Chapter 8
ACIDS, BASIS, AND IONIC
OMPOUNDS
Definitions
• Electrolyte is a "medical/scientific" term for
salts, specifically ions. The term electrolyte
means that this ion is electrically-charged and
moves to either a negative (cathode) or positive
(anode) electrode
Defintions
• Ions that move to the cathode
(cations) are positively
charged
• Ions that move to the anode
(anions) are negatively
charged
• Electrolyses is the passage of
electricity through a solution
holding dissolved ions
• Electrolyte is a solute that
enables a solution to conduct
electricity
• Electrodes are the plates or
wires that dip into the solution.
• For example, your body fluids
blood, plasma, interstitial fluid
(fluid between cells) -- are like
seawater and have a high
concentration of sodium
chloride (table salt, or NaCl).
The electrolytes in sodium
chloride are:
• sodium ion (Na+) - cation
• chloride ion (Cl-) - anion
• Ionization is the gain or loss of electrons. The loss of
electrons, which is the more common process in
astrophysical environments, converts an atom into a
positively charged ion, while the gain of electrons
converts an atom into a negatively charged ion
• The ionization energy of an atom measures how
strongly an atom holds its electrons. The ionization
energy is the minimum energy required to remove an
electron from the ground state of the isolated gaseous
atom.
•
•
•
As electrons are removed, the positive charge from the nucleus remains unchanged,
however, there is less repulsion between the remaining electrons
Zeff increases with removal of electrons
Greater energy is needed to remove remaining electrons (i.e. the ionization energy
is higher for each subsequent electron)
Notation for Degrees of Ionization
Suffi
x
Ionization
Examples
Chemist
's
Not
atio
n
I
Not ionized
(neutral)
H I, He I
H, He
II
Singly ionized
H II, He
II
H+, He+
Doubly ionized
He III, O
III
He++,
O++
III
What is the process of solutes
when dissolved in water?
• Through ionization,
solutes release ions in
water
• Molecules mix readily
because both types of
molecules engage in
hydrogen bonding. Since
the intermolecular
attractions are roughly
equal, the molecules can
break away from each
other and form new
solute
Ammonia Dissolves in Water
solute (NH3), solvent (H2O) hydrogen bonds
What is the process of solutes
when dissolved in water?
• Alcohol Dissolves in
Water: The -OH group
on alcohol is polar and
mixes with the polar
water through the
formation of hydrogen
bonds. A wide variety of
solutions are in this
category such as sugar in
water, alcohol in water,
acetic and hydrochloric
acids.
Do ions carry electricity in
water?
• Yes, ions in water do carry electricity.
How?
• Redox reactions primarily involve the transfer of
electrons between two chemical species.
• The compound that loses an electron is said to be
oxidized, the one that gains an electron is said to be
reduced.
• A compound that is oxidized is referred to as a reducing
agent, while a compound that is reduced is referred to as
the oxidizing agent.
Redox Reactions (Oxidation –
Reductions)
For example: 2Fe3+ + Sn2+ -> 2Fe2+ + Sn4+ (8+ each side of the equation)
• Redox reactions are the
transfer of electrons
from one reactant to
another...
•
•
•
•
• Sn2+ donated electrons
to the Fe3+ (an electron
transfer took place).
In these reactions, the valency (oxidation number) of the reactants
change.
The iron (iii) + tin (ii) have reacted to give iron (ii) + tin (iv) of course, this
reaction is carried out in the presence of HCl (Hydrochloric Acid), but the
oxidation reduction reaction is only between the iron (iii) and tin (ii).
Now, a redox reaction is the release and uptake of electrons.
So, the Fe3+ is reduced to Fe2+, and the Sn2+ is oxidized to Sn4+.
Oxidation-Reduction
• When there is
oxidation, there is also
reduction.
•
• The substance which
loses electrons is
oxidized.
•
• The substance which
gains electrons is
reduced.
•
For example: Fe (metal) + Cu2+ > Fe2+ + Cu (metal)
Fe donates two electrons to the
Cu2+ to form Cu (metal). The Fe
lost 2 electrons, so is oxidized.
The Cu2+ gained 2 electrons, so
is reduced (in its valency).
Fe
Oxidised
Reducing
Agent
+
Cu2+
Reduced
Oxidising
Agent
> Fe2+
+
Cu
Redox Reactions involving acid
and bases solutions
• Not only are there an exchange of electrons in
these reactions, but also an exchange of protons
(hydronium ions), as in any base system
• CuS + HNO3 -> Cu SO4 + NO (g) + H2O
(equation not balanced).
• 3CuS + 8HNO3 -> 3 CuSO4 + 8NO(g) + 4H2O
(equation balanced)
• 3CuS2+ + 3S2- + 8H+ + 8NO3- -> 3Cu2+ +
3SO42- + 8NO(g) + 4H2O
Redox Reactions
Split into 2 separate steps.
