1101:Chemistry of Alkali & Alkaline Earth Metals and Halogen Compounds

Chemistry of Alkali & Alkaline Earth Metals and Halogen Compounds

Alkali Metals

Except hydrogen the elements of group 1A of the periodic table form alkali solutions by reacting with water. These elements are called alkali metals such
as Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs) and Francium (Fr).

Example: 2Na +2H2O → 2NaOH + H2  

      Alkali

Of these alkali metals, francium (Fr) is a very rare, radioactive and unstable element. This makes it difficult to study.

Fig. Group IA – the alkali metals

Electron structure and reactivity of alkali metals

The alkali metals have inert gas configuration in the core and ns1 electron in the outermost shell. This means that:-

lithium 2,1	sodium 2,8,1	potassium 2,8,8,1

• They can easily obtain the stable configuration of an inert gas by losing one electron.
• They all lose their outer shell electron in reactions to form positive ions with a +1 charge.
• They have similar physical and chemical properties.

Element	Atomic number	Electronic structure	Valence electron
Li	3	[He] 2s1	2s1
Na	11	[Ne] 3s1	3s1
K	19	[Ar] 4s1	4s1
Rb	37	[Kr] 5s1	5s1
Cs	55	[Xe] 6s1	6s1
Fr	87	[Rn] 7s1	7s1

• The size of each element’s atoms and the number of full electron shells increase down the group.
• This means that down the group the electron in the outer shell gets further away from the nucleus and is shielded by more electron shells.
• The further an electron is from the positive attraction of the nucleus,
  the easier it can be lost in reactions.
• This means that reactivity increases as the size of the atom increases.

Physical properties of alkali metals

Alkali metals are different to typical (transition) metals, such as iron and copper. Unlike typical metals, alkali metals:-

• Are soft and can be cut by a knife – softness increases down the group;
• Have a low density. e.g. lithium, sodium and potassium float on water;
• Have low melting and boiling points.

However, alkali metals do share a few properties with typical metals, such as:

• They are good conductors of heat and electricity;
• They are shiny. This is only seen when they are freshly cut.

1. Nature:

All of them are metals.

2. Colour:

These metals look like white silver.

3. Physical state:

With the exception of Li, alkali metals are extremely soft readily fused. They are highly malleable and ductile.

4. Atomic volume, Atomic and ionic radii: increases from Li to Cs. Due to the addition of an extra shell of electrons to each element

Elements	Li	Na	K	Rb	Cs
Atomic volume (c.c.)	12.97	23.68	45.36	55.80	69.95
Atomic radii (Aº)	1.5	1.86	2.27	2.43	2.62
Ionic radii (Aº)	0.6	0.96	1.33	1.48	1.69

5. Ionization potential:

They have low ionisation potential. The ns1 electron is so weakly held with the nucleus that it can be removed very easily .The ionization potential of
 alkali metals decreases from Li to Cs due to increasing sizes of the atoms.

6. Electronegativity:

Alkali metals have a strong tendency to lose their ns1 electron and hence have electropositive character. Because of this electropositive character, these 
metals have very little electronegative character.

The alkali metals are the most electropositive having very low electronegativity.

Elements	Li	Na	K	Rb	Cs
Electronegativity	1.0	0.9	0.8	0.8	0.7

7. Melting and boiling points:

The melting and boiling point of alkali metals decreases down the group.

Element	Melting point (°C)
Lithium	181
Sodium	98
Potassium	64
Rubidium	39
Caesium	28

Melting points are lower than typical (transition) metals, because of the presence of weak interatomic bonds in the solid state of alkali metals.

8. Density:

The alkali metals generally become denser down the group, due to the large atomic volumes; but the trend is not perfect because potassium is less dense 
than sodium due to an unusual increase in atomic size of potassium.

Element	Density (g/dm3)
Lithium	0.53
Sodium	0.97
Potassium	0.86
Rubidium	1.53
Caesium	1.87

Water has a density of 1 g/dm3. Lithium, sodium and potassium are all less dense than water and so will float.

9. Electropositive character:

The elements which have a strong tendency to lose their outermost electrons to form cations are said to be electropositive.

M → M+ + e-

Thus, the alkali metals have a strong tendency to lose their outermost electrons to from M+ cations and hence show electropositive character. Li shows 
least electropositive character and Cs shows highest electropositive character. So, the electropositive character of alkali metals increases from Li to Cs.

10. Reducing properies:

Reducing agent is a substance which can lose electrons. Since the alkali metals have a strong tendency to lose their outermost electrons, they act as good 
reducing agents.

2M + 2HCl → 2MCl + H2

11. Colouration to the flame:

The alkali metals give colouration to the flame. For example,

Li   → Crimson red; Na → Golden yellow; K   → Violet;   Rb → Lilac red; Cs → Bright blue

This property of alkali metals is used to detect their presence in salt by a test known as flame test.

