A Short History of The Periodic Table

Johann Dobereiner - Law of Triads

Nature contained triads of elements

The middle element had properties that were an average of the other two members when ordered by

the atomic weight (the Law of Triads).

John Newlands - Law of Octaves

Any given element will show similar behavior to the eighth element following it in the table.

Lothar Meyer - Meyer’s Curve

Plotted the atomic volumes of the elements against the atomic weight, and found that the chemical

properties of the element recur periodically.

Also notice that the elements occupying the corresponding position of the curve show similar

chemical properties.

Mendeleev - First Periodic Table

arrange the elements according to the ascending order of atomic mass and put all the elements that

have same chemical properties in the same group.

left empty space in the periodic table for elements that haven't been discovery at that time.

Although he arranged the elements in ascending order of the atomic mass, he changed the order if the

chemical properties of the element did not match its group.

From the empty space present in his table, he even predicted the existence and properties of unknown

elements which he called eka-aluminum, eka-boron, and eka-silicon.

H.J.G. Moseley - Modernized the Periodic Table by Arranging The Elements According to the Ascending

Order of Atomic Number (Proton Number)

managed to measure the proton number of atoms.

arranged the elements in the periodic table according to the ascending order of the atomic number

(proton number), but not the atomic mass, as done by Mendeleev.

managed to predict the existence of four undiscovered elements from the proton number.

Modern Periodic Table

1. In modern periodic table, the elements are arranged in ascending order of proton number.

2. The vertical columns of the periodic table is called GROUPs whereas the horizontal rows is called the

PERIODs.Group

1. The vertical columns of the periodic table are called Group.

2. There are 18 groups in the periodic table.

3. All the elements in the same group have same number of valence electron.

4. Thus, elements in the same group exhibit similar chemical properties.

5. The Group is named according to their position in the periodic table. For example, the first group is

called Group 1, the second group is called Group 2 and so on.

6. Group 3 to Group 12 are categorised to another big group, called the Transition Metal Group.

7. Group 1, 2, 17 and 18 have trivial name, as shown in the table below.

Group Trivial NameGroup 1 The Alkali Metals

Group 2 The Earth Alkaline Metals

Group 17 The Halogens

Group 18 The Noble Gases

Period

1. The horizontal rows is called the PERIOD.

2. There are 7 periods in period table.

3. The first period only has 2 elements only.

4. The second and third period consist of 8 elements, are called the short period.

5. The forth and the fifth period consist of 18 elements, are called the long period.

6. The sixth and the seventh period has 32 elements.

Periodic Table and Electrons Arrangement1. The classification of Group and Period are totally related to the electron arrangement of the atoms. As

shown in the table above, the classification of elements in a periodic table is as follows:

a. All the elements in a same group have equal number of valence electron(s) (outer most

electron)

b. All the elements in a same period have equal number of electron shell(s) (orbit).

2. For example, the electron arrangement of calcium is 2.8.8.2. It has 2 valence electron and 4 electron

shell. Therefore, it is placed in Group 2, period 4.

3. Hydrogen, the simplest element atom, with 1 electron, does not fit into any group. (Even though some

of the periodic table place it in Group 1, due to its 1 (and the only one electron) valence electron.

Number of Valence Electron in a GroupGroup Number of Valence

ElectronGroup 1 1Group 2 2Group 13 3

Group 14 4

Group 15 5

Group 16 6

Group 17 7

Group 18 8

Group 18 Elements - Noble Gases1. The "Noble Gases" are the last group in the Periodic Table, they also known as "inert gas", due to their

non-reactive behavior.

2. This group consist of six elements, namely Helium, Neon, Argon, Kripton, Xenon and Radon.

3. They are non-metallic, colourless gases at room temperature and pressure with very low melting

points and boiling points.

4. They form 1% of air, and most of this is argon.

5. The size of atom increases down the group, due to the increase number of electron shell.

Physical PropertiesSolubility and Conductivity

1. All noble gas are insoluble in water.

2. They are not conductor of heat and electricity.

Melting Point and Boiling Point

1. As show in the table above, the melting and boiling point of noble gases are very low.

2. This is because all noble gases exist as monoatoms. The force in between all these atoms is the weak

van de Waals' Force.

