Sách Hóa Tiếng Anh 11

Chapter 1
DISSOCIATION
 What will happen when acids, bases and salts dissolve
in water?
 What are the characteristics of chemical reactions
which occur in aqueous solutions?

Arrhenius (S.Arrhenius, 1859 – 1927),
a Swedish, awarded the Nobel Prize in
Chemistry in 1903.

A type of pH-meters used in
laboratories.


Lesson 1:

DISSOCIATION
 To know what dissociation is and what an
electrolyte is.
 To know what a strong electrolyte is and what a
weak electrolyte is.

I. DISSOCIATION
1. Experiment
Prepare three beakers: (a) containing distilled water, (b) containing sucrose
solution (C12H22O11) and (c) containing sodium chloride solution (NaCl), then mix them
apparatus as shown in Figure 1.1.

Figure 1.1: The apparatus for proving the electrical conductivity of a solution
When connecting the terminals of the electrical wire to the same electrical source,

we see that only the lightbulb at the beaker containing NaCl solution lights up. Thus,
NaCl solution conducts electricity, but the distilled water and the sucrose solution do not.
In the similar experiment, it can be seen that, dried NaCl solid, dried NaOH solid,
ethanol solution (C2H5OH), glicerol (C3H5(OH)3) do not conduct electricity. In contrast,
solutions of acids, bases and salts conduct electricity.
2. The reason for the electrical conductivity of aqueous solutions of acids, bases and
salts


In 1887, Arrehenius proposed a theory which was later confirmed by experiments.
The theory states that: Solutions of acids, bases and salts conduct electricity because
there are charged particles called ions moving freely in their solution.
Since acids, bases and salts dissociate into ions when dissoleved in water, their
solutions conduct electrcity.
The process of separation of substances into ions is called dissociation. Substances
that give ions when dissolved in water are called electrolytes. Thus, acids, bases and salts
are electrolytes.
The dissociation is represented by a dissociation equation:
NaCl → Na+ + ClFor example:

HCl → H+ + ClNaOH → Na+ + OH-

II. CLASSIFICATION OF ELECTROLYTES
1. Experiment
Prepare two beakers: one containing 1.10M HCl solution, the other containig
0.10M CH3COOH solution and fix them into the apparatus as shown in Figure 1.1. When
connecting the terminals of electrical wire to the same electrical source, we see that the
lightbulb at the beaker containing HCl solution is brighter than the one at the beaker
containing CH3COOH solution.
This proves that the concentration of ions in HCl solution tn higher than that of

ions in CH3COOH solution, that is, the number of HCl molecules dissociating into ions is
greater than that of CH3COOH molecules separating into ions.
2. Strong electrolytes and weak elestrolytes
a. Strong electrolytes
A strong electrolytes is a substance that when dissolved in water, all of its
dissolved molecules completely dissociate into ions.
For example, NaCl is a strong electrolyte. If there are 100 molecules of NaCl
dissolved in a solution, all 100 molecules dissociate into ions.


Strong acids such as HCl,HNO3, HClO4, H2SO4…, strong bases such as NaOH,
KOH, Ba(OH)2, Ca(OH)2,… and most salts are strong electrolytes.
In the dissociation equations of strong electrolytes, a unidirectional arrow is used to
designate thedirection of the dissociation process.
For example:

Na2SO4 → 2Na+ + SO2-

Since Na2SO4 dissociates completety, the concentration of ions for dissociated
Na2SO4 can be easily calculated.
For example, in 0.01M Na2SO4 solution, the concentration of Na+ ions is 0.02M
and the concentration of SO2- is 0.01M.
b. Weak elestrolytes
A weak electrolytes is a substance that when dissolved in water, only a fraction of
their dissolved molecules dissociates into ions, the rest still exists as molecules in the
solution.
For example, in 0.043M CH3COOH, for every 100 dissolved molecules, only 2
molecules dissociate into ions, 98 other molecules do not dissociate. Thus, CH3COOH is
a weak electrolyte.
Weak acids such as CH3COOH, HClO, H2S, HF, H2SO3,… and weak bases such

as Bi(OH)3, Mg(OH)2,… are weak electrolytes.
A bidirectional arrow is used in the dissociation equations of weak electrolytes.
For example: CH3COOH → CH3COO- + H+
The dissociation of weak electrolytes is a reversible process. When the rate of the
dissociation of ions and that of the recombination of ions are equal, the equilibrium of
the dissociation is established. Dissociation equilibrium is a dynamic equilibrium. Like
other chemical equilibriums, dissociation equilibrium also obeys the Le Chatelier’s
principle.

