LAB PREPARATION OF OXYGEN FROM POTASSIUM CHLORATE AND MANGANESE IV OXIDE

my free school tannzania January 18, 2018 Materials
LAB PREPARATION OF OXYGEN FROM POTASSIUM CHLORATE AND MANGANESE IV OXIDE
Apparatus: Hard boiling tube, delivery tubes, gas jar, stopper, retort stand, beehive shelf, water trough.
Chemicals: Manganese IV Oxide, Potassium Chl.
Method: A known amount of potassium chlorate and a small amount of manganese IV oxide is put in a hard boiling tube.
The apparatus is set as shown in the diagram above.
The mixture is heated.
The manganese IV oxide acts as a catalyst.
Observation
Cracking sound occurs. A colourless gas O2.is given off. The gas is collected by downward displacement of water.
Note: It is possible to collect the gas by passing it through water because it is slightly soluble in    water.
What is the importance of oxygen being soluble in water?
Aquatic organisms i.e. fish can obtain the oxygen for respiration or breathing.

Physical properties of O2.
It is colourless, tasteless, odourless.

Chemical Properties
Test for oxygen
– Relights a glowing splint.
O2reacts and turns colourless Nitrogen Monoxide gas into reddish brown Nitrogen Dioxide gas.
Nitrogen Monoxide + Oxygen = Nitrogen Dioxide
Chemical Properties of Oxygen Gas
Metals and non-metals burn in O2to form metal oxides and non-metallic oxides respectively.
Metals with Oxygen
When burning, Mg is introduced in a gas jar of oxygen. It continues to burn with bright sparkling flame and white solid magnesium oxide remains as a residue.
Oxygen with non-metals
Oxygen with sulphur
Burning sulphur on a deflagrating spoon is put in a gas jar of oxygen.
Observation.
Sulphur burns with bright blue flame producing white fumes of sulphur dioxide. The fumes dissolve in water forming a colourless acidic solution.
Sulphur Dioxide + Water = Sulphuric Acid

Charcoal with Oxygen
Red hot carbon on a deflagrating spoon is introduced in a gas jar of oxygen. The charcoal burns with a lot of sparks giving off colourless gas.
Carbon + Oxygen = Carbon Dioxide
Carbon Dioxide + Water = Carbonic Acid

Phosphorus with oxygen
Red phosphorus burns in oxygen with a bright yellow flame producing white fumes.

Uses of Oxygen
– Used in LD process to manufacture steel.
LAB PREPARATION OF HYDROGEN
From Magnesium and dilute Hydrochloric acid.
Observation
Rapid effervescent occurs and colourless gas is given off.
Magnesium + Hydrochloric Acid = Magnesium Chloride + Hydrogen
The gas is collected by upward delivery or downward displacement of air.
All the gases which are less dense than air are collected by using the above method.
Heavier gases than air are collected by downward delivery or upper displacement of air. E.g. Carbondioxide, Hydrogen sulphate etc.

Chemical Properties of H2
Test for Hydrogen
If a burning splint is held on a mouth of a gas jar containing hydrogen, it burns with a pop sound.
Hydrogen + Oxygen + Water
Hydrogen is a reducing agent
CopperOxide + Hydrogen = Copper + Water

Reducing agent removes oxygen from substances and adds hydrogen. Reduction is addition of hydrogen and removal of oxygen.

2.FUELS AND ENERGY
Fuel is any substance which burns in air (O2) to produce energy in form of heat. The heat produces is used economically for domestic and industrial purposes. The main types of fuels used in Tanzania are coal, charcoal, petrol, diesel, kerosene from crude oil as fossil fuels.
How would you obtain fuel? E.g. Charcoal from the locally available materials.
Charcoal is made by the dry distillation of wood at a temperature of about 4000° – 4500°C in a pit kiln or earth mould kiln.
In a pit kiln, the wood is heaped in a hemispherical pile in a central pit. The wood is covered with soil/mud leaving a small air hole near the bottom. The wood is lit at the bottom and allowed to burn until the whole pile is on fire, producing CO2, water and volatile organic compounds which escape into the atmosphere. The holes for allowing air are then closed. The pit is kept closed till the fire goes off and the charcoal cools. The charcoal is about 20% by weight and 75% by volume of the wood.
In the earth kiln, the wood is heaped in a pile above the surface instead of a pit. Good charcoal is porous brittle and retains the form of the wood. It burns with a non-luminous flame.