3. Therefore very little energy is needed to overcome this force during melting and boiling.

Density

1. The density of noble gases are very low.

2. Nevertheless, the density increases steadily down the group.

3. Density of a substance is given by the equation "Density=Mass/Volume".

4. Down the group, both the mass and the volume increase, but increase of mass is faster than the

volume, hence the density increases down the group.

Chemical Properties

1. All the noble gases are non-reactive elements.

2. This is because their valence shell is full of electrons.

3. In the chemical world, an atom is in the chemically most stable state if their valence shell is full with

eight electrons (or 2 electrons for the first shell.).

4. Therefore all noble gases do not react with other elements, due to their stable electronic structure.

5. They exist as single atoms, that is they are monatomic.

Uses of the Group 18 ElementsHelium

Airship

1. The gas is much less dense than air. Therefore it is used in balloons and

'airships'.

2. Because of its inertness it doesn't burn in air UNLIKE hydrogen which

used to be used in large balloons with 'flammable' consequences.

Deep-sea Diving

1. Helium is also used in gas mixtures for deep-sea divers.

2. This is because the solubility of the helium is very low. Therefore it will not dissolve in the blood

even though the pressure of the surrounding is very high

Neon

1. Neon emits light when high voltage electricity is passed through it.

2. Because of this, it is used in glowing 'neon' advertising signs and fluorescent lights.

Argon

Filament Bulb

1. Argon, like all the Noble Gases is chemically inert.

2. It used in filament bulbs because the metal filament will not burn in Argon and it reduces evaporation

of the metal filament.

WeldingArgon also used to produce an inert atmosphere in high temperature metallurgical processes, eg in welding

where it reduces brittle oxide formation reducing the weld quality.

Fire Extinguisher

Argon is also used for extinguishing fires where damage to equipment is to be avoided

KryptonKrypton is used in fluorescent bulbs, flash bulbs and laser beams.

Xenon

Also used in fluorescent bulbs, flash bulbs and lasers.

Radon

Used in radiotherapy to kill cancer cell

Group 1 Elements - Alkali Metals

1. The Group 1 metals is called the Alkali Metals.

2. This is because they form oxides and hydroxides that dissolve in water to give

alkaline solutions.

3. As shown in the diagram on the right, elements in this group are lithium, sodium,

potassium, rubidium, caesium and francium.

4. They are the first element of a period, with one valence electron.

5. This similarity (1 valence electron) makes them chemically behave in a similar

manner.

6. All alkali metals are very reactive. They must be stored in oil prevent reaction with oxygen or water

vapour in air.

Physical Properties of Alkali Metals

Name Proton number

Electron arrangement

melting point

boiling point

Density g/cm3

Lithium 3 2.1 180ºC 1342ºC 0.53Sodium 11 2.8.1 98ºC 883ºC 0.97Potassium 19 2.8.8.1 63ºC 759ºC 0.86

Rubidium 37 2.8.18.8.1 39ºC 688ºC 1.48Caesium 55 2.8.18.18.8.1 29ºC 671ºC 1.87Francium 87 2.8.18.32.18.8.1 27ºC 677ºC > 1.87

1. All Group 1 metal exist as solid at room temperature and hence have all the typical metallic properties,

such as:

a. good conductors of heat

b. good conductors of electricity,

c. high boiling points,

d. shinny surface (but rapidly tarnished by air oxidation).

2. Nevertheless, Group 1 metals also show some non-typical metallic properties, such as:

a. low melting points,

b. low density (first three float on water),

c. very soft (easily squashed, extremely malleable, can be cut by a knife).

Important trends down the group:

1. size of atoms increases

2. the melting point and boiling point decrease

3. the density increases.

4. the hardness decreases.

Size of Atom

1. Down the group, the size of atom increases.

2. This is due to the increase of number of electron shells.

3. Atom with more shells is bigger than atom with less shells.

Boiling Point and Melting Point

1. The melting point and boiling point generally decrease

down the group.

2. All the atoms of Group 1 metals are bonded together by a

force called metallic bond.

3. The strength of metallic bond depends on the distance

between the atoms. The closer the atoms, the stronger the

bond.

4. Down the group, the size of the atoms increases, causing

the distance of the atoms increases.

