CHEMICAL BONDING

I. Choose the letter corresponding to the correct answer from the

options A, B, C and D given below:

(i) The type of bonding in X will be:

A. Ionic B. Electrovalent C. Covalent D. Molecular
(ii) X is likely to have a:
A. Low melting point and high boiling point
B. high melting point and low boiling point
C. Low melting point and low boiling point
D. high melting point and high boiling point
(iii) In the liquid state, X will:
A. Become ionic
B. Conduct electricity
C. Not conduct electricity

II. Name the following:

1. Process of formation of ions from molecules

2. A polar covalent compound
3. A gaseous molecule containing single covalent bond
4. Ions formed by loss of electrons
5. An inert gas which contains a triple bond
6. A regular tetrahedral molecule
7. A substance which conducts electricity in molten or aqueous state
8. A non-polar covalent compound
9. The number of electrons lost or gained by an atom refers to
10. During ionisation metals lose electrons, the change is called

III. Give reasons for the following:

1. Inert gases do not form ions.

2. Ionic compounds have a high melting point
3. Carbon tetrachloride does not conduct electricity
4. Hydrogen chloride can be termed as polar covalent compound
5. Solid sodium chloride does not conduct electricity
6. Bond formed in hydrogen molecule is non-polar.
7. Sugar solution does not conduct electricity.
8. Covalent compounds exist as gases, liquids and soft solids.
9. Why are ionic compounds crystalline solids?
10. Sodium atom is highly reactive, but sodium ion is not.
11. Why are polar covalent compound soluble in water

IV. Define the following:

1. Covalent bonding
 2. Ionic bonding
3. Co-ordinate bonding
4. Lone pair effect
5. Covalency
6. Electrovalency

V. Answer the questions:

1. State the type of bonding in the following molecules: [3]
a. Methane
b. Calcium oxide
c. Ammonium chloride
2. By drawing the dot diagram, show the lone pair effect leading to the
formation of ammonium ion from ammonia gas and hydronium ion from
water.
3. Element X, Y and Z have atomic number 6,9 and 12 respectively.
Which:
(i) Forms an anion
(ii) Forms a cation
4. What are the conditions for the formation of ionic bonds?
5. Write three differences between electrovalent compound and ionic
compound.
6. In the formation of magnesium chloride (by direct combination between
magnesium and chlorine), name the substance that is oxidized and the
substance that is reduced.
7.
Element W X Y Z
Electronic 2,8,1 2,8,7 2,5 1
configuration

(i) What type of bond is formed between:

(a) W and X
(b) X and Y
(ii) What is the formula of the compound formed between,
(a) X and Z
(b) W and X
8. M is a metal above hydrogen in the activity series and its oxide has the
formula has the formula M2O. This oxide when dissolved good
conductor of electricity. In the above context, answer the following:
(i) What kind of combination exist between M and O?
(ii) How many electrons are there in the outermost shell of M?
(iii) Name the group to which M belongs.
9. Water is a _____________[non polar/polar] covalent molecule in which
the atom of _____________[hydrogen/oxygen] attracts the electrons
strongly towards itself. The water molecule shows the presence of
_____________[double/single/triple] covalent bond/s and
 ____________ [one/two] lone pair of electrons present in the
____________ oxygen atom.
*********************************************THE END****************************************
 PERIODIC TABLE
1. State how do the following change in a period from left to right.
a. No. of shells
b. Valence electrons
c. Size of the atoms
d. Metallic character
2. Arrange the following as per instructions given in the brackets
a. He, Ar, Ne (Increasing order of electron shells)
b. Cl, Br, I (increasing electronegativity)
c. Na, K, Li (increasing electron affinity)
3. Name the following:
a. The element with the highest I.P
b. The largest size atom in the third period
c. A liquid non-metal
d. The noble gas with duplet
e. The family of elements to which chlorine belongs to
4. Match the atomic number 4,6,8,15 and 19 with each of the following
a. A solid non-metal of valency 3
b. A gas of valency 2
c. A metal of valency 1
d. A non-metal of valency 4
5. An element has atomic number 15. To which group and period of the
periodic table does this element belong.
6. Give reasons for the following:
a. Ionization potential increases across the period from left to right
b. Noble gases have zero electron affinity
c. The size of the cation is smaller than the size of the parent atom
 ELECTROLYSIS
I. Choose one correct answer to the questions from the given options:
1. Given below is an electrolytic cell diagram where copper, copper ions and
oxygen indicate the products formed on the respective electrodes. Choose the
correct diagram representing the electrolysis of copper sulphate solution using
active electrode:

a) U
b) T
c) V
d) W
2. The ionic equation for the chlorine on entering chemical reaction is:
a) Cl – e-  Cl+
b) Cl + e-  Cl-
c) Cl – 2e-  Cl2+
d) Cl + 2e-  Cl2-
3. The ions present in an electrolyte are Na+; H+; Pb2+; K+. Their order of
preferential discharge during electrolysis is:
a) Na+; H+; Pb2+; K+
b) H+; Pb2+; Na+; K+
c) Pb2+; H+; Na+; K+
d) K+; Na+; H+; Pb2+
4. During the electrolysis of acidified water which of the following takes place:
a) Oxygen is released at cathode.
b) Oxygen is released at anode.
c) Hydrogen is released at anode.
d) Sulphur dioxide is released at anode.
5. During the electrolysis of molten lead bromide, which of the following takes
places?
a) Bromine is released at the cathode
b) Lead is deposited at the anode
c) Bromine ions gain electrons
 d) Lead is deposited at the cathode
6. An aqueous electrolyte consists of the ions mentioned
below, which ion could be discharged most readily during
electrolysis?

a) Fe 2+
b) Cu2+
c) Pb 2+
d) H+
7. The diagram shows an electrolysis experiment.

During the electrolysis, hydrogen is formed at electrode P and oxygen is formed at

electrode Q. Which row correctly identifies P, Q and X?

P Q X
A anode cathode acidified water
B anode cathode dilute sodium chloride
C cathode anode acidified water
D cathode anode concentrated sodium chloride

II. Name the following:

1. The process by which an electrovalent compound breaks up into free mobile ions
in molten or aqueous form
2. The chemical compounds that contain only molecules
3. The process of formation cations and anions from the molecules of polar covalent
compounds on dissolving in water is
4. The product formed at the anode during the electrolysis of acidified water using
platinum electrodes
5. Electrode used as cathode in electrorefining of impure copper
6. The gas evolved at anode in the electrolysis of dilute sodium chloride solution
7. The reducing electrode in the electrolysis process
8. The process of deposition of a superior metal on the surface of a baser metal
 9. The oxidizing electrode in the electrolysis process
10. The gas evolved at anode in the electrolysis of concentrated sodium chloride
solution
III. Give reasons for the following:
1. Solid sodium chloride does not conduct electricity
2. The electrolysis of acidulated water is an example of catalysis
3. Conductivity of dilute hydrochloric acid is greater than acetic acid
4. The metals copper, silver and lead are electro refined but sodium is not electro
refined
5. In the electrolysis of the aqueous copper sulphate solution the blue colour does not
fades if copper electrode is used
6. Although copper is a good conductor of electricity it is a non- electrolyte.
7. Electrolysis is considered as a redox reaction.
8. In the electrolysis of dilute sodium chloride oxygen is liberated at the anode
however in the electrolysis of concentrated sodium chloride solution, chlorine gas
is liberated at the anode.
9. A bulb connected to a battery in the electrolysis of carbon tetrachloride does not
glow
10. In electrolysis of lead bromide, why is a graphite electrode used as the anode?
11. In the electrolysis of lead bromide why is the electrolytic cell in made up of silica?
12. Why dilute sulphuric acid is preferred to dilute nitric acid for acidification of
water?
13. Why sodium argentocyanide is preferred in the electroplating of an object with
silver?
14. Why should low current be passed for a longer period of time in the electrolysis
process?
15. Why does alternative current not be used in electrolysis process
IV. State one appropriate observation for the following:
1. At the anode, when molten lead bromide is electrolyzed using graphite
electrodes.
2. At the cathode when molten lead bromide is electrolyzed using graphite
electrodes.
3. At the anode when aqueous copper sulphate solution is electrolyzed using
copper electrodes.
4. At the cathode when aqueous copper sulphate solution is electrolyzed with
copper electrodes.
5. At the anode when an article is electroplated with nickel
V. Answer the following questions:
1. Mr. Ramu wants to electroplate his key with nickel to prevent from rusting. For
this electroplating.
(a) Name the electrolyte to be used.
(b) Name the cathode and anode.
(c) Write the reaction at anode and cathode.
(d) Give the observation at the anode and cathode
2. Observe the given diagram.
(a) Identify the electrode used in the given
diagram
(b) Write the equation at the anode and cathode.
(c) During the process will there be any change in
the anode. Give reason for your answer.