EXERCISES


1. Why is it that the solutions of acids such as HCl, bases such as NaOH and salts such as
NaCl to conduct electricity, while the solutions such as ethanol, sucrose and glyxerol
solutions do not?
2. What is the dissociation? What is an electrolyte?
What types of substances are electrolytes? What are strong electrolytes? What are weak
electrolytes? Give examples and write their dissociation equations.
3. Write the dissociation equations of the following substances:
a) Strong electrolytes: Ba(NO3)2 0.10M ; HNO3 0.02OM ; KOH 0.010M. Calculate the
molarity of each ion in the above solutions.
b) Weak electrolytes: HClO3, HNO2.
4. Choose the correct answer in the following sentences:
The solution of electrolyte conducts electricity due to
A. The moving of electrons.
B. The moving of cations.
C. The moving of dissolved molecules.
D. The moving of both cations and anions.
5. Which of the following substances does not conduct electricity?
A. Dried, solid KCl

B. Molten CaCl2
C. Molten NaOH
D. HBr dissolved in water

Lesson2:

ACIDS, BASES AND SALTS


 To know what are acids, bases, amphoteric hydroxides
and salts according to Arrhenius theory and write their
dissociation equations.
I. ACIDS
1. Denifition
According to Arrhenius theory, an acid is a substance that dissolves in water and
dissociates to produce H+ cations.
For example:

HCl → H+ + ClCH3COOH

€

CH3COO- + H+

Acidic solutions have some common properties, which are the properties of H+
cations in the solutions.
2. Polyprotic acids
From two above examples, it can be seen that, the molecule of HCl, as well as the
molecule of CH3COOH, only dissociates into one H+. They are monoprotic acids.
When dissolved in water, acids whose molecules dissociate stepwise into H+ ions

are called polyprotic acids.
For example:

H3PO4
H2PO

4

HPO4

2-

€
€
€

H2PO4- + H+
HPO42- + H+
PO43-

+ H+

H3PO4 molecules dissociate through 3 steps to ptoduce H+ ions, so H3PO4 is a triprotic
acid.
II. BASES
According to Arrhenius theory, when dissolves in water, a bases is a substance
which dissociates into OH- anions.
For example:

NaOH → Na+ + OH-

Bases solutions have some some common properties, which are the properties of
OH anions in the solutions.
-

III. AMPHOTERIC HYDROXIDES
An amphoteric hydroxide is a hydroxide which can dissociate as an acid and also
as a base when dissolves in water.
For example, Zn(OH)2 is an amphoteric hydroxide:
Dissociation as a base:
Zn(OH)2

€

Zn2+ + 2OH-

Dissociation as an acid:
Zn(OH)2

€

ZnO22- + 2H+

Zn(OH)2 is often written as H2ZnO2, which implies that it is an acid.
Cammon amphoteric hydroxide are Zn(OH)2, Al(OH)3, Sn(OH)2 and Pb(OH)2.
They are all moderately soluble in water anh their acidity (ability to dissociate into ions)
and basicity are low.
IV. SALTS
1. Denifition

A salt is a compound which dissociates into cations of a metal (or NH4+ cation) and
acidic anions when dissolving in water.
For example:
(NH4)2SO4 → 2NH4+ + SO42NaHCO3 → Na+ + HCO3Salts whose acidic anions do not contain hydrogen atoms that can dissociate into
H ions (acidic hydrogen) are called neutral salts. For example: NaCl, (NH4)2SO4 ,
Na2CO3.
+


If acidic anions of salts still contain hydrogen atoms which can dissociate into H+
ions, the salts are called acidic salts. For example: NaHCO3, NaH2PO4, NaHSO4.
2. Dissociation of salts in water
Most salts, when dissolved in water, dissociate completely into cations of metals
(or NH4+ cation) and acidic anions (except some salts such as HgCl2, Hg(CN)2…are
weak electrolytes).
For example:

K2SO4 → 2K+ + SO42NaHSO3 → Na+ + HSO3-

If acidic anions contain acidic hydrogen atoms, these anions further dissociates
weakly to produce H+ ions.
For example:
HSO3-

→ H+ + SO32-

EXERCISES
1. State the definitions of acids, monoprotic acids, polyprotic acids, bases, amphoteric
hydroxides, neutral salts, and acidic salts. Give the illustrative examples and write their
dissociation equations.