Types Of Fuels
Type
Example
Comment
1
Solid
Wood, Coal, Charcoal, Coke.
Coal and coke are used mostly in industries. Coke is residue left from the destructive distillation of coal.
Wood and Charcoal are obtained from plants. Coal is a fossil and other organisms that lived many million years ago.
When solid fuels are burnt, they leave solid ashes behind.
2
Liquid
Petrol, Diesel, Kerosene, Ethanol (alcohol)
Liquid fuels have an advantage over solid fuels because they leave no solid residue when burnt. They can be regulated by automatic devices and are relatively more convenient to handle.
3
Gaseous
Biogas/ natural gas, Coal gas and producer gas are mostly used in industries. They are obtained from coke and steam and/or coke and air.
They are easier to handle then liquids and solids.
Bottled gas delivered to our homes is liquefied propane or butane or a mixture of their two.

Characteristics of a good fuel
All fuels give out energy in form of heat but their efficiency and quality differ. A good fuel should have most of the following characteristics:-
(a) Should have a high heat content i.e. must burn easily and produce a lot of energy.
(b) It must be cheap.
(c) It should have a little or no waste products like ash and smoke.
(d) It must not give off dangerous byproducts i.e. poisonous fumes.
(e) It should be easily controlled.
(f) It should be easily stored and transported.

Efficiency of a fuel
A good fuel burns easily to produce a large amount of energy.
Calorific Value
The amount of heat given out when 1gm or 1kg of fuel burns completely in air.
Energy Value
Is the amount of heat given out when one mole burns completely in air.
Experiment 1: To determine energy value of kerosene.
Procedure
– Put a tin on a burner.
– Put 100cm3 of water in the tin and put the thermometer in position.
– Record the temperature of the water when it’s steady.
– Half fill the burner with kerosene.
– Weigh the burner and its contents.
– Light the burner and immediately put the tin on top of the burner.
– Stir the water using thermometer until the temperature is 15°C – 20°C.
– Extinguish the flame and not the final temperature.
– Re-weigh the burner.

From the experiment
(a) Mass of burner and kerosene (Initial)
(b) Mass of burner and kerosene (Final)
(c) Mass of kerosene burnt
(d) Final temperature of the water
(e) Initial temperature of the water
(f) Rise in temperature
(g) Mass of water

Heat given out = Mass of water X specific heat of water X rise in temperature

Uses of Fuel In Our Daily Life
Transport                     Aeroplanes                              Kerosene
Vehicles                               Petrol, Diesel, Gas
Ships                                                Diesel
Train                                     Coal

Industries                       Power station
Heavy duty + Light duty

Domestic Uses                        Cooking                      Gas
Kerosene
Charcoal

Environmental Effect of Using Charcoal and Firewood As Sources of Fuel
Charcoal and firewood come from trees. Use of these as a source of fuel leads to:-
Soil erosion
Cutting down the trees leaves the land bare and exposed to agents of soil erosion.
Deforestation
Forests are habits for organisms therefore cutting down trees destroy habitats of organisms i.e. leopards, monkeys to other places.
Desertification
Trees are a catchment for rains, cutting them down disrupts the rain cycles and thus decrease in rainfall. Continuous cutting down of trees without planting may lead to a development of desert.
Global warming
Trees consume carbondioxide in the process of photosynthesis. Their removal leads to accumulation of carbondioxide in the atmosphere, which causes global warming. Some effects of global warming are:-
(i)         Melting of ice belts.
(ii)        Rising of sea level.
(iii)       Extreme climate: – drought, floods.
(iv)Spread of diseases.
Burning charcoal gives a lot of smoke, which pollutes the air. Also the smoke solid particles block the stomata of plants thus decrease diffusion of gases in and out of plants. This leads to low productivity.
Air pollution
Incomplete combustion of charcoal and firewood produces carbonmonoxide, which causes respiratory problems.
– The carbondioxide given out contributes to global warming.
– The carbondioxide causes acid rain i.e. carbondioxide dissolves in the rainwater forming carbonic acid.
Water pollution
The solid residue may cause water pollution if washed in water bodies e.g. rivers.