3. In the electroplating of an article with silver

(a) Name the electrolyte to be used.
(b) Name the cathode and anode.
(c) Write the reaction at anode and cathode.
(d) Give the observation at the anode and cathode
4. An aqueous electrolyte consists of the ions Fe2+, Cu2+, Pb2+, H+. Name the ion
which could be discharged most readily during electrolysis.
5. On electrolysis, Ag+ and H+ ions migrate to the _________ (cathode/anode) and
___________ (Ag+ /H+) are discharged.
6. As we descend electrochemical series containing cations, the tendency of cations
to get ________(oxidized/reduced) at the cathode increases
7. Electrolysis of aqueous sodium chloride solution will form ____________ at the
cathode [hydrogen gas/sodium metal]
8. Name the particles present in:
(a) Strong electrolyte
(b) Non- electrolyte
(c) Weak electrolyte
9. Write the reaction at anode and cathode for the following:
(a) Electrolysis of acidified water
(b) Electrolysis of aqueous copper sulphate solution using active electrodes
(c) Electrolysis of molten lead bromide
 CHEMISTRY METALLURGY WORKSHEET

I. Choose the correct answer:

1. The reason for using Aluminium in the alloy duralumin is:
a. Aluminium is brittle.
b. Aluminium gives strength.
c. Aluminium brings lightness.
d. Aluminium lowers melting point.
2. Which one of the following is the main ore of Aluminium?
a. Cryolite
b. Bauxite
c. Fluorspar
d. Corendum

3. The ratio of electrolyte taken in Hall Heroult’s process is 1:3:1. Identify

the electrolyte used:
a. fused alumina: cryolite: fluorspar
b. cryolite: fused alumina: fluorspar
c. fluorspar: fused alumina: cryolite
d. fused alumina: fluorspar: cryolite
4. Duralumin is used for:
a. Construction of aircrafts and ships
b. Resistance in heaters
c. Making coins
d. Soldering
5. A substance used in metallurgy to remove the rocky impurities is called:
a. Slag
b. Flux
c. Oxidizing agent
d. Gangue
6. The percentage composition of magnalium is:
a. Al = 75%, Mg = 25%
 b. Al = 80%, Mn = 10%, Mg=10%
c. Cu = 4%, Al = 95%, Mg = 0.5%, Mn = 0.5%
d. Al = 95%, Mg = 5%
7. Which of the following reactions occurs at the anode during the
electrolysis of aluminum?
a. 2O–2+4e-  O2
b. 2O2– O2+4e–
c. Al3+ Al + 3 e –
d. O2 2O2–+4e–
8. Duralumin is an alloy of:
a. Al & Cu
b. Cu & Sn
c. Al & Ag
d. Al & Fe
9. The metals zinc & tin are present in alloy:
a. Solder
b. Brass
c. Bronze
d. Duralumin
10. The main ore used in the extraction of iron is:
a. Haematite
b. Calamine
c. Bauxite
d. Cryolite
II. Name the following.
1. The process of concentration of bauxite using sodium hydroxide
2. The process by which sulphide ore is concentrated
3. The chemical formula for bauxite
4. The naturally occurring minerals from which metals can be extracted
profitably.
5. The crystals added in the hydrolysis of sodium aluminate
6. Name the cathode & the anode in Hall Heroult’s process
 7. The purity of aluminium obtained by this method
8. The method used to separate ore from gangue by preferential wetting
9. The compound added to lower the fusion temperature of electrolytic
bath in the extraction of aluminium
10.The ore of zinc containing its sulphide
11.A metal present in cryolite other than sodium
12.A black powdery substance used for the reduction of zinc oxide to
zinc
13.The substance is an alloy of zinc, copper and tin
14. Electrolytic deposition of a superior metal on a base metal
15.The process of heating the concentrated ore to a high temperature in
the presence of excess air
16. Separation of ore and gangue due to difference in the density of
particles
17.Process of heating the concentrated ore in limited supply of ore or
absence of air
18. The process of crushing the ore into a fine powder is called

III. Answer the following:

A. Answer the following questions related to the extraction of aluminium from

its ore:

1. Name the above process of extraction of aluminium.

2. Name the ores of aluminium and its chemical formula.
3. Write the steps and the equations in the Bayers process
4. Why is aluminium reduced by electrolytically and not using
reducing agents?
5. State the difficulties in the electrolysis of fused alumina. How was
that overcome?
 6. Give reason for the addition of cryolite and fluorspar to the
electrolytic reduction of alumina.
7. During the process of electrolytic reduction of alumina, a layer of
powdered coke is added. Give reason
8. The graphite anode is periodically replaced. Give reason
9. Write the equation at the anode & cathode
10.Name the products formed at the anode & cathode
11.State the reason for addition of caustic alkali to bauxite ore during
purification of bauxite. Give a balanced equation for the above
reaction.
12.Graphite is used as the electrode. Why?
B. Answer the following questions:

1. Define: metallurgy, flux, matrix & alloy.

2. Carbonate and sulphide ores are usually converted into oxides
during the process of extraction. Give reason

**************************THE END*****************************
 HYDROCHLORIC ACID

a. Name the following:

1. Drying agent for hydrogen chloride

2. An acidic gas which gives dense white fumes with NH3

3. The gas obtained when rock salt reacts with Concentrated sulphuric
acid.

4. The experiment which demonstrates extreme solubility of hydrogen

chloride gas

5. Name the arrangement used for the prevention of back suction of

water during the dissolution of hydrogen chloride gas in water
b. Choose the right option:

1. Dry hydrogen chloride gas can be collected by

__________________ (upward, downward) displacement of air

2. Quick lime is not used to dry HCl because

__________________(CaO is alkaline, CaO is acidic, CaO is
neutral)

3. Hydrogen chloride can be obtained by adding concentrated

sulphuric acid to _________________ (NaCl, Na2SO4, Na2CO3,
NaNO3)

4. The aim of the fountain experiment is to prove ______________

(HCl turns blue litmus red, HCl is denser than air, HCl is soluble in
water, HCl fumes in moist air)

5. Hydrogen chloride gas being highly soluble in water is dried by

_________________ (anhydrous calcium chloride, phosphorous
pentoxide, quicklime, concentrated sulphuric acid)

c. Fill in the banks with the correct answer:

1. Dry hydrogen chloride gas can be collected by

__________________ displacement of air
 2. The aim of the fountain experiment is to prove ______________

3. Hydrogen chloride gas being highly soluble in water is dried by

_________________

d. Name the gas evolved in each of the following

1. Dilute hydrochloric acid reacts with copper (II)carbonate
2. Dilute hydrochloric acid reacts with magnesium sulphite
3. Dilute hydrochloric acid reacts with iron (II)sulphide
4. Dilute hydrochloric acid reacts with Magnesium
5. The gas produced when magnesium sulphite reacts with HCl
e. Write the chemical equation for the following:
1. Dilute hydrochloric acid reacts with copper (II)carbonate
2. Dilute hydrochloric acid reacts with magnesium sulphite
3. Dilute hydrochloric acid reacts with iron (II)sulphide
4. Dilute hydrochloric acid reacts with Magnesium
5. Dilute hydrochloric acid reacts with Iron (II) oxide
6. Dilute hydrochloric acid reacts with Iron (III)oxide
7. Silver nitrate with dilute hydrochloric acid
8. Dilute hydrochloric acid reacts with lead (II) nitrate
f. State one observation for each of the following:
1. Ammonia is reacted with hydrochloric acid
2. Action of dilute hydrochloric acid on iron(II)sulphide
3. Hydrogen chloride gas is passed through silver nitrate solution
4. Hydrogen chloride gas is passed through lead nitrate solution and the
product thus formed is heated
5. Reaction of concentrated hydrochloric acid with manganese dioxide
6. Reaction of silver nitrate solution with hydrochloric acid
7. Dilute hydrochloric acid reacts with copper oxide
8. Dilute hydrochloric acid reacts with copper (II) hydroxide
9. Excess of ammonium hydroxide is added to a substance by adding
hydrochloric acid to silver nitrate solution
g. Answer the following:
1. Laboratory method of preparation of HCl
(i) Write the equation for the reaction.
(ii) How would you check whether the gas jar is filled with hydrogen
chloride?
(iii) What does the method of collection tell you about the
density of hydrogen chloride?
(iv) Name the chemical used for drying the above gas.
(v) While dissolving the above gas in water, what is the arrangement
done?
(vi) Why do we use such an arrangement?
(vii) Why is Con. H2SO4 used instead of Con.HNO3?
(viii) What would happen if quicklime is used as a drying agent?
(ix) What is the method of collection of dry HCl gas. Give reason.
(x) State a safety precaution you would take during the preparation of
hydrochloric acid.
(xi) Calcium oxide and phosphorous pentoxide are good drying agents but
they are not employed to dry hydrogen chloride gas. Give reasons for
each.
(xii) Give a chemical test to distinguish between hydrogen chloride gas
and hydrogen sulphide gas
h. Give reasons for the following:
1. Direct absorption of HCl in water is not preferred
2. Hydrogen chloride gas cannot be dried over quicklime
3. Hydrogen chloride is not collected over water
4. Hydrogen chloride gas is collected by the upward displacement of air
5. Dry HCl does not change the colour of blue litmus
6. HCl gas has fumes in moist air
i. Answer the following questions:
1. The solution in the flask is blue litmus and water
a. Identify the gas Y
b. What property of gas does this experiment demonstrate.
c. Name another gas which has the same property and can be
demonstrated through this experiment.
d. Identify the colour of the water that has entered the round
bottomed flask
2. The diagram shows an apparatus for the laboratory preparation of
hydrogen chloride.