2. Write the dissociation equations of the following substances:
a) Weak acids: H2S, H2CO3.
b) Strong base: LiOH
c) Salts: K2CO3, NaClO, NaHS.
d) Amphoteric hydroxide: Sn(OH)2
3. Which conculsion is correct according to the Arrhenius theory?
A. A compoud which contains hydrogen atoms as its component is an acid.
B. A compoud which contains OH groups as its component is a base.
C. A compoud which is able to dissociate into H+ cation in water is an acid.


D. A base does not have to contain OH groups in its molecules.
4. For the 0.01M of weak CH3COOH acid solution, if ignoring the dissociation of water,
which of the following judgements about molarity of ions is correct?
A. [H+] = 0,01M

C. [H+] > [CH3COO-]

B. [H+] < [CH3COO-]

D. [H+] < 0,01M

5. For the 0.01M of strong HNO3 acid solution, if ignoring the dissociation of water,
which of the following judgements about molarity of ions is correct?
A. [H+] = 0,01M

C. [H+] > [NO3-]

B. [H+] < [NO3-]

D. [H+] < 0,10M


Lesson 3 : THE

DISSOCIATION OF WATER. pH.

ACID-BASE INDICATORS
 To know how to evaluate acidity and basicity of solutions based on
the concentration of H+ and pH.
 Know the colors of some indicators in the solution at different
ranges of pH.
I. WATER IS A VERY WEAK ELECTROLYTE
1.The dissociation of water
By using sensitive equipment, it can be seen that water conducts electricity very weakly.
Water dissociates very minimally.
H2O

ƒ

H+ + OH-

(1)

Experiments have verified that at ambient temperature, for every 555 millions of H2O
molecules, there is one molecule dissociating into ions.
2. Ion product of water
From the dissociation equation of H2O (1), it can be seen that one H2O molecule ionizes
into one H+ ion and one OH- ion, that is, in water, the concentration of H+ equals to that
of OH-. Water is neutral, so that we can define:

Neutral medium is the medium in which [H+] = [OH-].
By experiment, their concentrations can be determined as follows:
[H+] = [OH-] = 1.010-7 (mol/L) ở 250C


Assign KH2O ( 25oc ) = [H+] . [OH-] = 1.010-7 1.010-7 = 1.0 10-14
The product KH2O = [H+] . [OH-] is called ion product of water. This product is a constant
at definite temperature. How ever, the value of ion product of water
1.0 10-14 is usually used in calculation at around 25oC.
Approximately, the value of ion product of water can be considerded as a constanteven in
dilute solutions of different substances.
3. The significance of the ion product of water
a) In acidic medium
When dissolving an acid in water, the concentration of H+ increases, therefor the
concentration of OH- must decrease so that the ion product of water is constant. For
example, dissolving HCl acid in water to get the H+ concentration of
1.0 10-3M , the concentration of OH- is:
[OH-] = = M
Thus, an acidic medium is a medium with:
[H+] > [OH-] or [H+] > M
b) In alkaline medium
When dissolving a base in water, the concentration of OH- increases, therefor the
concentration of H+ must decrease so that the ion product of water is constant. For
example, dissolving an acid in water to get the OH- concentration of 1.0 10-5M, the
concentration of H+ is:
[H+] = = M
Thus, an alkaline medium is a medium with:


[H+] < [OH-] or [H+] < M

The examples above indicate that, if the concentration of H+ in a aqueous solution is
given, the concentration of OH- can be calculated and vice versa. Thus, the acidity and
basicity of a solution can be evaluated only by the concentration of H+:
In neutral medium:

[H+] = M

In acidic medium:

[H+] < M

In alkaline medium:

[H+] > M

II. THE CONCEPT OF pH. ACID-BASE INDICATORS
1. The concept of pH
As mentioned above, the acidity and basicity can be evaluated by the concentration of
H+. However, the H+ concentration of common solutions is usually small. To avoid
writing the concentration of H+ with negative power, the value of pH is used according to
the following convention.
[H+] = M. If [H+] = M then pH=a
For example:
[H+] = M

pH = 2.00 : acidic medium

[H+] = M

pH = 7.00 : neutral medium

[H+] = M

pH = 10.00 : alkaline medium

The common pH scale has the values from 1 to 14.
The value of pH is of grate practical significance. For instance, pH of human and animal
blood has the almost constant value. Every kind of plants can grow normally when the


pH of solutions in soil is in the definite range. The speed of corrosion of metals in natural
water depends much on the pH value of water.
2. Acid-base indicators
Acid-base indicators are substances which change color with the pH value of solutions.
For example, the colors of two acids-base indicators, litmus and phenolphthalein, in
solutions with different values of pH are given in the table 1.1.
Litmus
Phenolphthalein

Red
Purple
pH6
pH = 7.0
pH < 8.3 pH < 8.3
colorless

Blue
pH 8
pH 8.3 pH 8.3
pink

When mixing some indicators with colors changing sucessively over the pH values,
universal indicator is obitaned. Using a band of paper immerged in this mixture, we can
determine the approximate value of the solution pH (figure 1.2)

Figure 1.2. The color of universal indicator ( MERCK reagent from Germany)
over different pH values
pH- Meters are used to determine the pH value of solutions relatively precisely.
EXCERCISES
1. What is the ion product of water? What is its value at 25oC?


2. State the difinitions of acidic, neutral and alkaline (basic) solutions based on H+
concentration and pH.
3. What is an acid-base indicator? Give the colors of litmus and phenolphthalein in the
solutions over different ranges of pH.
4. A solution has [OH -] of 1.5 x 10-5M. This solution is
A.
B.
C.
D.

acidic
neutral
alkaline
unknown

5. Calculate the concentrations of H+, OH- and pH of a 0.10M HCl solution and a 0.010M
NaOH solution.
6. In the 0.010M HCl solution, the ion product of water is

A.
B.
C.
D.

[H+] [OH-] > 1.0 x 10-14
[H+] [OH-] = 1.0 x 10-14
[H+] [OH-] < 1.0 x 10-14
unknown

Material
pH VALUES OF SOME COMMON LIQUID SOLUTIONS
Sample

pH

Gastric juice

1.0 2.0


Lemon juice
Vinegar

3.0

Grape juice
Orange juice
Urine

4.8 7.5

Water in the air

5.5

Saliva

6.4 6.9

Milk

6.5

Blood

7.3 7.45

Tears

7.4

Lesson 4 : ION-EXCHANGE

REACTIONS

IN THE SOLUTIONS OF ELECTROLYTES


 To understand the nature of ion-exchange reactions and the

conditions for ion-exchange reactions in the solutions of
electrolytes and write the net ionic equations of the reactions.
I - THE CONDITION OF ION-EXCHANGE REACTIONS IN THE SOLUTIONS
OF ELECTROLYTES
1. Reactions forming precipitates
Experiment. When a sodium sulfate, ( Na2SO4), solution was dropped into a
test-tube containing barium chloride, ( BaCl2 ), solution, a white precipitate, appeared.
Na2SO4 + BaCl2 → BaSO4 ↓ + 2NaCl

(1)

Explanation. Both Na2SO4 and BaCl2 are soluble and completely dissociate in water.
Na2SO4 → 2Na+ + SO42BaCl2

→ Ba2+ + 2Cl-

Among the four dissociated ions, only Ba2+ and SO42- ions combine to form the BaSO4
precipitate (Figure 1.3), therefor, the reaction in the solution indeep is:
Ba2+ + SO42- → BaSO4 ↓

Figure 1.3. BaSO4 precipitate

(2)


The equation (2) is called net ionic equation of the reaction (1).
The net ionic equation shows the nature of the reaction in the solutions of electrolytes.
The conversion of molecular equation into net ionic equation is done as follows:
Convert all soluble substances with strong electrolytes into ions, and write gases,
precipitates and weak electrolytes in molecular form. The obitained equation is called

full ionic equation, for instance, for the reaction (1), we have:
2Na+ + SO42- + Ba2+ + 2Cl- → BaSO4↓ + 2Na+ + 2ClReducing ions which do not participate in the reaction, we obtain net ionic equation:
Ba2+ + SO42- → BaSO4 ↓
From this equation, it can be seen that, to prepare BaSO4 precipitate, it is necessary to
mix two solutions, one with Ba2+ ions and the other with SO42- ions.
2. Reaction forming weak electrolytes
a) Reaction forming water
Experiment. When some drops of a phenolphthalein solution were added to a beaker
containing a 0.10M NaOH solution, the solution had a pink color ( Figure 1.4).