Contribution of Vegetation to Balance of Atmospheric Gases
Air is a mixture of gases i.e. N2, CO2, O2 and noble gases. The CO2 and O2 enter the plants through stomata. During day time, plants consume CO2 in the process of photosynthesis and give out O2. In the dark (night), plants consume O2 and give out CO2 in the process of respiration. In photosynthesis and respiration process, the plants bring the level of gases to the required amount in the air.
Some plants can convert the atmosphere N2 into nitrates and nitrates are used by plants to make protein.

Alternative to firewood and charcoal as a source of fuel.
Sources of fuel can be divided into renewable and non-renewable sources.
Renewable sources are those which are continuously being replaced within a short period of time. These include wind energy, solar energy, firewood and charcoal.
Non-renewable sources are those which cannot be replaced within a short period of time. These include fossil fuels i.e. oils, natural gas, nuclear and coal.
Most of the energy used today comes from non-renewable sources.
The alternative to firewood and charcoal would be solar energy and fossil fuels such as coal, natural gas. This type of energy cannot be exhausted. It is also clean as it does not release harmful gases.

Solar energy can be tapped in many ways like:-
– Generating electricity.
– Heating and cooking using parabolic mirrors.
– Heating and cooling use solar chimneys.
– Geothermal energy obtained by tapping the heat in the earth’s crust.

Wind energy
Wind is moving air. The energy is usually hammered by wind mills. This energy causes no pollution.

Water power
H20 possesses energy in form of kinetic energy. The forms of water energy include: Hydroelectric energy and Tidal energy.



3.CONSERVATION OF ENERGY
Energy – Is the ability/capacity of a body to do work.
– S1 unit of energy is Joules (J).
– Energy exists in two forms: –
(1) Potential Energy – Is the energy in matter due to its position or state.
(2) Kinetic Energy – Is the energy possessed by a body due to motion. The motion could be waves, electrons, atoms, molecules or the object itself.
Mechanical energy is the sum of kinetic and potential energy.

Some forms of potential and kinetic energy
Potential Energy
Kinetic Energy
1
Chemical energy is possessed by matter due to its chemical makeup i.e. arrangement of atoms and molecules.
Electrical energy is possessed by electrical charges in motion. E.g. Electricity and lighting.
2
Elastic energy is stored in objects by application of force. E.g. Compressed springs, rubber bands.
Radiant energy is electromagnetic energy that travels in transverse waves. E.g. Solar energy, visible light, x-rays, radio waves.
3
Nuclear energy is possessed by an atom in its nucleus. Nuclear energy holds the nucleus together. Energy is released when the nucleus are combined or split open.
Thermal (heat) is the internal energy in substances caused by vibration of atoms within the substance.
4
Gravitational energy is possessed by body due to place. E.g. When an object is lifted, it possesses gravitational energy.
Sound energy is the movement of energy through substances in longitudinal waves. Sound is produced when force causes a substance to vibrate.

From the above, it shows that energy is conserved. Principle/law of conservation of energy states that “Energy can neither be created nor destroyed, it can only be change from one form to another”.

Transformation of Energy – is the process of changing energy from one form to another.
Biogas
Is a fuel gas derived from decomposing biological waste. It can be easily produced from both industrial and domestic waste such as paper production and sugar production waste, sewage and animal waste. The waste matter is put together and allowed to ferment naturally, producing biogas. This can be done by converting the existing waste disposal channels into biogas plants. When all the methane has been extracted by a plant, the remains can be used as fertilizer.
Use of biogas in environmental conservation.
It is an efficient fuel that burns completely to produce a large amount of heat energy leaving no solid waste products like ashes which when washed to water bodies causes pollution.
After extraction of all the biogas, the remaining byproducts can be used as fertilizers thus enriching the soil with nutrients, which will be absorbed by the plants. Biogas raw materials are waste products from sugar, paper industries, sewage and animal waste. These, if left in the environment would accumulate and cause terrestrial pollution.
Biogas is a useful alternative fuel for firewood and charcoal thus spares the trees and vegetation, which are useful in preventing soil erosion.

