1. Periodic Table of Elements

2. Analyse the Periodic Table of Elements
Analyse Group 18 elements
Analyse Group 1 elements
Analyse Group 17 elements
Analyse elements in a period
Understand transition elements

3. Antoine Lavoisier (1743 – 1794)
Classify substances into
metals and non-metals
Unsuccessful because light,
heat and some other
compounds where not
elements.

4. Johann Dobereiner (1780 - 1849)
Introduced triads.
Elements were classified
into groups of three elements with
same chemical properties
The atomic mass of middle
elements was approximately the
average atomic mass of the other
two elements

5. Lothar Meyer (1830 - 1895)
Plotted a graph of the atomic
volume against atomic mass.
Elements with similar chemical
properties occupied same
positions.
Successful in showing the
properties of elements formed a
periodic pattern against their
atomic masses.

6. John Newlands (1837 - 1898)
Arranged elements in order of
increasing atomic mass.
Elements with similar properties
recurred at every eight element.
This was known as the Law of
Octaves
Failed because only obeyed by
first 17 elements only

7. Dimetri Mendeleev (1839 – 1907)
Arranged elements in order of
increasing atomic mass
Elements with similar chemical
properties are grouped together
He left empty spaces in the table
for undiscovered elements

8. Henry J. G. Moseley (1887-1915)
Concluded that proton
number should be the bases
for the periodic change of
chemical properties
Arranged the elements in
order of increasing proton
number in the Periodic Table.

9. Elements are arranged according their increasing
order of proton number.
Vertical columns = groups(according to their number
of valence electrons)
Horizontal rows = periods (number of electron shells
filled by electrons)

11. Known as noble gases/inert gases(chemically
unreactive elements)
Non-metals that exist as monoatomic colourless
gases.
Members : Helium(He)
Neon(Ne)
Argon(Ar)
Krypton(Kr)
Xenon(Xe)
Radon(Rn).

12. Very small atomic sizes.
Low melting and boiling points
Weak van der Waals’ forces of attraction
between atoms.
Low densities
Very small masses but huge volumes.
Melting and boiling points of elements increase
down the Group 18.

13. All Group 18 elements are chemically
inert/unreactive.
The outermost electron shell of each member is fully
occupied by electrons.
This is a stable electron arrangement which in
Helium, it is said achieve duplet electron
arrangement.
Other than Helium, it is said achieve octet electron
arrangement.

14. Helium
To fill airships and weather
balloons.
used as artificial atmosphere in
oxygen tank for divers.

15. Neon
Advertising lights
Used in aeroplane runway
lights

16. Argon
To fill light bulbs.
Krypton
Used in lasers to repair
the retina of the eye.
To fill photographic flash
lamps.

17. Xenon
Making electron tubes and
stroboscopic lamps
Used in bubble chambers
in atomic energy reactors.
Radon
Used to treat cancer

18. Consists of lithium(Li), sodium(Na), potassium(K),
rubidium(Rb), caesium(Cs) and francium(Fr).
Li Na K Rb Cs Fr
They are known as alkali metals because they react
with water to produce alkaline solution.

19. Soft
Low melting points
Low densities
Shiny and silvery surface
Good conductor of heat
Good conductor of electricity

20. Hardness, melting point and boiling point
of the elements decreases going down the
group.
When go down Group 1, size of atom becomes
larger. The positive nucleus gets further away from
the negative sea of electrons.
The force of attraction between the metal ions and
the sea of electrons gets weaker down the group.
Less energy is needed to overcome this
weakening force of attraction.

21. 1. All react with water to produce alkaline metal
hydroxide solution and hydrogen gas.
2X(s) + 2H2O(l)  2XOH(aq) + H2(g)

22. 2. All burn in oxygen gas to produce white solid metal
oxides.
4X(s) + O2(g)  2X2O(s)
The oxide dissolve in water to form alkaline metal
hydroxide solution.
X2O(s) + H2O(l)  2XOH(aq)
3. All burn in chlorine gas to produce white solid
metal chlorides.
2X(s) + Cl2(g)  2XCl(s)

23. Why the reactivity of elements increases down the Group
1?
Atomic size of Group 1 elements increases from lithium to
francium//Number of shells occupied by electrons increases.
Distance between the valence electron in the outermost shell
and positive nucleus increases down the Group 1.
Attraction between nucleus and valence electron decreases.
It is easier for the atom to lose the valence electron to achieve
stable electron arrangement.
Why all elements in Group 1 have same chemical
properties?
Chemical reaction is all about the activity of electrons
All the elements have one valence electron.
Each of them reacts by donating one valence electron to form
an
ion with a charge of +1 to achieve stable electron arrangement.

24. Members are fluorine(F2) , chlorine(Cl2),
bromine(Br2), iodine(I2), and astatine(At2)
F Cl Br I At
The elements are also known as halogens which exist
as diatomic molecules.

25. They have low melting and boiling points because
molecules are attracted to each other by weak van der
Waals’ forces of attraction.
The melting and boiling points of the elements increases
down Group 17.
This change the states of elements from gas to solid and
the colour of elements from lighter colour to darker
colour.