a) Identify A and B.
b) Write the equation for the reaction.
c) How would you check whether or not the gas jar ls filled
with hydrogen chloride?
d) What does the method of collection tell you about the
density of hydrogen chloride?
 Class 10 Chemistry Revision – useful equations.
Chapter 1 - Periodic Table.
Property IA IIA ....→....’across a period’.. → . .. VIIA Group 0
Atomic Size Large → → → → → → → → → small large
Valency with Oxygen 1 2 3 4 5 6 7 0
Valency with Hydrogen 1 2 3 4 3 2 1 0
Electronegativity, EA, IP. Low → → → → → → → → High Electro-0. EA- 0. IP – v. high
Oxides Basic → → → → → → → Acidic No Oxides
Metallic Nature Metallic → Semi-metals → Non-metals
Definitions. Electronegativity – ‘Tendency to attract electrons in a compound’.
EA (Electron Affinity) – ‘Energy released when an electron is added’
IP (Ionisation Potential) – ‘Energy taken to pull off an electron’
Q. Why does Atomic Size decrease across a period?
A. Because the charge on the Nucleus increases but the number of electron shells remains the same.
Going down a Group – Increases: Atomic Size, metallic nature. Decreases: Electro, EA, IP.
Valency remains the same.
Chapter 2 – Chemical Bonding -
Covalent Bond – ‘Two atoms sharing a pair of electrons, one electron comes from each
atom’. Eg. Molecules like Methane, organic compounds, O2, N2,
Polar covalent compounds have small + and – charges at each end. Eg. Water, Ammonia.
Coordinate Bond- ‘Two atoms sharing a pair of electrons which is a ‘lone pair’ donated by
one atom’. Eg H3O+ (Hydronium Ion), NH4+ (Ammonium Ion)
Formation of Ionic bond (aka- Electrovalent Bond)
Na (2,8,1) ─ e─ → Na+ (2,8) and Cl (2,8,7) + e─ → Cl─ (2,8,8)
Na+ and Cl─ stick together by Electrostatic Attraction.
Weak organic acids (eg. Acetic Acid) contain both molecules and ions.
Chapter 3 – Acids, Bases and Salts.
Acid Reactions. (1) Acid + Base → Salt + Water (called ‘Neutralisation’ or ‘Titration’)
H2SO4 + CuO → CuSO4 + H2O
(2) Acid + Carbonate → Salt + Water + CO2
2HCl + ZnCO3 → ZnCl2 + H2O + CO2
(3) Acid + Sulphite → Salt + Water + SO2
HNO3 + CaSO3 → Ca(NO3)2 + H2O + SO2.
(4) Acid + Reactive Metal → Salt + Hydrogen gas
H2SO4 + Zn → ZnSO4 + H2.
Note: Nitric Acid (conc. or dil.) does not give Hydrogen with any metals.
(5) Acid + Sulphide → Salt +H2S. eg. 2HCl + FeS → FeCl2 + H2S
Making Salts – (1) Methods for making Soluble Salts.
(a) Neutralisation reaction – Acid + Base (can also use Acid + Carbonate)
(aka.’ Titration’) HCl + NaOH → NaCl + H2O
(b) Acid + Reactive Metal (Fe, Mg, Zn) H2SO4 + Zn → ZnSO4 + H2
(c) Direct Combination. 2Fe + 3Cl2 → 2FeCl3.
Note: Iron + Chlorine (Direct Combination) produces Iron (III) Chloride FeCl3 but:
Iron + Hydrochloric Acid produces Iron (II) Chloride FeCl2
 (2) Method for making Insoluble Salts. Double Displacement (Precipitation)
Soluble Salt of Metal (Nitrate) + Soluble Salt (Sodium or Acid) of the Anion.
Eg. Making Lead Sulphate: Pb(NO3)2 + H2SO4 → PbSO4↓ + 2 HNO3
Note: if starting with an insoluble salt eg. PbCO3 first dissolve it in Nitric Acid.
Types of Salts.
(1) Normal Salt (One Metal Cation + Anion) eg. NaCl, CuSO4, Ca(NO3)2.
(2) Acid Salt (One Metal Cation + Hydrogen + Anion) eg. NaHCO3
(Sodium Hydrogen Carbonate) Note: Name contains ‘Hydrogen’
(3) Mixed Salt (Two metals with one Anion) eg. NaKCO3 (Sodium Potassium Carbonate)
(4) Double Salt (two Simple Salts mixed) eg. Alum
