ACIDS AND BASES
GENERAL PROPERTIES
OF ACIDS
- Acids have sour taste
- Acid causes color change in plant dyes. It changes blue litmus paper to red.
- Acids react certain metals such as zinc, magnesium and iron to produce hydrogen gas.
Ex. 2HI(aq)+ Mg(s) → MgCl2(aq)
+ H2(g)
- Acids react with carbonates and bicarbonates such as Na2CO3, CaCO3, and NaHCO3 to produce carbon dioxide gas.
Ex.
2HCl(aq) + CaCO3(aq) →CaCl2(aq) + H2O(1)
+ CO2(g)
- Aqueous acid solutions conduct electricity
GENERAL PROPERTIES OF BASES
- Bases have bitter taste
- Bases feel slippery
- Bases causes color change to plant dyes. It changes red litmus paper to blue
- Aqueous base solutions conduct electricity
Release of H+ or OH-
and the Arrhenius Acid-Base Definition
Arrhenius acid-base definition,
acids and bases are classified in terms of their formulas and their behavior in
water. Svante Arrhenius (father of ionization theory) postulated that when
molecules of electrolytes dissolve in water, positive and negative ions are
formed.
An acid is a substance that has H in
its formula. Arrhenius acid is a substance that dissolve in water to give H+
ion or substance that dissociate in water to produce hydronium ion, H3O+.
A base is a substance that has OH in
its formula. Arrhenius base is a substance that dissociates in water to yield
OH-.
Some typical Arrhenius acids are HCl,
HNO3, and HCN, and some typical bases are NaOH, KOH, and Ba(OH)2.
When an acid and a base react, they undergo neutralization. In the Arrhenius
sense, neutralization occurs when
the H+ ion from the acid and the OH- ion from the base combine to
form H2O.
Proton Transfer and the
Bronsted-Lowry Acid-Base Definition
A major shortcoming of the
Arrhenius acid-base definition: many substances that yield OH- ions
when they dissolve in water do not contain OH in their formulas, like ammonia.
Another limitation of the Arrhenius definition was that water had to be solvent
for acid-base reactions. J.N. Bronsted and T.M. Lowry suggested definitions
that remove these limitations. According to the Bronsted-Lowry acid-base
definition, an acid is a proton
donor, any species that donates an H+ ion. A base is a proton acceptor,
any species that accepts an H+.
From the Bronsted-Lowry perspective,
the only requirement for an acid-base reaction is proton transfer process.
LEWIS ACIDS AND BASES
Gilbert N. Lewis suggests a more
general definition of acids and bases. He defined Lewis base as a substance
that can donate a pair of electrons (electron pair donor). A Lewis acid is a
substance that can accept a pair of electrons (electron pair acceptor).
- Amphoteric/Amphiprotic substances – substances that can function as an acid and as a base, like water.
HCl(g)
+ H2O(l) ⇌
Cl-(aq)
+ H3O+(ag) water
acts as a base
NH3(aq) + H2O(l)
⇌
NH4+(aq)
+ OH- (aq)
water act as acid
- Water is a unique solvent because of its special properties to act as either acid or base.
STRENGTHS OF ACID AND BASES
Strong acids/bases are strong electrolytes
which ionized completely in water (100% dissociated in water), while weak
acids/bases ionized only to a limited extent in water (slightly ionized).
A. Strong
Acids
1.
The hydrohalic acids HCl, HBr and HI
2.
Oxoacids in which the number of O atoms
exceed the number of ionizable H atoms by two or more such as HNO3, H2SO4
and HClO4.
B. Weak
Acids
1.
The hydrohalic acid HF
2.
Acids in which H is not bonded to O or
to halogen such as HCN and H2S.
3.
Nionizable H atoms such as HCLO, HNO2
and H3PO4.
4.
Organic acids with the general formula
RCOOH like CH3COOH.
C. Strong
Bases
1.
M2O or MOH; where M = group
IA metals; such as NaOH and KOH
2.
MO or M(OH)2; where M = Group
IIA metals; such as Mg(OH)2 and Ba(OH)2.
D. Weak
Bases
1.
Ammonia (NH3)
2.
Amines with the general formula RNH2,
R3N such as CH3NH2, (CH3)2NH
and (C3H7)3N.
pH – A MEASURE OF ACIDITY
The pH of a solution is defined as
the negative logarithm of the hydrogen ion concentration (in mol/L):
pH = -log[H+]
Like equilibrium constant, the pH of
a solution is a dimensionless quantity.
Since pH is simply a way to express
hydrogen ion concentration, acidic and basic solutions at 25 ºC can be distinguished
by their pH values, as follows:
Acidic
solutions: [H+]> 1.0 x 10-7 M, pH<7.00
Basic solutions: [H+] <
1.0 x 10-7 M, pH> 7.00
Neutral solutions: [H+] =
1.0 x 10-7 M, pH = 7.00
Notice
that pH increases as [H+] decreases.
A pOH scale analogous to the pH
scale can be devised using the negative logarithm of the hydroxide ion
concentration of a solution.
pOH
= -log [OH-]
THE SELF-IONIZATION OF WATER
Water is a unique solvent. It is a
very weak electrolyte and therefore a poor conductor of electricity, but it
undergoes ionization to a small extent.
H2O
H+ OH-
This reaction is sometimes
called the autoionization of water. We can write the equilibrium constant for
the autoionization of water:
Kw = [H+][OH-]
The equilibrium constant Kw is
called the ion-product constant for water or water ionization constant,
which is the product of the molar concentration of H+ and OH- ions at a
particular temperature.
In pure water at 25 C, [H-] = [OH-]
= 1.0 x 10-7 M. Thus from the equation
Kw = [H+] [OH-]
= (1.0 x 10-7) (1 x 10-7)
Kw = 1.0 x 10-14
Consider
the ion product for water:
Kw = [H+][OH-]
= 1.0 x 10-14
Taking
the negative log:
(-log[H+])
+ (-log[OH-]) = - log 1.0 x 10-14
(-log
1.0 x 10-7) + (-log 1.0 x 10-7) = - log 1.0 x 10-14
pH + pOH = 14
Practice Problems:
1.
The concentrated OH- in a certain
household ammonia cleaning solution is 0.0038 M. Calculate the concentration of
the H+ ions.
2.
In a NaOH solution, [OH-] is
2.9 x 10-4 M. Calculate the pH of the solution.
PREDICTING THE ACID- BASE
PROPERTIES OF SALT SOLUTIONS
Cation from
|
Anion from
|
|
Neutral
|
Strong Base
|
Strong acid
|
Acidic
|
Weak base or
metal ion with high + charge
|
Strong Acid
|
Basic
|
Strong base
|
Weak acid
|
Example:
Predict
whether the following aqueous solutions will be acidic, basic or neutral.
a. KI
b. NH4Cl
c. CH3COOK
Practice Exercise:
Predict
whether the following aqueous solutions will be acidic, basic or neutral.
a. NH4I
b. CaCl2
c. KCN
d. Fe(NO3)
RECORDED LECTURES
Acids and Bases Chemistry - Basic Introduction
