1. The solvent is generally in excess.
AQUEOUS SOLUTIONS
A solution is a homogeneous mixture of a solute dissolved in a solvent. In aqueous solutions solvent is WATER.
The solvent is generally in excess.
Example
The solution NaCl(aq) is
sodium chloride NaCl(s) dissolved in water H2O(l)
The solute is NaCl(s) and the solvent is H2O(l)

2. Aqueous Reactions
Aqueous reactions can be grouped into three general categories, each with its own kind of driving force:
Precipitation reactions
Acid base neutralization reactions
Oxidation-reduction reactions.

3. Electrolyte and Non-electrolyte
Electrolyte: a substance that conducts electricity when dissolved in water.
Acids, bases and soluble ionic solutions are electrolytes.
Non-electrolyte: a substance that does not conduct electricity when dissolved in water.
Molecular compounds and insoluble ionic compounds are non-electrolytes.

4. Electrolytes Some solutes can dissociate into ions.
Electric charge can be carried.

5. Types of solutes Strong Electrolyte - 100% dissociation,
high conductivity
Strong Electrolyte -
100% dissociation,
all ions in solution
Na+
Cl-

6. Types of solutes Weak Electrolyte - partial dissociation,
slight conductivity
Weak Electrolyte -
partial dissociation,
molecules and ions in solution
CH3COOH
CH3COO- H+

7. Types of solutes Non-electrolyte - No dissociation,
no conductivity
Non-electrolyte -
No dissociation,
all molecules in solution
sugar

9. Table 4.1 Electrolyte Classification of Some Common Substances
Strong Electrolytes Weak Electrolytes Nonelectrolytes
HCl, HBr, HI CH3COOH H2O
HClO4 HF CH3OH
HNO C2H5OH
H2SO C12H22O11(sucrose)
KBr Most organic compd
NaCl
NaOH, KOH
Other soluble
ionic compounds

10. Ionic Equations
The molecular equation does not tell us that the reaction actual involves ions in solution. So the soluble ionic substances in solution should be represented by their separate ions.
Other ions are SPECTATOR IONS that do not take part in reaction.
A strong electrolyte:
MgCl2(s) → Mg2+(aq) + 2 Cl- (aq)
Strong – complete dissociation
Weak – reversible
A weak electrolyte:
CH3COOH(aq) ← CH3COO -(aq) +H+(aq) →
CH3COOH(aq), CH3OH(aq), H2O(l), CO2(g), Zn(s), MgCO3(s)
A non-electrolyte:

11. Writing Net Ionic Equation
Overall /Molecular Equations: Complete formulas are written for all the reactants and products, no ions are written.
AgNO3(aq) +NaI (aq) → AgI(s) + NaNO3(aq)
Complete ionic equation: Strong electrolytes are written in their ionized forms and weak/non-electrolytes as unionized form.
Ag+(aq) + NO3-(aq) + Na+(aq) + I-(aq) →
AgI(s) + Na+(aq) + NO3-(aq)
Spectator ions
Ag+(aq) + NO3-(aq) + Na+(aq) + I-(aq) →
AgI(s) + Na+(aq) + NO3-(aq)
Net ionic equation: Write only those chemical species which are involved in a chemical reaction. All spectator ions are eliminated.
Ag+(aq) + I-(aq) → AgI(s)

12. Example of Ionic equations
Suppose copper (II) sulfate reacts with sodium sulfide.
Write out the chemical reaction and name the precipitate.
CuSO4 (aq) + Na2S (aq) CuS (s) + Na2SO4 (aq)
Write out the net ionic equation.
Cu+2 (aq) SO4-2 (aq) + 2Na+ (aq) + S-2 (aq) CuS (s) + 2Na+ + SO4-2 (aq)
Cu+2 (aq) + S-2 (aq) CuS (s)

13. Precipitation Reactions
Precipitation reactions are process in which soluble reactants yield an insoluble solid product that falls out of solution. Most precipitations take place when certain cations and anions combined to produce an insoluble ionic solid called a precipitate.
2Ag NO3 (aq) + Na2CO3 (aq)  AgCO3(s) + 2NaNO3
Soluble Cations
Alkali metals: Li+, Na+, K+, Rb+, Cs+,
NH4+
Mostly insoluble
Metal (other than alkali metals) sulfides, hydroxides carbonates, phosphates
Soluble Anions:
Halides: Cl-, Br-, I- except of Ag+, Hg2+2, Pb+2
NO3-, ClO4-, CH3CO2-, SO4-2 except SO4-2 of Ba+2, Hg2+2, Pb+2