5. As the distance between the atoms increases, the metallic bond between the atoms decreases.

6. Therefore, less energy is needed to overcome the metallic bond during melting process.

7. Consequently, the melting point of Group 1 metal decreases down the group.

Density

1. The densities of Group 1 metals are low compare with the other metals.

2. The densities of the first 3 elements (Lithium, Sodium and Potassium) are lower than water. Thus,

they can float on the surface of water.

3. Nevertheless, the density increases steadily down the group.

4. Density of a substance is given by the equation "Density=Mass/Volume".

5. Down the group, both the mass and the volume increase, but increase of mass is faster than the

volume, hence the density increases down the group

Chemical Properties of Alkali Metals1. Group 1 metals are very reactive metals.

2. They all show the same chemical properties.

3. They can react with water and non-metal such as oxygen and chlorine to form

a new compound.

4. The table to the right shows the electron arrangement of all the Group 1

metals. All the atoms of Group 1 metal consist of 1 valence electron.

5. When an alkali metal atoms react, it loses the valence electron to form a positively charged ion. 

Example:

1. They tend to react mainly with non-metals to form ionic compounds.Safety Precaution

2. Alkali metals are very reactive.

3. Therefore it must be kept in paraffin oil to prevent them from reacting with oxygen and water vapour

in the air.

4. We must avoid to hold group 1 metals with bare hand because they may react with water on our hand.

5. We must wear safety goggles and gloves during handling experiment involving group 1 metal.

Alkali Metals React with Chlorine

1. All alkali metals react with chlorine gas to form white metal chlorides salt.

Group 1 Metals + Chlorine Gas → Metal Chloride

2. The metal chlorides salt formed is soluble in water to give a neutral solution of pH 7.

3. The reactivity increases down the group from lithium, sodium to potassium.Example:

Lithium + Chlorine

2Li + Cl2 → 2LiCl

Observation

Lithium burned slowly with a reddish flame . A white solid is produced.

Sodium + Chlorine

2Na + Cl2 → 2NaCl

Observation

Sodium burned brightly with a yellowish flame. A white solid is produced.

Potassium + Chlorine

2K + Cl2 → 2KCl

Observation

Potassium burned very brightly with a purplish flame . A white solid is produce

Alkali Metals React with Oxygen1. Group 1 metals react with oxygen gas produces metal oxides. These metal oxides dissolve in water

produces alkalis.

Group 1 Metals + Oxygen Gas → Metal Oxide

2. Lithium, sodium and potassium form white oxide powders after reacting with oxygen.

3. The white powder is the oxide of lithium, sodium and postassium.

4. When the white powder is dissolved in water, it produces a solution which turned red litmus paper

blue. Which means, these oxides dissolve in water to form strong alkali.

5. The reactivity increases down the group from lithium, sodium to potassium.Example

Lithium + Oxygen

4Li + O2→ 2Li2O

Dissolve in water

Li2O + H2O → 2LiOH

Observation

Lithium burns with red flame and produces white powder immediately after reaction.

Sodium + Oxygen

4Na + O2→ 2Na2O

Dissolve in water

Na2O + H2O → 2NaOH

Observation

Sodium burned with bright yellow flame, forming white powder immediately after reaction.

Potassium + Oxygen

4K + O2→ 2K2O

Dissolve in water

K2O + H2O → 2KOH

Observation

Potassium burned with very bright purplish flame, forming white powder immediately after reaction.

Alkali Metals React with Water

Group 1 metals react vigorously with water produces alkali and hydrogen gas

Group 1 Metals + Water → Alkali + Hydrogen gas

Common Observation

1. Lithium, sodium or potassium floats and move around on the surface of the water and then dissolve in

the water.

Conclusion: Lithium, sodium and potassium are less dense than water.

2. Colourless gas is released around the metal. The gas produces a “pop” sound when ignited with

lighted wooden splinter.

Conclusion: The colourless flammable gas is hydrogen.

3. The solution turn blue when it is tested with universal indicator.

Conclusion: the solution produced is an alkali.

Lithium + Water

2Li + 2H2O → 2LiOH + H2

Observation

Lithium metal moves slowly on the surface of water with 'fizzing' sound.

Sodium + Water

2Na + 2H2O → 2NaOH + H2

Observation

The lump of sodium moves swiftly on the surface of water with 'fizzing' sound

Potassium + Water

2K + 2H2O → 2KOH + H2

Observation

Potassium reacts violently with water, move very fast on the surface of water and burn with lilac flame.