Geothermal energy
2 Greek words
Geo=Earth
Therme=Heat
The energy is obtained by tapping the heat in the earth’s crust. The temperature at the core is very high. This heat sometimes finds its way to earth’s surface in form of volcanoes, hot springs and geysers. This heat can be directly used for heating, cooking and bathing.

Wind energy
Wind is moving air. It’s usually harnessed using windmills. The wind turns the blades of the windmills, which runs turbines and produces energy. Areas where winds are high like high altitude sites are preferred locations.

4.ATOMIC STRUCTURE
The Atom
An atom is the smallest particle of an element that has all the chemical properties of the element.
The Atomic Theory
In 1803, Dalton developed his theory about the atom. The five main points of Dalton’s atomic theory are:-
Matter is made up of tiny particles called atoms, which cannot be split into smaller particles. (In Greek atom = unsplittable)
Atoms cannot be created or destroyed.
The atoms of any one element are identical and have the same chemical properties and same mass.
The atoms of a given element are different from those of any other element. The atoms of different elements can be distinguished from one another by their respective weights.
Atoms of one element can combine with the atoms of another element to form compound atoms otherwise known as Molecules. The atoms always combine in simple ratios.

Modification of Dalton’s Atomic Theory
Atoms can be created or destroyed or split by means of nuclear reactions. For example an atom of uranium – 235 can be split into two separate atoms by nuclear fission.
Some elements have atoms of more than one kind, which differ slightly in mass. Such atoms are called Isotypes of the element. For example, carbon has three isotypes known as Carbon-12, Carbon-13 and Carbon-14.
An atom is made up of even smaller sub-atomic particles called protons, neutrons and electrons.
Atoms of different elements may combine in many different ratios to from complex compounds.

Sub-Atomic Particles
An atom consists of a very small and dense region called Nucleus. Nucleus consists of protons and neutrons. The nucleus is surrounded by shells/orbit. In every orbit there are electrons. The main sub-atomic particles are protons, neutrons and electrons.
Protons
These are positively charged particles.
One proton has a mass of one atomic mass unit, which is equal to that of hydrogen.
Protons are found in the nucleus and are denoted by the symbol P.
Neutrons
They are denoted by the letter n.

The properties of neutrons are:
a) They have no charge – Neutral.
b) They are located in the nucleus of an atom.
c) They have nearly the same mass as corresponding protons.
d) They have a mass nearly1840 times the mass of an electron.

Sub-Atomic
Symbol
Location
Charge
Real Mass
Relative Mass
Proton
P
In the nucleus
+1
1.6726 X 10-24
1
Neutron
n
In the nucleus
0
1.6750 X 10-24
1
Electron
e–
Outside the nucleus
-1
  9.109 X 10-28
1/1840

ELECTRONIC CONFIGURATION
Is the arrangement of electrons in different energy levels in an atom.
The electrons orbit the nucleus in special regions called energy levels.
The energy levels are fixed at a distance from the nucleus.
The shells can hold a maximum number of electrons, which are determined by the formula 2n2 where n is the position of the energy level from the nucleus.
The energy levels are represented by the letters K, L, M, N from the nucleus respectively according to the formulae.

According to this formulae:-
– The first energy level can hold (2 X 12) = 2 electrons
– The second level can hold (2 X 22) = 8 electrons
– The third level can hold (2 X 32) = 18 electrons
– The fourth level can hold (2 X 42) = 32 electrons

Electronic Structure Of An Atom In Writing And Diagrams.
Atomic Number
The atomic number is the number of protons in an atom. It is also known as the proton number e.g. the atomic number of hydrogen is 1 since it has only one proton. A sodium atom has 11 protons in the nucleus. It’s atomic number is therefore 11.
Mass Number
Protons and neutrons are found in the nucleus of an atom and are called nucleons. The sum of protons and neutrons in one atom of an element is called mass number or nucleon number. Thus:
Number of Protons + Number of Neutrons = Mass Number
Example
Hydrogen has 1 proton and 0 neutrons. Therefore, it’s atomic number is 1 and mass number is 1 + 0 = 1.