14. Example (a) Al2(SO4)3 + NaOH 
i) write down the reactants and interchange of anions to get product
Al2(SO4)3 + 6NaOH  2Al(OH)3 + 3Na2SO4
All common Na compounds are water soluble Na+ remain in solution. The combination of Al3+ and OH- produce insoluble Al(OH)3. Then the ionic equation is
2Al3+ +3SO Na+ + 6OH-2Al(OH)3(s)+ 6Na+ + 3SO42-
The net ionic equation is :
Al OH-Al(OH)3(s)

15. Acid-Base Reactions
An acid is a substance that provides hydrogen ions (H+) (increase the concentration of H+) in aqueous solution. H+ is too reactive to exit by itself, it attaches to water to give the more stable hydronium ion, H3O+.
A base is a substance that produces hydroxide ions (OH-) (increase the conc. of hydroxide ions) in aqueous solution.
HA (aq)  H+(aq) + A-(aq)
an acid HA is a general formula for an acid
MOH(aq)  M+(aq) + OH-(aq)
a base MOH is a general formula for a base

16. Strong and Weak Acids and Bases
A strong acid is an acid that is almost completely ionized in aqueous solution.
A weak acid is an acid that only partially ionized (as result of an equilibrium reaction with water) in aqueous solution.
HCl(aq)  H+(aq) + Cl-(aq) strong acid
CH3CO2H(aq)  H+(aq) + CH3CO2-(aq) weak acid
A strong base is a base that dissociate nearly completely in aqueous solution.
A weak base is a base that is only partially ionized (as result of an equilibrium reaction with water) in aqueous solution.
NaOH(s) Na+(aq) + OH-(aq) strong base
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) weak base

17. Weak Acids
CH3COOH     HCOOH         HF                    HCN                HNO2              H3PO4-       
Weak Bases
NH3
CH3NH2  

18. Acid-Base Neutralization Reaction
In a neutralization reaction, an acid and a base react to form water and an aqueous solution of an ionic compound called a salt.
A neutralization reaction:
HA(aq) + MOH(aq)  H2O(l) + MA(aq)
acid base water salt
HCl(aq) + NaOH(aq) NaCl(aq) +H2O(l)
H+(aq) + Cl-(aq) +Na+(aq) + OH-(aq) Na+(aq) +Cl-(aq) + H2O(l)
By eliminating the spectator ions, we discover the actual Reaction of the neutralization of strong acid by a strong base
The net ionic equation:
H+(aq) + OH-(aq) H2O(l)
or H3O+(aq) + OH-(aq) 2H2O(l)

19. Oxidation-Reduction Reaction
Historically, OXIDATION referred to the combination of an element with oxygen to yield an oxide, and the word REDUCTION referred to the removal of oxygen from an oxide to yield the element.
Today, by using broader definitions, we look at the change of oxidation state of the element involved in the reaction.
An oxidation is defined as the lose of one or more electrons by a substance. (Increase in Oxidation number)
A reduction is the gain of one or more electrons by a substance. (Decrease in Oxidation number)
As a remembrance LEO says GER
Loss Electrons = Oxidation
Gain Electrons = Reduction

20. What is Redox? REDOX stands for REDuction/Oxidation
30/09/99
REDOX stands for REDuction/Oxidation
An oxidation and a reduction must occur together, and such a reaction is called an oxidation-reduction reaction or REDOX reaction. A redox reaction is a process in which electrons are transferred between substance or in which atoms change oxidation number.
Not all the redox reactions involve oxygen.

21. Oxidizing and Reducing Agents
An oxidizing agent: contains an element whose oxidation number decreases in a redox reaction by gaining electrons (it make possible for some other substance to be oxidized and itself reduced).
A reducing agent: contains an element whose oxidation number increases in a redox reaction by losing electrons (it make possible for some other substance to be reduced and itself oxidized).
In general, a substance with an element in one of its highest possible oxidation state is an oxidizing agent. If an element is in its lowest possible oxidation state, the substance is a reducing agent.

22. OXIDATION NUMBER Rules for determining oxidation numbers (ON):
An oxidation number is the APPARENT charge on an atom in a molecule or a compound.
Rules for determining oxidation numbers (ON):
1. The ON of any element in free (uncombined) state is zero. e.g. Na0, H20, O20 etc
2. The ON for any simple, monatomic ion is equal to the charge on the ion and it is same in compound as was in monoatomic ion e.g. +1 for Sodium in Na+,Na2CO3and NaCl
3. The sum of all the oxidation numbers of the atoms for neutral species must be equal to zero and for ions must be equal to the charge on the ion.
4. In its compounds, fluorine always has an ON of –1.
5. In its compounds, hydrogen has an ON of +1 except for metal hydrides, where the ON is -1.
6. In its compounds, oxygen has an ON of -2 except for the peroxides (-1), superoxides (-1/2) and OF2 (+2).