Explaining the Reactivity Trend of the Alkali Metals1 When an alkali metal atom reacts, it loses its valence electron to form a positively charged ion.

Example

Li → Li+ + e

Na → Na+ + e

K → K+ + e

2 As we go down the group from one element down to the next, the atomic radius gets bigger due to an

extra filled electron shell.

3 The valence electron is further and further from the nucleus. Thus the attraction force between the

nucleus and the valence electron become weaker and weaker.

4. This causes the valence electron is easier to be released to form an ion when the atom takes part in a

reaction.

Solubility of the Oxide, Hydroxide and Salt of Alkali Metals1 All the oxide and hydroxide of group 1 metal are soluble in water to form an alkali solution.

2 All the salts (salt of chloride, nitrate, sulphate, carbonate....) of group 1 metals are soluble in water.

The solutions formed are neutral.

Group 17 Elements – Halogens1. Group 17 elements are typical non-metals and also known as halogens.

2. The elements in this group are fluorine. chlorine, bromine, iodine and

astatine.

3. In nature, all halogens exist as diatomic molecules. They are written as F2,

Cl2, Br2, I2 and As2.

4. Most of the halogens exist in the nature as halide salts.

5. Halide is the name given to the ion of halogens. Table below shows the

corresponding halide of the halogen.

Halogen HalideFluorine FluorideChlorine CHlorideBromine BromideIodine Iodide

Physical Properties of Halogens

1. All group 17 elements are non-metals. Therefore they are heat and electricity insulator.

2. Table below shows the electron arrangement and physical properties of group 17 elements.

Size of Atom and Density

1. The atomic size of group 17 elements increases down the group.

2. This is due to the increase of number of electron shell down the group.

3. The density of group 17 elements is also increases down the group.

4. This is because the rate of increment of the atomic mass is higher than the rate of increment of the

volume.

Melting Point and Boiling Point1. As shown in the graph to the right, the melting points and boiling

increase steadily down the group.

2. The physical state at room temperature also change from gas to liquid

and then to solid.

3. This is because the intermolecular attractive force (van der Waals

force) increase with increasing size of atom or molecule.

Chemical Properties of Halogens1. Group 17 elements are very reactive non-metals.

2. The atoms all have 7 valence electrons, makes them have very similar chemical properties.

3. During chemical reaction, the atom gains one electron to form an ion with charge of -1.

4. The reactivity of group 1 decreases down the group.

5. All group 17 elements are poisonous.

6. Astatine is very radioactive.

Safety Precaution

1. Fluorine, chlorine and bromine gases are poisonous.

2. Therefore all the experiments involving these gases should be carried out in a fume chamber.

3. The experiments involve fluorine are nor done in school.

4. This is because fluorine is so reactive that it will react with most of the substance it comes into contact

with.

5. It is very difficult to conduct experiments involving fluorine.

Halogens React with MetalsChlorine + Iron

3Cl2 + 2Fe → 2FeCl3

Observation

The iron wool burns vogorously with bright flame, forming a brown solid after reaction.

Discussion

1. Chlorine react with iron to form brown iron(III) chloride.

2. The reaction of potassium with concentrated hydrochloric acid produces chlorine gas.

2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 5Cl2 + 8H2O

Bromine + Iron

3Br2 + 2Fe → 2FeBr3

Observation

The iron wool glows brightly but less vigorously. A brown solid is formed.

Discussion

Bromine react with iron to form brown iron(III) bromide.

Iodine + Iron

3I2 + 2Fe → 2FeI3

Observation

The iron wool glows slowly with dim light. A brown solid is formed after reaction.

Discussion

Iodine react with iron to form brown iron(III) iodide.

Note:

1. Halogen is a reactive non-metal. It form salt when react with metal.

2. The reactivity decreases down the group.

3. Chlorine gas is a poisonous. Excess chlorine gas is absorbed by the soda lime (sodium hydroxide) so

that it does not escape to the surrounding.