(5) Complex Salt (contains metal in the Anion) eg. NaAlO2 (Sodium Aluminate)
OR: The Cation has non-metal groups attached eg. [Cu(NH3)4]SO4 (called Tetrammine
Copper Sulphate – a dark blue solution made by adding excess Ammonia (NH4OH) to Copper
Sulphate.
Salt Analysis.
Metal Ions Few drops NaOH Excess NaOH Few drops NH4OH Excess NH4OH
Na+, K+, NH4+, none none none none
Ca2+ No ppt. White ppt. No ppt. No ppt.
Zn2+ White ppt dissolves White ppt. dissolves
Al3+, Pb2+ White ppt. dissolves White ppt. Insoluble
Cu2+ Blue ppt. Insoluble Blue ppt. Dark blue soln.
Fe2+Green, Fe3+brown Coloured ppt Insoluble Coloured ppt Insoluble
Amphoteric Metals dissolve in both acids and alkalis eg. Zinc, Lead and Aluminium. They give white
ppt . with few drops of NaOH which dissolves in excess to produce a complex salt.
Only Zinc does the same thing with NH4OH.
Zn(OH)2 + 2NaOH → Na2ZnO2 + 2H2O (Same for Pb instead of Zn)
Al(OH)3 + NaOH → NaAlO2 + 2H2O
Barium Chloride gives a white ppt with Sulphates. BaCl2 + CuSO4 → BaSO4↓ + CuCl2
Silver Nitrate gives a white ppt with Chlorides which dissolves in xs Ammonia
AgNO3 + NaCl → AgCl↓ + NaNO3 and: AgCl + 2NH3→[Ag(NH3)2]Cl (Di-ammine Silver Chloride)
Nitrates give the ‘Brown Ring Test’ producing ‘Nitroso Ferric Sulphate’.
Identification of gases. (i) Acidic Gases – turn blue Litmus red – CO2, SO2, Cl2, H2S, HCl, NO2
CO2 – no smell, turns lime water milky and then clear with excess.
SO2 – sharp smell, turns lime water milky and clear with excess, turns orange K2Cr2O7 green.
Cl2 – sharp smell of swimming pools, bleaches litmus, turns Starch Iodide paper blue/black
H2S – smell of rotten eggs, turns Lead Acetate paper black.
HCl – sharp smell, dense white fumes with Ammonia. NH3 + HCl → NH4Cl (Ammonium Chloride)
NO2 – brown gas, sharp smell, turns Starch Iodide blue/black
(ii) Neutral gases (no smell, no change in Litmus) – O2, H2,
O2 – relights glowing stick.
H2 – light gas, collect a tube-full, burns with a ‘pop’.
(iii) Alkaline gas – turns red Litmus blue. NH3 – dense white fumes with conc. HCl.
5A. Mole Concept.
1 mole = 6 x 1023 atoms/molecules of a substance (Avagadro’s Number)
1 mole = Molecular Weight of the substance in grams. eg. 1 mole of H2O = 2 + 16 = 18 grams.
1 mole of any gas occupies 22.4 litres at STP (‘Standard Temperature and Pressure’ = 0°C and 1 atm.)
Avagadro’s Law. Equal Volumes of gases contain equal numbers of molecules. (at same temp.
and pressure)
Vapour Density of a gas = ½ Molecular Mass.
In the reaction: Mg + 2HCl → MgCl2 + H2 (Molecular Weights: Mg=24; Cl=35.5; H=1)
One mole of Mg reacts with 2 moles of HCl to produce one mole of MgCl2 and one mole of H2.
ie. 24 grams of Mg reacts with 2 x 36.5 = 73 grams of HCl to make 95 g of MgCl2 and 2 g of H2.
5B. Stoichiometry and Empirical Formula. 2 types of problems.
(1) Formula of a compound given. Find the percentage of each element.
Percentage of an element = Number of atoms of the element x Atomic Weight x 100
Total Molecular Weight of the compound
(2) Percentage of each element in a compound given. Find the formula.
Method: Divide each percentage by the Atomic Weight of that element. Then find the
simplest ratio between the answers. This gives the ‘Empirical Formula’. The actual
formula may be a multiple of this.
Balancing Equations: Do metals first, non-metals next and Oxygen and Hydrogen last.