Element
Symbol
No. of Electrons
No. of Protons
Electronic Structure
Hydrogen
H
1
1
.1
Helium
He
2
2
.2
Lithium
Li
3
3
2.1
Beryllium
Be
4
4
2.2
Boron
B
5
5
2.3
Carbon
C
6
6
2.4
Nitrogen
N
7
7
2.5
Oxygen
O
8
8
2.6
Fluorine
F
9
9
2.7
Neon
Ne
10
10
2.8
Sodium
Na
11
11
2.8.1
Magnesium
Mg
12
12
2.8.2
Aluminium
Al
13
13
2.8.3
Silicon
Si
14
14
2.8.4
Phosphorus
P
15
15
2.8.5
Sulphur
S
16
16
2.8.6
Chlorine
Cl
17
17
2.8.7
Argon
Ar
18
18
2.8.8
Potassium
K
19
19
2.8.8.1
Calcium
Ca
20
20
2.8.8.2

STRUCTURE OF ATOMS
– Number of Protons = Atomic Number = Number of Electrons
– Number of Protons + Number of Neutrons = Mass Number/Atomic Mass
– Protons and neutrons are in the nucleus and are called nucleons.

Draw a diagram to show the structures of:
Sodium (Na23)
Phosphorus – 31 (15P31)
Chlorine – 37 (11Cl37)

Nuclide Notation
Atoms of different elements can be represented by symbols that indicate their respective atomic numbers and mass numbers. Eg.Using element X.
Atomic number (Z) as a subscript and atomic mass (A) is superscript.
AzX – Nuclide Notation

Isotopes
Isotopes are atoms of the same element, having same atomic number but different atomic mass due to different numbers of neutrons.
Isotopy is the existence of atoms of same elements having same atomic number but different atomic mass.

Examples of Isotopes
Element
Symbol
Atomic No.
Isotopes
Abundance
Hydrogen
H
1
11H
99.99%
21H
0.01%
31H
Very rare
Carbon
C
6
126C
98.9%
136C
1.1%
146C
Trace
Chlorine
Cl
17
3517Cl
75%
3717Cl
25%
Oxygen
O
8
168O
99.8%
178O
0.037%
188O
0.20%
Neon
Ne
10
2010Ne
90.5%
2110Ne
0.3%
2210Ne
9.2%

Relative Atomic Mass
An atom is very small and it’s mass would be very difficult to measure. To overcome this difficulty, chemists made a simpler way to express the mass of an atom. This involved expressing the mass of an atom in relation to a chosen standard atomic mass.
The carbon atom was chosen as the standard atom and it’s mass was arbitrarily chosen as 12 units. Then using a machine called mass spectrometer, all the other atoms were compared tothis standard atom. This reference was called carbon – 12 scale. For example, it was found that:
Magnesium atom was twice as heavy as reference atom, so it’s mass was 24.
The hydrogen atom was 1/12 as heavy as reference atom, so it’s mass was put at 1.
Helium atom was 1/3 as heavy as the reference atom so it’s mass was 4.

Note: The relative atomic mass of an element is the average mass of one atom of the element relative to 1/12th the mass of 1 atom of Carbon – 12 i.e.
Ar = Average mass of an atom of an element
1/12th the mass of Carbon – 12 atom

The relative atomic mass of the elements is calculated from the atomic masses of the different isotopes and their abundances.
Eg.
1) Chlorine elements exist in 2 isotopic forms
Chlorine
Cl
17
3517Cl
75% abundance
3717Cl
25% abundance

Ar = (Atomic mass of Isotope I X % of abundance) + (Atomic mass of Isotope II X % of abundance)
= (35 X 75/100) + (37 X 25/100)
= 35 X 75 + 37 X 25
100
= 2625 + 925            =          3550                =          35.5
100                              100

Uses of Isotopes
Isotopes can be used to trace the path of certain elements in a biological process e.g. photosynthesis uses 14C.
Used to determine leakage in underground petroleum lines.
They are used in dating fossils.
Radioactive isotopes can be used to produce nuclear energy.
Agricultural applications – Irrigation.
Medical uses – Used to evaluate organ function.
Used in smoke detectors.
Used in scientific research.
Used in industries.

