Kw/[OH-] = 10-14/1.0 = 10-14M
i.e. pH = -log(10-14) = 14
The range of possible [H+] is about 16 orders of magnitude wide.
T h e re f o re ,
logarithmic scale is convenient.
pH = -log[H+]
Range of pH values: -1 to 15.
Neutral solutions:
Ac i d i c s o l u ti o n s :
Al k a l i n e s o l u ti o n s :
[H+] = 10-7M,
low pH,
h ig h p H
pH 7
pH<7
pH>7
For strong acids & bases in concentrations that are exact power of 10,
pH is self-evident:
[HCl] = 10-2M, [H+] = 10-2M, pH 2
[KOH] = 10-3M. [OH-] = 10-3M, p[OH-]=3
where by definition:
pOH = - log[OH-]
Since Kw=[H+][OH-] = 10-14
pH + pOH = 14
 pH =14 - pOH = 14 - 3 = 11
When the concentration differs from integer power of 10, use
calculator.
[HCl] = 3.4x10-5M. pH-? pH = -log(3.4x10-5) = 4.47
[KOH] = 0.001 M
KOH  K+ + OH- [OH-] = 0.001 M = 10-3 M
pOH = 3, pH = 14-3=11 or [H+] = 10-11 M
pH INDICATORS
are weak acids HIn  H+ + Inwhich have different colors of their neutral molecular form
HIn as compared to that of anion InDuring titration of an alkali, KOH, with phenolphtalein, initial
KOH solution has color of In- anion (pink), since excess of
alkali shifts the equilibrium
HI n  H + I n
+
-
When alkali is neutralized by the acid, pH decreases sharply, &
the color changes to that of HIn (colorless for phenolphtalein).
This moment is equivalence point.
Phenolphthalein is an example of an indicator which establishes this type of equilibrium in aqueous solution:
In-
HIn
-H+
+H+
Colourless (Acidic)
Rasberry red (Alkaline)
Transition occurs at pH9
Other indicators have different transition interval.
Colour
Indicator
Acid
Base
pH range
Thymol Blue - 1st change
red
yellow
1.2 - 2.8
Methyl Orange
red
yellow
3.2 - 4.4
Bromocresol Green
yellow
blue
3.8 - 5.4
Methyl Red
yellow
red
4.8 - 6.0
Bromothymol Blue
yellow
blue
6.0 - 7.6
Phenol Red
yellow
red
6.8 - 8.4
Thymol Blue - 2nd change
yellow
blue
8.0 - 9.6
Phenolphthalein
colourless
pink
8.2 - 10.0
How pH is estimated?
Paper strips are pre-socked with indicator solutions covering the entire pH range:
Another way: use an electronic instrument: pH meter
(was invented by Beckman in California, in 1930-s, who combined the glass electrode with then new electronic amplifyer. This
instrument, initially developed for CA orange juice companies to determine the quality of juice by estimating the content of
ascorbic acid in it, started his multi-billion buisiness. Beckmann was a major benefactor of American research & education:
Beckman institutes, etc.)
voltameter
Reference
electrode
pH sensitive glass electrode
Dr. Arnold Beckman (1900-2004)
10” long