Figure 1.4. The color of phenolphthalein in the alkaline solution


Slowly pour a 0.10M HCl solution into the above beaker as you stir until the color of the
solution disappears. The reaction occurs as follows:
HCl + NaOH → NaCl + H2O
Explanation. Both NaOH and HCl are soluble and dissociate completely in water:
NaOH → Na+ + OHHCl

→ H+ + Cl-

OH- in the solution causes phenolphthalein to turn pink. When a HCl solution was added,
H+ ions of HCl reacted with OH- ions of NaOH to form the weak electrolyte H2O. The
net ionic equation is:
H+ + OH- → H2O
The color of the solution disappears when the H+ ions of HCl react completely with the
OH- ions of NaOH.
The reactions between acid solutions and basic hydroxides occur easily because of the
formation of the week electrolyte H2O. For instance, Mg(OH)2 is moderately solution in
water but easily soluble in a strong acid solution:

Mg(OH)2 (s) + 2H+ → Mg2+ + 2H2O
b) Reaction forming weak acids
Experiment. When a HCl solution was dropped into a test-tube containing CH3COONa
solution, the weak CH3COOH acid was formed:
HCl + CH3COONa → CH3COOH + NaCl
Explanation. HCl and CH3COONa are well-soluble and well-dissociated substances:


HCl → H+ + ClCH3COONa → Na+ + CH3COOIn the solution, H+ ions will combine with CH3COO- to form the weak CH3COOH
electrolyte (with vinegar smell). The net ionic equation is:
H+ + CH3COO- → CH3COOH
2. Reaction forming a gas
Experiment. Pouring HCl solution into a beaker containing Na2CO3, we see some gas
bubbles escaping:
2HCl + Na2CO3 → 2NaCl + CO2↑ + H2O
Explanation. Both HCl and Na2CO3 are well soluble ang well dissociated:
HCl → H+ + ClNa2CO3 → 2Na+ + CO32H+ and CO32- ions in the solution combine into the weak acid H2CO3, which is unstable
and decomposed into CO2 and H2O.
H+ + CO32- → HCO3H+ + HCO3- → H2CO3
H2CO3 → CO2↑ + H2O
Net ionic equation:
2H+ + CO32 → CO2↑ + H2O


The reactions between carbonate salts and acid solutions occur easily because of the
formation of the weak electrolyte H2O and the formation of CO2 gas separate from the
reaction solutions. For instance, carbonates are moderately soluble in water but well
soluble in acid solutions. For example, limestone (CaCO3) dissloves easily in the HCl
solution (Figure 1.5):
CaCO3(r) + 2H+ → Ca2+ + CO2↑ + H2O

Limestone

Figure 1.5. Reaction forming CO2 gas

II – CONCLUSION
1. The reaction in solutions of electrolytes is the reaction between ions.
2. The ion-exchange reaction in solutions of electrolytes occurs only when the ions
combine into at least one of the following substances:
−
−
−

a precipitate
a weak electrolyte
a gas


EXCERCISES
1. What is the condition for ion-exchange reactions in the solutions of electrolytes to
occur? Give illustrative examples.
2. Why do the reactions between acidic solutions and basic hydroxides, and those
between carbonates and acid solutions occur easily?
3. Give some examples to prove that the nature of the reactions in solutions of
electrolytes is the reaction between ions.
4. A net ionic equation shows:
A.
B.
C.
D.

What ions exist in the solutions.
Ions with the highest concentretion in a solution.
The nature of the reaction in the solution of electrolytes.
Molecules do not exist in the solution of electrolytes.

5. Write the molecular equation and net ionic equation of the reactions (if any) occurring
in the solutions of the following pairs of substances:
a) Fe2(SO4)3 + NaOH

d) MgCl2 + KNO3

b) NH4Cl + AgNO3

e)FeS(s) + HCl

c) NaF + HCl

f) HClO + KOH

6. Which of the following reactions occurring in the solutions to from Fe(OH)3
precipitate (Figure 1.6)
A.
B.
C.
D.

FeSO4 + KMnO4 + H2SO4
Fe2(SO4)3 + KI
Fe(NO3)3 + Fe

Fe(NO3)3 + KOH


Figure 1.6. Fe(OH)3 precipitate
7. Give the examples and write the molecular and net ionic chemical equations for the
following reactions:
a)
b)
c)

Forming precipitates.
Forming weak electrolytes.
Forming gases.