•
•
•
•
2Fe3+ + 2e- -> 2Fe2+ (reduction)
(6+) + (2-) -> (4+) (balanced for
charges)
Sn2+ -> Sn4+ + 2e- (oxidation)
(2+) -> (4+) + (2-)
•
•
Add the two half equations:
2Fe3+ + 2e- + Sn2+ -> 2Fe2+ +
Sn4+ + 2eThe electrons cancel each other
out, so equation is: 2Fe3+ +
Sn2+ -> 2F2+ + Sn4+
By breaking down the equation
into half cells, the oxidation or
reduction of each chemical can be
determined. The atom which
gains electrons reduces its
valency, therefore is reduced and
is called the oxidizing agent.
The atom which loses electrons,
increases its oxidation number,
therefore is oxidized, and is
called the reducing agent.
Another Example
•
•
•
At the cathode: Cu2+(aq) + 2e- ----------) Cu(s)
At the anode:
2Br-(aq) ----------) Br2(l) + 2eSum:
Cu2+(aq) + 2Br-(aq) ----------) Cu(s) + Br2(l)
= Reduction
= Oxidation
= Electrolysis
• Acids taste sour, are corrosive to metals,
change litmus (a dye extracted from
lichens) red, and become less acidic when
mixed with bases
• Bases feel slippery, change litmus blue,
and become less basic when mixed with
acids
[H+]
Acids
Neutral
Bases
pH
Example
1 X 100
0
HCl
1 x 10-1
1
Stomach acid
1 x 10-2
2
Lemon juice
1 x 10-3
3
Vinegar
1 x 10-4
4
Soda
1 x 10-5
5
Rainwater
1 x 10-6
6
Milk
1 x 10-7
7
Pure water
1 x 10-8
8
Egg whites
1 x 10-9
9
Baking Soda
1 x 10-10
10
Tums® antacid
1 x 10-11
11
Ammonia
1 x 10-12
12
Mineral Lime - Ca(OH)2
1 x 10-13
13
Drano®
1 x 10-14
14
NaOH
Water as a base
• When an acid reacts with
water, the water behaves
as a proton acceptor to
form the hydronium ion.
• When a base (like
ammonia) reacts with
water, a proton is
transferred from water to
the ammonia molecule to
form the ammonium ion.
Therefore, water is
behaving as a proton
donor
Salts
• Ionic solids (or salts) contain positive and negative ions, which are
held together by the strong force of attraction between particles with
opposite charges. When one of these solids dissolves in water, the
ions that form the solid are released into solution, where they
become associated with the polar solvent molecules.
• H2ONaCl(s)Na+(aq) + Cl-(aq)
• We can generally assume that salts dissociate into their ions when
they dissolve in water. Ionic compounds dissolve in water if the
energy given off when the ions interact with water molecules
compensates for the energy needed to break the ionic bonds in the
solid and the energy required to separate the water molecules so
that the ions can be inserted into solution
Solubility Rules
1. Salts containing Group I elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this rule
are rare. Salts containing the ammonium ion (NH4+) are also soluble.
2. Salts containing nitrate ion (NO3-) are generally soluble.
3. Salts containing Cl -, Br -, I - are generally soluble. Important exceptions to this rule are halide
salts of Ag+, Pb2+, and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2 are all insoluble.
4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2) are common soluble salts of silver;
virtually anything else is insoluble.
5. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4,
Ag2SO4, and CaSO4.
6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble.
Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of
transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble.
7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, Ag2S are all
insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
8. Carbonates are frequently insoluble. Group II carbonates (Ca, Sr, and Ba) are insoluble. Some
other insoluble carbonates include FeCO3, PbCO3. Carbonates become soluble in acid solution.
9. Chromates are frequently insoluble. Examples: PbCrO4, BaCrO4
10. Phosphates are frequently insoluble. Examples: Ca3(PO4)2, Ag2PO4
11. Fluorides are frequently insoluble. Examples: BaF2, MgF2 PbF2.
Solubility example
• Take for example the reaction of lead(II) nitrate with sodium chloride
to form lead(II) chloride and sodium nitrate, shown below:
• Pb(NO3)2(aq) + 2 NaCl(aq) PbCl2(s) + 2 NaNO3(aq)
• This complete equation may be rewritten in ionic form by using the
solubility rules, lead(II) nitrate is soluble and therefore dissociated,
same about NaCl. As products, sodium nitrate is predicted to be
soluble and will be dissociated.
• The lead(II) chloride, however, is insoluble. The above equation
written in dissociated form is:
• Pb2+(aq) + 2 NO3-(aq) + 2 Na+(aq) + 2 Cl-(aq) PbCl2(s) + 2
Na+(aq) + 2 NO3-(aq)
• At this point, one may cancel out those ions which have not
participated in the reaction. Notice how the nitrate ions and sodium
ions remain unchanged on both sides of the reaction.
• Pb2+(aq) + 2 NO3-(aq) + 2 Na+(aq) + 2 Cl-(aq) PbCl2(s) + 2
Na+(aq) + 2 NO3-(aq)
• What remains is the net ionic equation, showing only those chemical
species participating in a chemical process:
• Pb2+(aq) + 2 Cl-(aq) PbCl2(s)