6. Electrolysis.
An Electrolyte conducts electricity accompanied by a chemical change when dissolved in
water or in the fused state (melted). Electrolytes are Salts, Acids and Alkalis.
Weak electrolytes are weak acids eg Acetic Acid and weak bases eg Ammonia.
Non-electrolytes - Organic compounds eg sugar, urea. Pure water does not conduct electricity.
Metals form + Cations which go to Cathode (- electrode) where Cations become metals again.
Anode reactions: Either the metal of the Anode dissolves or Oxygen is liberated.
Electrolysis of Copper Sulphate using Copper metal as the Anode.
Cathode reaction: Cu2+ + 2e– → Cu (metal). Anode reaction: Cu – 2e– → Cu2+
Electrolysis of water. Use Platinum electrodes. H2SO4 is added so water conducts electricity.
Cathode Reaction for water: 2H+ + 2e– → H2 (gas) Anode Reaction: 4OH– – 4 e– → O2 + H2O.
Fused salts eg Lead Bromide PbBr2: Cathode: Pb2+–2e– → Pb (metal) Anode: 2Br––2e– → Br2 (gas)
7. Metallurgy.
Making Iron from Iron Ore Fe2O3 (Haematite).
Mixture put in Blast Furnace – Iron Ore + Coke (Carbon) + Limestone (Calcium Carbonate)
Coke burns to form CO - Carbon Monoxide, which is the reducing agent.
Reduction: Fe2O3 + 3CO → 2Fe (metal) + 3 CO2.
Calcium Carbonate CaCO3 decomposes to CaO + CO2.
Removing impurities: CaO + SiO2 → CaSiO3 (Calcium Silicate –‘Slag’)
 Making Aluminium from Bauxite – Al2O3 (aka. Alumina).
Purification of Aluminium Ore, Al2O3 (Bauxite) – Baeyer’s Process.
Step 1: Al2O3 dissolved in NaOH. Al2O3 + 2NaOH → 2NaAlO2 + H2O.
Step 2: Solution is diluted to precipitate Al(OH)3 which is then heated to 1000° to reform
pure Al2O3. 2Al(OH)3 → Al2O3 + 3H2O
Hall’s Process is similar but Al2O3 is reacted with Na2CO3 to form NaAlO2 + CO2 and
CO2 is used to precipitate the Al(OH)3. (Impurities are SiO2 and Fe2O3)
Extraction: Alumina is dissolved in Cryolite Na3AlF6 at 950°C and electrolysed using
Graphite (Carbon) rods as Anode. The rods need replacing often as they react with the
Oxygen formed. Anode reaction: O2─-2e─ → [O] and C + 2[O] → CO2
Cathode reaction: Al3+ + 3e─ → 3Al (metal)
(Note: metals are always formed at the Cathode) Liquid Aluminium (mp. 660°C) is tapped off.
Making Zinc from Zinc Blende ZnS.
Zinc Blende is purified by Froth Floatation – only works for Sulphides.
ZnS is ‘roasted’ in air to produce Zinc Oxide and SO2. 2ZnS + 3O2 → 2ZnO + 2SO2
Zinc Oxide reduced with carbon. 2ZnO + C → 2Zn (metal) + CO2
Zinc vapours are condensed with molten Lead.
8. Sulphuric, Nitric and Hydrochloric Acid
Copper with Nitric Acid. Copper does not normally react with acids but Nitric Acid is a very
strong Oxidising Agent.
(1) Conc. Nitric. Cu + 4HNO3 → Cu(NO3)2 + 2NO2 + 2H2O
(2) Dilute Nitric. 3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O
Conc. Nitric and Sulphuric Acids both oxidise Carbon and Sulphur to oxides:
C + 4HNO3 → CO2 + 4NO2 + 2H2O and C + 2H2SO4 → CO2 + 2SO2 + 2H2O
S + 4HNO3 → SO2 + 4NO2 + 2H2O and S + 2H2SO4 → 3SO2 + 2H2O
Making HCl and HNO3 with conc. Sulphuric Acid.
2NaCl + H2SO4 → Na2SO4 + 2HCl
2KNO3 + H2SO4 → K2SO4 + 2HNO3 (for Nitric Acid - Apparatus must be all glass)
Sulphuric Acid is used because it is a non-volatile acid (High Boiling Point = 198°C)
Conc. H2SO4 + Sugar (or paper) turns black as conc H2SO4 is a strong dehydrating agent-
C6H12O6 (+ conc. H2SO4) → 6C (black) + 6H2O
Aqua Regia is 3 parts conc. HCl + 1 part conc. HNO3 (dissolves gold)
