Periodic Table of Elements

Analyse the Periodic Table of ElementsAnalyse Group 18 elementsAnalyse Group 1 elementsAnalyse Group 17 elementsAnalyse elements in a periodUnderstand transition elements

Antoine Lavoisier (1743 – 1794)Classify substances into

metals and non-metalsUnsuccessful because light,

heat and some other compounds where not elements.

Johann Dobereiner (1780 - 1849)Introduced triads. Elements were classified

into groups of three elements with same chemical properties

The atomic mass of middleelements was approximately the average atomic mass of the other two elements

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.

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

OctavesFailed because only obeyed by

first 17 elements only

Dimetri Mendeleev (1839 – 1907) Arranged elements in order of

increasing atomic massElements with similar chemical

properties are grouped togetherHe left empty spaces in the table

for undiscovered elements

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.

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)

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).

Very small atomic sizes. Low melting and boiling points

Weak van der Waals’ forces of attraction between atoms.

Low densitiesVery small masses but huge volumes.

Melting and boiling points of elements increase down the Group 18.

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 inHelium, it is said achieve duplet electron

arrangement.Other than Helium, it is said achieve octet

electron arrangement.

HeliumTo fill airships and weather

balloons.

used as artificial atmosphere in oxygen tank for divers.

NeonAdvertising lights

Used in aeroplane runway lights

ArgonTo fill light bulbs.

KryptonUsed in lasers to

repair the retina of the eye.

To fill photographic flash lamps.

XenonMaking electron tubes

and stroboscopic lamps

Used in bubble chambers in atomic energy reactors.

RadonUsed to treat

cancer

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.

Soft Low melting pointsLow densitiesShiny and silvery surfaceGood conductor of heatGood conductor of electricity

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.

1. All react with water to produce alkaline metal hydroxide solution and hydrogen gas.

2X(s) + 2H2O(l) 2XOH(aq) + H2(g)

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)

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 electronsAll the elements have one valence electron.Each of them reacts by donating one valence electron

to form anion with a charge of +1 to achieve stable electron

arrangement.

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.

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.

Elements State Colour

Fluorine Gas Pale yellow

Chlorine Gas Greenish-yellow

Bromine Liquid Reddish-brown

Iodine Solid Purplish-black

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.

All members have similar chemical properties but differ in the reactivity.

1. React with water to form two acidsX2(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)

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)

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.

Elements across a period exhibit a periodic change in properties.

Proton number increases by one unit from one element to the next element

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

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.

Uses of metalloidMake diodes and transistorsA diode A transistor

Both are commonly used in the making of microchipsMicrochips 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.

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 waterAcidic oxides – react with alkalis to form salt and

waterAmphoteric oxides – react with both acids and alkalis

to form salt and water.

Elements from Group 3 to Group 12 in the Periodic Table.

Common characteristicsSolid metal with shiny surface. Good conductor of heat and electricity.High melting and boiling points.Hard, malleable and ductile.

Special characteristics;

Show different oxidation numbers in their compounds

Form coloured ions or compoundsUse as catalystsForm complex ions

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

Form coloured ions or compounds

Element Ion Colour

Chromium

Cr3+ Green

CrO42- Yellow

Cr2O72- Orange

ManganeseMn2+ Pale pink

MnO4- Purple

IronFe2+ Pale green

Fe3+ Yellowish brown

Cobalt Co2+ Pink

Nickel Ni2+ Green

Copper Cu2+ Blue Green

Form coloured ions or compoundsGemstone Transition metal Colour

EmeraldNi and Fe Green

AmethystFe and Mn Purple

SapphireCo and Ti Blue

RubyCr Red

TopazFe Yellow

As catalyst

Process CatalystTo

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

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-