12. Conductivity:

Alkali metals are good conductors of heat and electricity.

Chemical properties of alkali metals

1. Reaction with H2:

Alkali metals combine directly with hydrogen on gentle heating and form crystalline colourless compound known as metal hydrides. These ionic 
compounds are composed of alkali metal positive ion (M+) and negative hydride ion (H–).

2M + H2 → 2MH

2Na + H2 → 2NaH

These hydrides are easily decomposed by water producing hydrogen.

MH + H2O → MOH + H2

NaH + H2O → NaOH +H2

2. Reactions with oxygen:

All alkali metals react with O2 or air rapidly to form metal oxides. This produces a layer of dull oxide on the surface of the metal and gets tarnished.

The speed with which alkali metals react with air increases down the group:

• lithium – tarnishes slowly;
• sodium – tarnishes quickly;
• potassium – tarnishes very quickly.

Why are alkali metals stored in oil?

The kerosene or paraffin oil prevents them from reacting with air and tarnishing. That’s why these metals are stored under kerosene or paraffin oil.

The reaction between an alkali metal and air is an example of an oxidation reaction:

Lithium + Oxygen → Lithium monoxide

4Li (s) + O2 (g) → 2Li2O (s)

Sodium + Oxygen → Sodium peroxide

2Na (s) + O2 (g) → Na2O2 (s)

Other alkali metals react with O2 to form super oxide.

M + O2 → MO2 (M = K, Rb, Cs)

3. Reaction with water:

All alkali metals react readily with water. The reaction becomes more vigorous down the group, and creates a lot of heat.

2M + 2H2O → 2MOH + H2

This reaction creates alkaline hydroxide ions. This is why the group IA elements are called the alkali metals. The reaction also produces hydrogen gas
 that can be ignited by a lighted splint.

Lithium is the least reactive of the alkali metals. When added to water, it fizzes and moves around slowly across the surface of the water.

2Li (s) +2H2O (l) → 2LiOH (aq) +H2 (g)

When added to water, sodium fizzes more than lithium, and moves quickly across the surface of the water. The sodium melts as it reacts, and it becomes 
spherical and shiny, like a ball bearing. The hydrogen sometimes catches fire because of the heat from the reaction.

2Na (s) +2H2O (l) →2NaOH (aq) +H2 (g)

When added to water, potassium burns with a lilac flame and the hydrogen catches fire immediately. The potassium moves across the surface of the 
water very quickly. Like sodium, it melts with the heat of the reaction.

2K (s) +2H2O (l) →2KOH (aq) + H2 (g)

4. Reaction with halogens:

All alkali metals react violently when heated with halogens to form metal halides. For example,

2Li (s) +Cl2 (g) → 2LiCl (s)

2Na (s) +Cl2 (g) → 2NaCl (s)

5. Reaction with acid:

The alkali metals react vigorously with dilute acid.

2M + 2HCl → 2MCl + H2

Uses of alkali metals

1. Li

• Lithium carbonate is sometimes used to treat mental illnesses such as depression.
• LiHCO3 is used as removing agents in gout.

2. Na & K

Na

1. Na+ is the most abundant cation in extracellular fluids.
2. It is the main contributor to osmotic pressure and hydration; and plays a fundamental part in the electrical activities of the body (e.g., nerve impulse
    transmission and muscular contraction).
3. A decreased concentration of sodium in the blood produces Hyponatremia. If serum sodium levels drop rapidly, seizures may occur.
4. An increased level of sodium in the serum is called Hypernatremia.

K

1. A mineral element that serves as both the principal cation in intracellular fluid and an important electrolyte in extracellular fluid.
2. An abnormally low concentration of potassium in the blood is called Hypokalemia.
3. An excessive amount of potassium in the blood is called hyperkalemia.
4. Along with other electrolytes (e.g. magnesium, calcium, chloride), sodium and potassium participates in many functions, including cell membrane 
   homeostasis, nerve impulse conduction, and muscle contraction.

3.NaCl & KCl

1. Hypertonic solution of NaCl is used to prevent the corneal edema.
2. Isotonic solution of NaCl is used as irrigation fluid.
3. 6% NaCl solution is used as preservative.
4. KCl is used as dirutics.
5. NaCl and KCl are the major components of ORS.
6. Oral potassium chloride is given in attacks accompanied by hypokalemia.
7. Iodized salt: A salt containing 1 part sodium or potassium iodide to 10,000 parts of sodium chloride. Its use prevents goiter due to iodine deficiency
   
8. Lactated R.’s solution: A crystalloid electrolyte sterile solution of specified amounts of calcium chloride, potassium chloride, sodium chloride, and 
   sodium lactate in water for injection. It is used intravenously to replace electrolytes.