4. Iron wool rather than iron piece is used to increase the rate of reaction.

Halogens React with Water

Chlorine + Water

Cl2 + 2H2O → HCl + HOCl

Observation:

1. Chlorine gas dissolves in water to form a pale yellow solution.

2. The solution turn blue litmus paper red before it is bleached.

Discussion:

1. The solution formed is acidic and contains bleaching agent.

2. Hypochlorous(I) acid is a strong bleaching agent. It decolourises the colour of litmus paper.

3.

Bromine + Water

Br2 + 2H2O → HBr + HOBr

Observation:

1. Bromine liquid dissolves slowly in water to form a yellowish-brown solution.

2. The solution turn blue litmus paper red before it is bleached slowly.

Discussion:

Hypobromous(I) acid is a weak bleaching agent.

Iodine + Water

I2 + 2H2O → HI + HOI

Observation:

1. Only a little iodine dissolves in water to form a yellowish solution.

2. The solution turn blue litmus paper red but does not bleach it.

Discussion:

Hypoiodous(I) acid has very weak bleaching characteristic.

Notes:

1. Chlorine, bromine and iodine are soluble in water to form an acidic solution.

2. The solubility decreases down the group.

3. Aqueous chlorine and bromine are bleaching agent.

4. Aqueous iodine does not act as bleaching agent.

Halogens React with Sodium HydroxideChlorine + Sodium Hydroxide

Cl2 + 2NaOH → 2NaCl + 2NaOCl + H2O

Observation:

The greenish chlorine gas dissolves quickly in sodium hydroxide, forming a colourless solution.

Discussion:

Chlorine gas react with sodium hydroxide to form salt of sodium chloride, sodium chlorate(I) and water.

Bromine + Sodium Hydroxide

Br2 + 2NaOH → 2NaBr + 2NaOBr + H2O

Observation:

The reddish-brown liquid bromine dissolves in sodium hydroxide, forming a colourless solution.

Discussion:

1. Bromine liquid react with sodium hydroxide to form salt of sodium bromide, sodium bromate(I) and

water.

2. The reaction is less reactive compare with chlorine.

Iodine + Sodium Hydroxide

I2 + 2NaOH → 2NaI + 2NaOI + H2O

Observation:

The black iodine crystals dissolve slowly in sodium hydroxide, forming a colourless solution.

Discussion:

1. The reddish-brown liquid bromine dissolves in sodium hydroxide, forming a colourless solution.

2. Solid iodine react slowly with sodium hydroxide to form salt of sodium iodide, sodium iodate(I) and

water.

3. The reaction is least reactive among the three.

Safety Precaution

1. Fluorine, chlorine and bromine gases are poisonous.

2. Therefore all the experiments involving these gases should be carried out in a fume chamber.

3. The experiments involve fluorine are nor done in school.

4. This is because fluorine is so reactive that it will react with most of the substance it comes into contact

with.

5. It is very difficult to conduct experiments involving fluorine.

Explaining the Reactivity Trend of the Halogens

1. The reactiveness of halogens decreases down the group.

2. This can be explained as below:

When a halogen atom reacts, it gains an electron to form a singly negative charged ion.

As we go down the group from F => Cl => Br => I, the size of the atom increases due to an extra

filled electron shell.

The valence electrons are further and further from the nucleus, the attraction force between the

electrons and the nucleus become weaker and weaker.

Therefore the ability of the atom to attract electron to fill the outermost shell reduces., which means

the reactiveness of the atom reduces.

The Periods1. Period is the horizontal rows of elements in the Periodic Table.

2. The modern periodic table has 7 periods.

3. The period number indicates the number of electron shell.

4. Elements in the same period have same number of electron shells.

5. The proton number of elements increases from left to right crossing the period.

6. The number of electrons is also increases from left to right crossing the period.

Lanthanide and Actinide Series

1. The sixth period has 32 elements. Due to short of space, 14 elements in the transition metal group are

removed from the same horizontal row and is placed below the main table, These elements are called

the Lanthanide Series.

2. The seven period also has 32 elements. With the same reason, 14 elements are removed from the same

horizontal row and is placed below the main table, These elements are called the Actinide Series.

Physical Change Across a Period

There are a few major changes in physical properties across a period. For example,

a. the state of matter changes across the period.b. The size of the molecule changes across the period, too.

State of Matter

1. The diagram above shows that the state of matter of the elements in period 2 and period 3 change from

solid to gas across the period (At room temperature).