HCl is a very soluble gas – does Fountain experiment. Funnel arrangement for dissolving.
Heating Nitrates – all Nitrates give off Oxygen and NO2 gas when heated.
2Pb(NO3)2 → 2PbO + 4NO2 + O2 (PbO is yellow when cold and brown hot)
The exceptions are Sodium and Potassium Nitrates which give only oxygen....
Eg. Heating Potassium Nitrate 2KNO3 → 2KNO2 + O2
… and Ammonium Nitrate NH4NO3 +heat→ N2O (Nitrous Oxide) + 2H2O.
Making Sulphuric Acid – Contact Process. Catalytic oxidation of SO2.
2SO2 + O2 → 2SO3. Catalyst - Vanadium Pentoxide V2O5 at 500°C
Step 2: SO3 + H2SO4 → H2S2O7 (Oleum) Step 3: H2S2O7 + H2O → 2 H2SO4. (Dilution of Oleum)
8B. Making Ammonia – Haber’s Process (industrial process).
N2 + 3H2 → 2NH3 – Catalyst- Iron with traces of Molybdenum at High Pressure (500atm)
and 500°C. Collect Ammonia by liquefaction at pressure.
Also making Ammonia: Magnesium or Aluminium Nitride with water
Mg3N2 (or AlN) + 6H2O → 2NH3 + 3Mg(OH)2 (or Al(OH)3)
Properties of Ammonia – Very soluble in water - does fountain experiment
Basic – turns red litmus blue.
Lighter than air – Collected by ‘downward displacement’ of air.
Making Nitric Acid – Ostwalt’s Process. – *Catalytic Oxidation of Ammonia. 4NH3 +
5O2 → 4NO + 6H2O – Catalyst – Platinum at 800°.
Step 2: Spontaneous Oxidation of NO to NO2. 2NO + O2 → 2NO2
Step 3: Spontaneous Oxidation of: 4NO2 + 2H2O + O2 → 4HNO3
*Note: Burning Ammonia in pure Oxygen gives different products:-
4NH3 + 3O2 → 2N2 + 6H2O - burns with a green/ yellow flame.
9. Organic Chemistry.
Saturated Hydrocarbons – all single bonds – name ends in –ane. Eg Ethane, Butane.
Unsaturated Hydrocarbons – Double bond – name ends in –ene. Eg. Ethene, Butene
Triple bond – name ends in –yne eg. Ethyne, Butyne.
(1) Halogenating Hydrocarbons (with Chlorine or Bromine).
(a) Substitution reactions: CH4 + Cl2 → CH3Cl (Chloromethane)+ HCl
(b) Addition reactions: C2H4 + Cl2 → C2H4Cl2 (Dichloro-ethane)
(2) Making Alcohols from Chloro (or Bromo) Alkanes- Treat with *aqueous NaOH or KOH
C2H5Cl + NaOH → C2H5OH (Ethanol) + NaCl
*Alcoholic NaOH or KOH removes HCl from C2H5Cl and gives C2H4 + KCl + H2O
(3) Oxidation of alcohols – (a) Mild oxidation (pass over hot Copper Oxide)
C2H5OH + CuO → CH3CHO (Ethanal) + Cu + H2O – makes aldehydes
(b) Strong Oxidation – (Acidified Potassium Dichromate K2Cr2O7 or KMnO4)
C2H5OH + O2 → CH3COOH (Acetic Acid) + H2O – makes Carboxylic Acids
(4) Dehydration of alcohols – using conc. H2SO4 or pass over hot Alumina.
C2H5OH → C2H4 (Ethene) + H2O - produces unsaturated Hydrocarbons.
(5) Decarboxylation of Acids. – make the Sodium salt of acid (eg. Propionic Acid)
and then treat with ‘Soda Lime’ (chemically this is just NaOH).
Step 1: C2H5COOH + NaOH → C2H5COONa (Sodium Propionate) + H2O
Step 2: C2H5COONa (Sodium Propionate) + NaOH → C2H6(Ethane) + Na2CO3
- produces a saturated Hydrocarbon with one less C-atom than original Acid.
(6) Esterification - Carboxylic Acid + Alcohol → Ester + Water – conc. H2SO4 must be added to
pull off the water. CH3COOH + C2H5OH → CH3COOC2H5 + H2O
Acetic Acid Ethanol Ethyl Acetate
Other important reactions:
o CaC2 (Calcium Carbide) + 2H2O → C2H2 (Acetylene (Ethyne)) + Ca(OH)2
o Addition of water (steam) to Ethene: C2H4 + H2O → C2H5OH
o Combustion: Hydrocarbons burn to form Carbon Dioxide and water.
Eg. Combustion of Methane: CH4 + 2O2 → CO2 + 2H2O