Lesson 5

REVIEW
ACIDS, BASES AND SALTS
ION – EXCHANGE REACTIONS
IN THE SOLUTIONS OF ELECTROLLYTES


 To consolidate the knowledge about acids, bases and
the condition for ion-exchange reactions in the solutions
of the electrolytes.
 To practice the skill of writing net ionic equations of
the reactions.
I.KNOWLEDGE TO MASTER
1.
2.

3.

When dissolved in water, acids dissociate into H+ ions.
When dissolved in water, bases dissociate into OH- ions.
When dissolved in water, amphoteric hydroxides can dissociate as acids as well as

4.

bases.
Most of salts when dissolved in water, dissociate completely into ions of a metal
(or NH4+ ions) and anions of an acidic radical.
If the acidic radical contains acidic hydrogen, it further dissociates moderately into
H+ cations and anions of an acidic radical.

5.

The ions product of water is KH2O=[H+].[OH-]=1.0*10-14 (at 25oC). Approximately,
the value of this product can be considered a constant even in dilute solutions of

6.

other substances.
The typical [H+] and pH values in media are:
In neutral medium: [H+] = 1.0*10-7M or pH = 7,00.
In acidic medium: [H+] > 1.0*10-7M or pH < 7,00.
In alkaline medium: [H+] < 1.0*10-7M or pH > 7,00.

7.

The colors of litmus, phenolphthalein and universal indicators in solutions over

8.

different pH values are shown in Table 1.1 and Figure 1.2.
The ion-exchange reaction in solutions of electrolytes occurs only when ions
combine into at least one of the following substances:
- Precipitates.
- Weak electrolyates.


Gases.
Net ionic equations show the nature of reactions in the solutions of electolytes.
-

9.

In the net ionic equations, ions which do not participate in the reactionare
removed, and precipitates, weak electrolytes and gases are kept in the molecular
form.
II. EXERCISES
1.

Write the dissocication reactions of the following substances:
K2S, Na2HPO4, NaH2PO4, Pb(OH)2, HBrO, HF, HClO4.

2.

A solution has the [H+] concentration of 0.010M. Calculate the [OH-] and the
pH of the solution acidic, neutral or alkaline? Show the color of purple litmus

3.

in this solution.
A solution has the pH of 9.0. Calculate the molarity of H+ and OH-ion in the

4.

solution. Show the color of phenolphtalein in this solution.
Write the molecular and net ionic equations of the reactions (if any) which
occur in the solutions of the following substances:

a, Na2CO3 + Ca ( NO3 ) 2
b, FeSO4 + NaOH ( dilute )
c, NaHCO3 + NaOH
d , NaHCO3 + NaOH
e, K 2CO3 + NaCl
g , Pb(OH ) 2( s ) + HNO3
h, Pb(OH ) 2( s ) + NaOH
i, CuSO4 + Na2 S
5.

An ion-exchange reaction in the solution of electrolytes occurs only when
A. The reactants are well soluble
B. The reactants are strong electrolytes
C. Some ions in the solutions combine to reduce their ionic concentration


6.

The reaction is nonreversible

Which of the following pairs of substances forms CdS precipitate (Figure 1.7)

7.

in the solution?
A. CdCl2 + NaOH
B. Cd(NO3)2 + HCl
C. Cd(NO3)2 + HCl
D. CdCl2 + Na2SO4
Write the (molecular and net ionic) chemical equations of the ion-exchange

D.

reactions in the solutions forming each of the following precipitates: Cr(OH)3,
Al(OH)3, Ni(OH)2 (Figure 1.7b, c, d).

Chapter 2 : NITROGEN

– PHOSPHORUS

 How are the positions of nitrogen and phosphorus in the
Periodic Table related to their atomic and moclecular
structures?
 What are the basic properties of nitrogen and phosphorus
elements and their compounds ? How to explain the properties
based on the theory you have studied ?
 How to prepare nitrogen , phosphorus and important
compounds ?

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