4.NaHCO3

1. Sodium bicarbonate is given in metabolic acidosis.
2. It is used in the treatment of tubular acidoses as well as Severe or progressive hyperkalemia.  
3. In the pancreatic juice it neutralizes the acidity of the chyme entering the duodenum from the stomach and prevents irritation of the duodenal mucosa.
   
4. In the renal tubular acidosis sodium bicarbonate is used as Alkali for effective treatments.
5. Sodium bicarbonate is used as uricosuric agent to alkalinize urine and keep uric acid crystals in solution.

NaHCO3  is also used

• As an antacid.
• As a buffering agent in the biological system.
• As a component of ORS.
• In the preparation of effervescent agents.
• In the treatment of CH3OH poisoning.
• As backing powder.

5.Na-CMC
Na-CMC is used as suspending agent.

6.KI

It is used

1. In Lugol’s solution :A strong iodine solution used in iodine therapy.
2. As resorbent to promotes the absorption of abnormal matters, as exudates or blood clots
3. As a component of iodized salt.
4. Orally administered potassium iodide is used to protect the thyroid from radioactive iodine.

Alkaline Earth Metals

Group II A of the long form of the periodic table consists of six elements which are: Breyllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), 
Barium (Ba), and Radium (Ra).

The oxides of the three metals viz., Ca, Sr and Ba were known much earlier than the metals themselves and were called alkaline earths, since they were 
alkaline in character and occurred in nature as earths [lime (CaO), strontia (SrO) and baryta (BaO)]. Later, when Ca, Sr and Ba were discovered, they were 
named alkaline earth metals.

Ca +2H2O → Ca(OH)2 + H2

Fig. Group IIA – the alkaline earth metals

Electronic structure of alkaline earth metals:

The alkaline earth metals have inert gas configuration in the core and ns2 electron in the outermost shell.

Element	Atomic number	Electronic structure	Valence electron
Be	4	[He] 2s2	2s2
Mg	12	[Ne] 3s2	3s2
Ca	20	[Ar] 4s2	4s2
Sr	38	[Kr] 5s2	5s2
Ba	56	[Xe] 6s2	6s2
Ra	88	[Rn] 7s2	7s2

These configurations show that the ultimate shell of these elements has two s-electrons (s2 electrons) while the penultimate shell contains eight (s2p6 electrons
.) When the two s-electrons of the valence-shell are lost, alkaline earth metals are converted into M2+ ions which have inert gas configuration. Since the 
additional electron enters the ns-orbital, these elements are called s-block elements.

The Physical Properties

1. Nature:

These elements all are metals.

2. Physical state:

These elements have a greyish white lustre when freshly cut and are malleable and ductile.

3. Atomic volume, atomic and ionic radii:

There is an increase of atomic and ionic radii as we move from Be to Ra.

Elements	Be	Mg	Ca	Sr	Ba	Ra
Atomic Volume(c.c)	4.90	13.97	25.9	34.54	36.7	38.0
Atomic radii(Aº)	0.96	1.36	1. 74	1.91	1.98	2.20
Ionic radii(Aº)	0.31	0.65	0.99	1.13	1.35	1.52

4. Ionization potential:

The 1st and 2nd ionization potentials of these elements decrease with the increase of atomic radii from Be to Ba. However, both these values for Ra are 
slightly higher than those of Ba.

Elements		Be	Mg	Ca	Sr	Ba	Ra
Ionisation energies (KJ/mole)	I1	899.5	737.7	589.8	549.5	502.9	509.4
	I2	1757.1	1450.7	1145.4	1064.3	965.2	979.06

Ionisation energy of an element is defined as the amount of energy required to remove the most loosely bound electron (i.e. outermost shell electron) from isolated neutral gaseous atom
 in its lowest energy state (i.e. ground state) to convert it into a unipositive gaseous cation.

Ionisation energy is also called ionisation potential, since it represents the amount of potential (or voltage) required to remove the most loosely bound electron from the gaseous atom
 of an element to get a gaseous cation.

In addition to first ionisation potential (I1) defined above, second, third etc. ionisation potentials are also known. We know that the cations are formed by removing the electrons from 
the atom/cation one by one.

M(g)M+(g)M2+(g)Mn+(g)

Thus a given cation has lesser number of electrons than its parent atom or cation. With the decrease of the number of electrons, the magnitude of the effective nuclear charge increases 
from M(g) to Mn+(g). With the increase of effective nuclear charge, the magnitude of the force of attraction between the nucleus and the outer-most shell electron cloud also increases. 
With the increase of force of attraction, the magnitude of energy to remove the outermost shell electron
(i.e. ionisation potential) also goes on increasing from M(g) to M n+((g) and hence I1 < 12 < 13 < ....... . and so on.

