Reactions in Solution
The most important substance on earth is water.
In chemistry, water is necessary for many reactions to take place.
Table salt (NaCl) when put into water dissolves into its ions, Na+ and Cl-.
Water is the solvent and NaCl, is the solute.
The mixture is an aqueous solution.
The Water Molecule allows many
substances to be dissolved in them.
One side of the water molecule is
negatively charged, and the other side is
positively charged. Water is a polar
molecule.
Electricity and Solutions
A useful characteristic of solutions is the ability to
conduct electricity. To determine if a solution has the
ability to conduct electricity, an electrical conductivity
apparatus is used. An electrical conductivity apparatus
is basically a battery and light bulb setup which lights
up when electricity is conducted through the solution.
Electrolytes are substances that produce ions upon
dissolving. There are two ways to provide these
mobile ions for conducting purposes.
1. Dissociation of Ionic Compounds
2. Ionisation of Polar Covalent Molecular Substances
Electrolytes
1.
Dissociation of Ionic Compounds:
Ionic compounds are made of cations and anions.
These ions are locked into position in their crystal structure and are not
able to move around.
In water, the water molecules, are attracted to the ions. The ions are said to
be dissociated, and able to carry electrical particles to conduct current.
K3PO4 + H2O ----> 3K+ (aq) + PO4-3 (aq)
Such substances are said to be electrolytes.
Salts that are completely soluble in water are strong electrolytes.
Salts that are slightly soluble are weak electrolytes at best.
The strength of an electrolyte is measured by its ability to conduct
electrical current.
Electrolytes
2.
Ionisation of Polar Covalent Molecular Substances
Polar molecular substances are substances whose atoms are covalently
bonded. Each molecule has a net molecular dipole moment and thus a
positive and a negative end.
Polar water molecules can line up around polar molecule. If this dipoledipole interaction can overcome the dissociation energy of a bond the
molecule will fragment with bonding electrons going with the most
electronegative atom in the broken bond, creating ions.
(Electronegativity is the electron attracting ability of an atom)
Such polar molecular compounds are called electrolytes.
An example of a strong electrolyte is any of the strong acids, such as HBr.
H-Br + H2O ----> H3O+ (aq) + Br- (aq)
Electrolytes
Some polar molecular substances have such strong covalent bonding that water
is only able to overcome these stronger dissociation energies in a portion of the
molecules.
CH3COOH + H2O H3O++CH3COOFor example,a weak acid such as ethanoic acid, CH3-COO-H, dissolves in
water with only a small percentage of the molecules being ionized.
Non-electrolytes are substances that do not
produce ions when they dissolve.
This results when polar molecular substances
are large enough and their covalent bonding is
strong enough so that water is not able to break
any of the covalent bonds during the solvation
process. As a result, the neutral molecules are
solvated (separated by solvent water molecules)
without any ionization occurring.
Acids and Bases
The properties of acids include the following:
•
Taste sour (but don't taste them!!)
•
Their water solutions conduct electrical current (electrolytes)
•
They react with bases to form salts and water
•
Turns Blue Litmus Paper to Red
The properties of bases include the following:
•
Have a slippery feel between the fingers
•
Have a bitter taste (but don't taste them!!)
•
React with acids to form salts and water
•
Turns Red Litmus Blue
•
Their water solutions conduct electrical current (electrolytes)
Acids and Bases
Arrhenius in 1884 discovered that acids give off H+ ions and
allow for a good flow of electricity through a solution.
Arrhenius also discovered that bases give off OH- ions and
OH- ions also allow for a good flow of electricity through the
solution.
Traditionally Professor Arrhenius defined:
Acid released Hydrogen ion (as Hydronium ions,
H3O+) in water solution.
Base produced Hydroxide ion in water solution.
The limitations on these definitions were:
1. The need for water
2. The need for a protic acid
3. The need for Hydroxide bases
Bronsted/Lowry acids and bases
Bronsted and Lowry defined these two terms the following:
Acid-Proton donor Base-Proton acceptor
These definitions are not as restrictive as Arrhenius’ definitions.
1. No need for water although it can be present, it need not be.
2. Bases do not have to be Hydroxide compounds.
However, one restriction still remaining is the need for a protic acid.
Each Bronsted acid is coupled to a
conjugate base to constitute a
CONJUGATE ACID-BASE PAIR
CH3COOH + H2O H3O++CH3COO-
Lewis Acids and Bases
G.N. Lewis defined these in an even less restrictive manner:
Acid- Electron pair acceptor
Base- Electron pair donor
In this set of definitions there is no longer a need for a protic
acid. In other words only electron exchange must occur.
These definition sets are NOT contradictory. A Proton donor is the same as
an electron acceptor. A Proton acceptor is the same as an electron donor. Also
the first set of definitions are less inclusive so that all of the Arrenhius acids
are found under the Bronsted definition but not all Bronsted acids will be
Arrenhius acids. All Arrenhius and Bronsted acids will be under the Lewis
definition but not all Lewis acids will be Bronsted or Arrenhius acids.
Acid and Base Strength
Strong acids (memorise) dissociate completely in water
HClO4, HCl, HBr, HI, HNO3 and H2SO4
Strong bases are the metal hydroxides of Group 1 and 2
E.g.
LiOH, NaOH, KOH, Ba(OH)2, Mg(OH)2 etc
Weak acids and bases are not completely ionised in solution
CH3COOH + H2O H3O++CH3COO-

H O CH COO 


Ka
3

3
CH3COOH
Ka is an equilibrium constant
called the
acid dissociation
constant
Acid and Base Strength
(a molecular base)

NH OH 


:NH3 + H2O NH4++OH-
Kb

4
: NH3 
The magnitude of the Ka or Kb, using water as a
common proton donor/acceptor, determines the
strength of the acid or base
In general (for acids)

H O A 


HA + H2O H3O++A-
Ka

3
HA 
Water is AMPHOTERIC. It can act as an acid or a base
Acid and Base Strength
Ka
Stronger
Acid
~1010
HClO4
H2SO4
HCl
H3O+
1x10-2
HSO46.8x10-4
HF
CH3COOH 1.75x10-5
9.5x10-8
H2S
5.7x10-10
NH4+
HCO34.7x10-11
H2O
1.8x10-16
ClO4Levelling
HSO4Effect
ClEach acid
H2O
will transfer
SO42a proton to a
Fbase below it
CH3COOin a mixed
HSsolution
NH3
CO32Stronger
OHBase