5.PERIODIC CLASSIFICATION
The Periodic Table
– It is the table showing the arrangement of elements in order of increasing atomic numbers.
– It is a table of elements arranged in order of increasing atomic number to show the similarities in chemical properties in relation to electronic structure.
The Periodic Law
‘The properties of the elements are a period periodic function of their atomic number.’
The periodic table has 7 periods and 8 groups. Metals appear on the left of the table. Metals have 1, 2, 3 electrons in the outmost shell. Non-metals have 4, 5, 6, 7, 8 electrons in the outmost shell and appear on the right hand side of the table.

There are 5 blocks of similar elements in the table
Note: Elements with the same number of shells appear in the same period.
Elements in the same group appear to have similar properties.

Therefore, the chemical and physical properties of the elements depend on the number and arrangement of electrons.
Groups
Periods
I
II

III
IV
V
VI
VII
VIII
1
H1
He2
2
Li3
Be4
B5
C6
N7
O8
F9
Ne10
3
Na11
Mg12
Al13
Si14
P15
S16
Cl17
Ar18
4
K19
Ca20
Transition Elements


Families of the PT
Alkali Metals – Elements in group I
They appear in the block of reactive metals as they are very reactive.
They catch flames when exposed to air and react with cold water to form an alkaline solution.

Alkaline Earth Metals
These are group II elements. Also appear in the reactive block. Most of their compounds are found in rocks. Some react with water to form an alkaline solution e.g. Ca, Mg.

Transition elements
– Are metals with high tensile strength.
– Have variable valency.
– Form coloured compounds

Poor Metals (Metalloids)
Elements in this group have some metallic and non-metallic properties e.g. Ge, Si, As.

Non-Metals
– These have 5, 6, 7, 8 electrons in outer shell.
– Some are very reactive and others are not.
– Noble gases are gases at room temperature.
– Have 2 or 8 electrons in the outmost shell.
– Are very stable and most are unreactive.

Halogens
Are group VIII elements.
They have seven electrons in the outmost shell and very reactive.
React with metals to form salts. Thus halogen means salt producer.

Properties Of The Elements In The Periodic Table
The trends observed include variations in:-
(i) Melting point – Is the temperature where solid melts to liquid.
(ii) Boiling point – Temperature at which liquid boils to form gas.
(iii) Density – Is the degree of compactness of a substance, which means it is the mass per unit volume of a substance.
(iv) Electronegativity – Is the ability of an atom to attract an electron.
(v) Ionization energy – Is the energy required to remove electrons from an atom or ion.
(vi) Atomic radius – Is the distance between the nucleus of an atom and the outmost stable energy level.
(vii) Reactivity – Refers to how likely an atom of a given element reacts with other substances.

Properties Of The Elements In The Periodic Table
Atomic Size
Atomic size increases down the group. This is due to increase in number of shells. E.g. 3Li = 2.1, 11Na = 2.8.1, 19K = 2.8.8.1.
Atomic radius increases across the table due to increase in number of electrons in the outmost shell, which leads to increase in force of attraction towards positive nucleus.
Ionization Energy
Is the energy required to remove an electron from an atom. It decreases down the group due to increase in number of shells. It increases across the period due to increase in nuclear attraction.
Electronegativity
Is the tendency of an element to attract electrons to itself. It increases across the period due to increase in number of electrons while the number of shells remains constant.
Metallic Character
Electropositivity is the tendency of an element to lose an electron when supplied with energy. It increases down the group to increase in number of shells. It decreases across the period due to increase in number of electrons in the outer shell. When atoms loose or gain electrons, they form ions.

Specific Trends In Groups
Group I – Alkali Metals
Consists of five metals namely: Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs). They have one electron each in their outer energy level.

Group I
Name
Atomic No.
Electronic Configuration
Atomic Radius
1st Ionization Energy
Melting Point
Density
Electronegativity
Lithium
3
2.1
152
526
180
0.54
1.0
Sodium
11
2.8.1
186
504
98
0.97
0.9
Potassium
19
2.8.8.1
231
425
64
0.86
0.8
Rubidium
37
2.8.18.8.1
244
410
39
1.5
0.8
Caesium
55
2.8.18.18.8.1
262
380
29
1.9
0.7