26. Elements State Colour
Fluorine Gas Pale yellow
Chlorine Gas Greenish-yellow
Bromine Liquid Reddish-brown
Iodine Solid Purplish-black

27. Why the melting and boiling points of elements
increases down Group 17?
Molecular size/relative molecular mass of the
elements increases down Group 17.
Forces of attraction between
molecules/Intermolecular forces of attraction
increases.
More heat is needed to overcome the stronger forces
of attraction between the molecules.

28. All members have similar chemical properties but differ in
the reactivity.
1. React with water to form two acids
X2(g) + H2O(l) HX(aq) + HOX(aq)
Example:
Cl2(g) + H2O(l) HCl(aq) + HOCl(aq)
hydrochloric hypochlorous
acid acid
Hypochlorous acid is a bleaching agent (bleach both blue and
red litmus paper)

29. 2. Halogens in gaseous state react with hot iron to form brown
solid.
2Fe(s) + 3X2(g)  2FeX3(s)
Example:
2Fe(s) + 3Cl2(g)  2FeCl3(s)
solid iron(III) chloride(brown)
3. Halogens react with sodium hydroxide solution to produce
sodium halide, sodium halate(I) and water
X2 + 2NaOH(aq)  NaX(aq) + NaOX(aq) + H2O(l)
Example:
Cl2 + 2NaOH(aq)  NaCl(aq) + NaOCl(aq) + H2O(l)
Sodium chlorate(I)

30. Why all halogens possess similar chemical properties?
Chemical reaction = lose or accept electrons
All halogens always gain one electron to achieve stable octet
electron arrangement.
Therefore, they have similar chemical properties.
Why chemical reactivity of halogens decreases down Group 17?
Atomic size/number of electron occupied shells of halogens
increases down Group 17.
The outermost shell becomes further from the nucleus of the
atom.
Strength to attract one electron into the outermost shell by the
nucleus becomes weaker.
Reactivity decreases.

31. Elements across a period exhibit a periodic change in
properties.
Proton number increases by one unit from one element to
the next element

32. All the atoms of the elements have three shells occupied
with electrons
The number of valence electrons in each atom increase from
1 to 8

33. All the elements exist as solid except chlorine and argon
which are gases
The atomic radius of elements decreases. This is due to the
increasing nuclei attraction on the valence electrons.
The electronegativity of elements increases. This is also due
to the increasing nuclei attraction on the valence electrons
and the decreases in atomic size.

34. Uses of metalloid
Make diodes and transistors
A diode A transistor
Both are commonly used in the making of microchips
Microchips are widely used in the manufacture of computers,
mobile phones, televisions, video recorders, calculators, radio
and etc.
Metalloid – semi-metal, reacts with acid only, weak
conductor, brittle and not malleable and ductile.

35. Oxides of elements change from basic to amphoteric
and then to acidic across the period towards the right.
Basic oxides – react with acids to form salt and water
Acidic oxides – react with alkalis to form salt and water
Amphoteric oxides – react with both acids and alkalis to
form salt and water.

36. Elements from Group 3 to Group 12 in the
Periodic Table.
Common characteristics
Solid metal with shiny surface.
Good conductor of heat and electricity.
High melting and boiling points.
Hard, malleable and ductile.

37. Special characteristics;
Show different oxidation numbers in their
compounds
Form coloured ions or compounds
Use as catalysts
Form complex ions

38. Show different oxidation numbers in their compound
Compound Formula Oxidation number
Chromium(III) chloride CrCl3 +3
Potassium dichromate(VII) K2Cr2O7 +6
Manganese(II) sulphate MnSO4 +2
Manganese(VI) oxide MnO2 +4
Potassium manganate(VII) KMnO4 +7
Iron(II) sulphate FeSO4 +2
Iron(III) chloride FeCl3 +3
Copper(I) oxide Cu2O +1
Copper(II) sulphate CuSO4 +2

39. Form coloured ions or compounds
Element Ion Colour
Chromium
Cr3+ Green
CrO4
2-
Yellow
Cr2O7
2-
Orange
Manganese
Mn2+ Pale pink
MnO4
-
Purple
Iron
Fe2+ Pale green
Fe3+ Yellowish brown
Cobalt Co2+ Pink
Nickel Ni2+ Green
Copper Cu2+
Blue Green

40. Form coloured ions or compounds
Gemstone Transition metal Colour
Emerald
Ni and Fe Green
Amethyst
Fe and Mn Purple
Sapphire
Co and Ti Blue
Ruby
Cr Red
Topaz
Fe Yellow

41. As catalyst
Process Catalyst
To
manufacture
Haber Process Iron fillings, Fe Ammonia
Contact Process
Vanadium(V)
oxide, V2O5
Sulphuric acid
Ostwald Process Platinum, Pt Nitric acid
Hydrogenation Nickel, Ni Margarine
HAI
CSV
ONiP

42. To form complex ions
Element Complex ions Formula
Iron
Hexacyanoferrate(II) ion [Fe(CN)6]4-
Hexacyanoferrate(III) ion [Fe(CN)6]3-
Chromium
Hexaamina chromium(III)
ion
[Cr(NH3)6]3+
Copper
Tetraamina copper(II) ion [Cu(NH3)4]2+
Tetrachlorocuprate(II) ion [CuCl4]2-