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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 Characteristics

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

1. high melting point and boiling point
2. hard,
3. high density,
4. high electrical conductivity,
5. high tensile strength ,
6. shinny surfaces,
7. ductile
8. 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.

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 Oxide

From 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 Metalloids

1. 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.

The Periods

1. 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.

Explaining the Reactivity Trend of the Halogens

1. The reactiveness of halogens decreases down the group.
2. This can be explained as below:
1. When a halogen atom reacts, it gains an electron to form a singly negative charged ion.
2. As we go down the group from F => Cl => Br => I, the size of the atom increases due to an extra filled electron shell.
3. The valence electrons are further and further from the nucleus, the attraction force between the electrons and the nucleus become weaker and weaker.
4. Therefore the ability of the atom to attract electron to fill the outermost shell reduces., which means the reactiveness of the atom reduces.

Reactivity of Halogens

Halogens React with Metals

hlorine + 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.

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 Hydroxide

Chlorine + 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.

Group 17 Elements - Halogens

1. 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	Halide
Fluorine	Fluoride
Chlorine	CHloride
Bromine	Bromide
Iodine	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.

Name	Proton Number	Electron arrangement	Colour	Melting point	Boiling point
Fluorine	9	2.7	pale yellow gas	-220ºC	-188ºC
Chlorine	17	2.8.7	Yellowish green gas	-102ºC	-34ºC
Bromine	35	2.8.18.7	dark red liquid, brown vapour	-7ºC	59ºC
Iodine	53	2.8.18.18.7	black solid, purple vapour	114ºC	184ºC
Astatine	85	2.8.18.32.18.7	black solid, dark vapour	302ºC	380ºC

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 Point

1. 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.

      

Explaining the Reactivity Trend of the Alkali Metals

1. 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.

Reactivity of Alkali Metal with Chlorine, Oxygen and Water

 Reaction with Chlorine

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

 Group 1 Metals (2X) + Chlorine Gas (Cl2)→ 

  Metal Chloride (2XCl)

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 

a) lithium -burned slowly with a reddish flame . A white solid is                 produced.

b) sodium -burned brightly with a yellowish flame. A white solid is produced

c) potassium- burned very brightly with a purplish flame . A white solid is produced.

Reaction with Oxygen

1. Group 1 metals react with oxygen gas produces metal oxides. These metal oxides dissolve in water produces alkali.
Group 1 Metals (4X) + Oxygen Gas (O2)→
 Metal Oxide(2X2O)
2. Lithium, sodium and potassium form white oxide powders after reacting with oxygen.
3. The white powder is the oxide of lithium, sodium and potassium.
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 

a) lithium-burns with red flame and produces white powder immediately after reaction.

b) sodium -burned with bright yellow flame, forming white powder immediately after                          reaction.

c) potassium- burned with very bright purplish flame, forming white powder immediately                        after reaction.

Reaction with Water

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

Group 1 Metals + Water → Alkali + Hydrogen gas

2X + 2H2O --> 2XOH + H2 (where X is alkali metal)

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.
4. The reactivity increases down the group from

a) Lithium - moves slowly on the surface of water with 'fizzing' sound.

b) Sodium - moves swiftly on the surface of water with 'fizzing' sound.

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

Solubility of the Oxide Hydroxide and Salt of Alkali Metals

1. 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 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	0.53
Sodium	11	2.8.1	98ºC	883ºC	0.97
Potassium	19	2.8.8.1	63ºC	759ºC	0.86
Rubidium	37	2.8.18.8.1	39ºC	688ºC	1.48
Caesium	55	2.8.18.18.8.1	29ºC	671ºC	1.87
Francium	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:
1. good conductors of heat
2. good conductors of electricity,
3. high boiling points,
4. shinny surface (but rapidly tarnished by air oxidation).
2. Nevertheless, Group 1 metals also show some non-typical metallic properties, such as:
1. low melting points,
2. low density (first three float on water),
3. 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 Metals

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

Li→Li++eNa→Na++eK→K++e
5. They tend to react mainly with non-metals to form ionic compounds.

Safety Precaution

1. Alkali metals are very reactive.
2. Therefore it must be kept in paraffin oil to prevent them from reacting with oxygen and water vapour in the air.
3. We must avoid to hold group 1 metals with bare hand because they may react with water on our hand.

Group 18 Elements - Noble Gases and its uses

1. The "Noble Gases" are the last group in the Periodic Table, they also known as "inert gas", due to their non-reactive behavior ( their arrangement of valence electron in the outermost shell already duplet or octet (full with electron, can't accept or donate electron)
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.

Name	Proton number	Electron arrangement	Melting point	Boiling point
Helium	2	2	-270°C	-269°C
Neon	10	2.8	-249°C	-246°C
Argon	18	2.8.8	-189°C	-186°C
Krypton	36	2.8.18.8	-157°C	-152°C
Xenon	54	2.8.18.18.8	-112°C	-108°C
Radon	86	2.8.18.32.18.8	-71°C	-62°C

Physical Properties

Solubility 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 Elements

Helium

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 Bulb)

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)

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.

(Welding)

Welding

Argon 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

(Laser)

Krypton

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

Xenon

(Camera Flash)

(Fluorescent Bulb)

Also used in fluorescent bulbs, flash bulbs and lasers.

Radon

Used in radiotherapy to kill cancer cell.