Group II
Name
Atomic No.
Electronic Configuration
Atomic Radius
1st Ionization Energy
Melting Point
Density
Electronegativity
Beryllium
4
2.2
112
899
14,849
1,280
1.5
Magnesium
12
2.8.2
160
738
7,730
651
1.2
Calcium
20
2.8.8.2
197
590
4,741
851
1.0
Strontium
38
2.8.18.8.2
215
549
4,207
800
1.0
Barium
56
2.8.18.18.8.2
217
503
3,420
850
0.9

Physical properties of Group I
Good conductors of heat and electricity.
Soft metals.
Low density.
Shiny surfaces when freshly cut.
Chemical properties of Group I
Burn in O2 or air with a flame color to form white solid oxides.
Metal + O2 = Metal oxide
React with water to give the alkaline solution and hydrogen.
Metal + Water = Metalhydroxide + Hydrogen

Physical properties of Group II
Harder than those in group I.
Are silver-grey when clean and pure.
Good conductors of heat and electricity.

Chemical properties of Group II
Burn in O2 or air to form a solid white oxide.
Metal + O2 = Metal oxide
The metals become more reactive as we move down the group.
React with water but less vigorously than those in group I.
Metal + Water = Metal hydroxide + Hydrogen

6.VALENCY, BONDING AND NOMENCLATURE
Valency is the combining power of an element.
For groups I, II, III and IV of PT, the valency of elements is number of shells.
For groups V to VIII, the normal valency is 8 – number of electrons in outmost shell. Eg. 147N 2.5 Valency 8 – 5 = 3
Other common elements and their valencies are:
Silver Ag         Valency 1
Copper Cu       Valency 1 or 2
Iron Fe            Valency 2 or 3
Mercury Hg     Valency 1 or 2
Zinc Zn           Valency 2

Radicals
A radical is a group of atoms with unpaired electrons.

Most radicals form the non-metallic part of a compound, so their ions are negatively charged. Examples are CO2-3 and SO2-4 ions. An exception is the ammonium radical, NH4, which behaves like the metallic part of a compound and forms a positive ion, NH+4.

The valency of a radical is the same as the numerical value that the group acquires when it loses or gains an electron to form an ion.
Radicals
Formula
Stable Ion
Valency
Nitrate
NO3
NO–3
1
Nitrite
NO2
NO–2
1
Sulphate
SO4
SO2-4
2
Hydrogen Sulphate
HSO4
HSO–4
1
Carbonate
CO3
CO2-3
2
Hydrogen Carbonate
HCO3
HCO–3
1
Hydroxyl
OH
OH–
1
Phosphate
PO4
PO3-4
3
Thiosulphate
S2O3
S2O2-3
2
Cyanide
CN
CN–
1
Permanganate
MNO4
MNO–4
1
Oichromate
Cr2O7
Cr2O2-7
2
Ammonium
NH4
NH+4
1

CHEMICAL FORMULA
A formula is a collection of symbols of elements and numbers, which indicate a molecule of an element or compound.
It is a representation that uses symbols to show the proportions of the elements present in a chemical compound.

Rules of writing a formula.
Positively charged ions are written before the negatively charged ions.
Any number written in the formula is written as subscript.
The valency of each element, radical is written ender each element respectively.
Radicals are treated as a single ion and if a need arise, a bracket is used.
The valency 1 is assumed not written in the formula.
Single elements are not bracketed.

Aluminium Oxide
Al                    O
3                      2
2                      3

OXIDATION STATE
Oxidation state is a measure of the degree of oxidation of an atom in a compound.

The oxidation number of an element indicates the number of electrons lost, gained or shared by an atom of the element, with respect of its neutral atom. The natural atom has no charge.

The following rules are used to assign oxidation states to elements:-
In free elements, each atom has an oxidation number of zero no matter how complicated it’s molecule is. For example, nitrogen, hydrogen, sodium and oxygen all have oxidation number zero.
In simple ions that consist of only one atom, the oxidation number is equal to the charge on the ion. For example, the oxidation of a sodium ion is +1, aluminium is +3, iron II is +2, and iron III is +3. In an oxide ion, the oxidation number of O2 is -2.
Hydrogen has an oxidation state of +1 in most compounds. The exception is in hydrides of active metals where oxidation number is -1. For example, the hydrogen atom gains an electron from the lithium atom in lithium hydride (LiH).
Oxygen has an oxidation state of -2 when present in most compounds except:                                                           a) In peroxides, eg. H2O2, where the oxidation number is -1.                                                                                           b) When bonded with fluorine to from F2O the oxidation number is +2 and of fluorine is -1.
All oxidation numbers must be consistent with the conservation of charge. This means that:
a) For all neutral molecules, the oxidation number of all the atoms must add up to zero. For example, in H2O, two hydrogen atoms each of charge +1 combine with one oxygen atom of charge -2. The charge of the molecule is +2-2=0. b) For complex ions, the oxidation numbers of all atoms must add up to the charge on the ion.