2. In period 2, the first 4 elements are solid whereas the last 4 elements are gases. The melting point of

lithium is low while the melting point of boron and carbon is very high.

3. In period 3, the first 6 elements are solid whereas the last 2 elements are gases.

Atomic Size

1. As shown in the diagram above, the atomic size of elements in period 2 and period 3 decreases across

a period.

2. This is because

1. The number of proton in the nucleus increases whereas the number of shells remains

unchanged across a period.

2. As the number of proton in the nucleus increases, the positive charge of the nucleus will also

increases.

3. The negative charge of the valence electrons will also increases due to the increase of

number of valence electrons across a period.

4. Thus, the attraction force between the nucleus and the valence electrons is getting stronger

and stronger across a period.

5. This force will pull the valence electrons closer to the nucleus and thus reduces the atomic

radius.

6. Therefore the size of atom decreases across a period from left to right.

3. Diagram below shows the change of the atomic size of the elements in all seven period. The trends

can be conclude as below:

1. The size of atom increases down the group.

2. The size of atom decreases across a period from left to right.

3. For transition metal, the size of atom does not have obvious change across a period.

Chemical Properties Change Across a Period

1. There is gradual chemical change across a period.

2. The acidity (or basicity) of the oxide of element changes across a period.

3. The metallic properties, the electronegativity of element also change across a period.

Acidic Oxide or Basic OxideFrom left to right across a period the oxides change from alkaline/basic (with metals e.g. Na2O) to acidic (with non-metals e.g. SO2).

Metal, Metalloid and Non-metal

1. As we go across a period from left to right, the elements change from metals to non-metals.

2. There are about 7 elements in the periodic table are classified as semi-metals.

3. The metals in the periodic table are mainly found in the left hand columns (Groups 1 and 2) and in the

central blocks of the transition elements.

4. On the right hand side of the periodic table, there are 7 semi-metals form a staircase like pattern, act as

a divider between metal and non-metal.

5. The semi-metals are also called the metalloid.

Uses of Metalloids1. The most widely used semi-metals are silicon and germanium.2. It is used to make diodes and transistors in electronic industry.

Electronegativity

1. Electronegativity is a measure of the potential of atoms to attract electrons to form negative ions.

2. Metals have low electronegativities.

3. Non-metals have high electronegativities.

4. Electronegativities of the elements increase across a period with increasing proton number. This is

because

1. as the proton number increases, the positive charge of the nucleus will increase accordingly.

2. this will increase the ability of the atom to attract electrons from the surrounding and thus

increase the electronegativity of the atom.

Transition Elements and Their Property

1. Transition metal is a block of metallic elements in between Groups 2 and 13 in the Periodic Table.

2. They are much less reactive than the alkali metals.

3. They do not react as quickly with water or oxygen as alkali metal.

General Physical CharacteristicsAll transition elements are metals. Therefore they have the all the physical properties of metal such as:

high melting point and boiling point

hard,

high density,

high electrical conductivity,

high tensile strength ,

shinny surfaces,

ductile

malleable ,

High Melting Point and Boiling Point

1. The bond between atoms of metal is called metallic bond and usually it is a very strong bond.

2. Thus all the transition metals have high melting points and boiling points.

3. For example: iron melts at 1535°C and boils at 2750°C.

4. Mercury is a transition metal, but with unusual low melting point ( − 39°C).

Special Properties of Transition Elements

Form Coloured Compounds and Ions in Solute

1. Transition elements tend toform coloured compounds either in solid form or dissolved in a solvent.

2. Table below shows the colours of some aqueous solutions of ions of transition elements.

Catalytic Properties

1. Transition elements or their compounds have catalytic properties.

2. A catalyst is a substance that speeds up a reaction but it itself does not change chemically after a

reaction.

3. Many transition metals are used directly as catalysts in industrial chemical processes.

Variable Oxidation State

1. transition metal can have a variable oxidation state, which means it can form more than one ion.

2. For example, iron(II), Fe2+ and iron(III) Fe3+; copper(I), Cu+ and copper (II) Cu2+.

Forming Complex Ions

1. Transition elements can form complex ions.

2. A complex ion is a polyatomic ions (positive or negative) consisting of a central metal ion with other

groups bonded to it.