Note: Valency is a fixed value while oxidation state is an arbitrary value.
Sulphur in Potassium Sulphate (K2SO4)
S. of K + O. S. of S + O. S. of O = 0
(+1X2) + S + (-2X4) = 0
+2 + S + (-8) = 0
S = +6

7.EMPIRICAL FORMULAE AND MOLECULAR FORMULAE
Empirical formula is the one, which expresses the composition of elements by mass.
Molecular formula is the one, which shows the exact number and kind of atom in a compound or molecule.
A certain organic compound contains 80% by weight carbon and 20% by weight hydrogen. Calculate:
(a) Empirical formula
Carbon
Hydrogen
% Age
80
20
% Age/Ar
80/12
20/1
EF = CH3

= 6.67
= 20
Divide by the smallest
6.67/6.67
20/6.67
(b) Molecular formula if it’s weight is 30.
Empirical Units (n) = Molecular formula mass
Empirical formula mass
(CH3)n = 30
12n + 3n = 30
15n = 30
n = 2                                        MF = C2H6

BONDING
Bonding is a method by which atom becomes stable.
Atoms can be stable by:
a) Donating and gaining electrons.
b) Sharing electrons.
Only the outmost electrons
Bonding occurs between:
a) Metal and non-metal elements.
b) Non-metal and non-metal elements.
A bond is electrostatic form of alteration between two or more atoms.
Atoms bond to attain the noble gases structure, thus noble gas stability.
Noble gases are stable because they have 2 or 8 electrons in the outermost shell. E.g. 2He = 2, 10Ne = 2.8, 18Ar = 2.8.8. The 2 and 8 electrons in the outermost shell is referred to as duplet and octet structure respectively.

Types of bonding
Covalent Bonding
Covalent bonding is the one, which involves sharing of electrons.
Properties of Covalent Compounds
Covalent compounds are gas or liquid at room temperature.
Covalent compounds have low melting point and boiling point (due to the weaker forces between the molecules)
Covalent compounds do not conduct electricity because they do not have ions.
They are soluble in organic solvents eg. petrol, kerosene.
They are in term of molecules.
Electrovalent/Ionic Bonding
This type of bonding occurs between molecules of non-metal elements.
Metals have 1, 2 or 3 electrons in the outermost shell. Therefore, they loose/donate these electrons thus becoming positively charged ions.
Eg.13Al = 13Al31 + 3e–
2.8.3      2.8
The non-metal element gains the electrons and become negatively charged.
Eg. O2
O + 2e- = O2-
2.6        2.8

The force of attraction between the opposite charged end is ionic bonding.
Electrovalent bond is the one formed by transfer of electrons.
Consider in NaCl
– Na has electronic configuration 2.8.1
– It becomes stable by loss of 1 electron and gain the octet structure.
Na = Na+ + e
2.8.1    2.8.1

Chlorine has electronic structure 2.8.7. To acquire stability it gains 1 electron from sodium.
Properties of Ionic Compounds
Ionic compounds occur in form of ions. Thus they are crystalline solids at room temperature.
They have high boiling and melting points.
They conduct electricity in molten form or in solution but not in solid form.
Are soluble in water and insoluble in organic solvents eg. petrol, ethanol etc.
Differences Between Covalent + Ionic Compounds
Covalent Compounds
Electrovalent Compounds
1.
They are liquids, gases at room temperature.
Solids at room temperature.
2.
Do not conduct electricity.
Conduct electricity in solid form.
3.
Usually have low melting and boiling point.
Have high melting and boiling point.
4.
Soluble in organic solvents.
Insoluble in inorganic solvents.
5.
Occur in form of molecules.
Occur in form of ions.