5. Density:

Like alkali metals the density of alkaline earth metals also decreases down the order but in an irregular manner due to difference in crystal structure of these 
elements.

Elements	Be	Mg	Ca	Sr	Ba	Ra
Density (gm/cm3)	1.86	1.75	1. 55	2.60	3.59	6.0

Why alkaline earth metals are denser than alkali metals?

Comparing with the alkali metals is denser than alkaline earth metals in the same period. This is due to the fact that , the elements of group II A are packed 
more tightly due to the greater nuclear charge and smaller size.

6. Electropositive character:

The alkaline earth metals show electropositive character which increases from Be to Ba.

M → M2+ + 2e-

Since the ionization energies of these metals are higher than alkali metals, they are not strongly electropositive as the alkali metals.

7. Electronegativity:

These metals have small electronegativity which decreases from Be to Ra.

Elements	Be	Mg	Ca	Sr	Ba	Ra
Electronegativity	1.5	1.20	1.0	1.0	0.90	0.97

8. Melting and boiling point:

The melting and boiling points of alkaline earth metals do not show any regular trend. They are, however, higher than alkali metals. This is because of two 
valence electrons and they are strongly bonded in the solid state than alkali metals.

Elements	Be	Mg	Ca	Sr	Ba	Ra
Melting Point(ºC)	1280	651	851	776	710	960
Boiling Point(ºC)	1500	1100	1487	1366	1537	1150

9. Colouration to the flame:

In flame test, their electrons absorb energy and excited to higher energy levels. When returning to lower energy level, they give out extra energy which 
appears as visible light of characteristic colour as shown below:

Elements	Colour of the flame
Ca	Brick red
Sr	Crimson red
Ba	Apple green

Be and Mg do not give any colour to the flame.

10. Conductivity:

Like alkali metals they are also very good conductor of heat and electricity.

Chemical properties of alkaline earth metals

1. Reducing property:

Alkaline earth metals have two electrons in their valence shell. Due to large size they are comparatively easily oxidised to M2+ ions. They are, therefore, 
strong reducing agents.

The oxidation potentials of alkaline earth metals are lower than those of alkali metals because of their higher sublimation and ionisation energies. Hence alkaline earth metals are 
weaker reducing agents than the alkali metals.

2. Reaction with oxygen:

The alkaline earth metals react readily on heating with O2 to form oxides, MO.

2M + O2 → 2MO

The peroxides of heavier metals (Ca, Sr, Ba etc.) can be obtained on heating the normal oxides with O2 at high temperature.

2BaO + O2 →2BaO2

3. Reaction with hydrogen:

Ca, Sr and Ba react with H2 when heated to produce crystalline metal hydrides. These hydrides react readily with water to give hydrogen.

MH2 + 2H2O → M(OH)2 + 2H2

4. Reaction with nitrogen:

These metals form nitrides on heating in presence of nitrogen which are hydrolysed by H2O to give ammonia.

3M+N2 → M3N2

M3N2+6H2O → 3M(OH)2 + 2NH3

5. Reaction with water:

The alkaline earth metals readily react with water giving off hydrogen and forming metal hydroxides.

M +2H2O → M(OH)2 + H2

Be and Mg do not react readily with water due to their low reactivity. Be does not react even at elevated temperature. Mg reacts only with steam.

6. Reaction with halogen:

The alkaline earth metals directly react with halogen to give metal halides.

M + X2 → MX2,  X = F, Cl, Br, I

Comparison between alkali and alkaline earth metal

Similarities:

The alkali and alkaline earth metals have quite a good deal of resemblance in their properties. The points of similarity are:

1. Like alkali metals, alkaline earth metals are electropositive and have great chemical activity.
2. The electropositive nature continuously increases from beryllium to radium in the second group, as in the first group from lithium to caesium.
3. They do not occur freely in nature.
4. Both are extracted by electrolytic method from their salts such as chlorides.
5. Both are soft and silvery white in untarnished condition and very light.
6. React with water to give hydroxides and hydrogen (gas).
7. Their hydrides produce strong base with water.

NaH + H2O → NaOH + H2

CaH2 + 2H2O → Ca(OH)2 +  2H2

8. They have strong reducing property and forms ionic compounds.
9. They react with water or oxygen with readiness and show much activity towards other substances.

Dissimilarities:

The points of the contrast between the alkali and alkaline earth metals arise mainly due to the presence of 2 valence electrons in the case of alkaline earth 
metals and 1 valence electron in the case of alkali metals.

Sl.	Features	Alkaline Earth Metals (AKM)	Alkali Metals (AM)
01.	Hardness	More harder	Less harder
02.	Softness	Less soft	More soft than AKM
03.	Heaviness	More heavier	Less heavier
04.	Melting & Boiling pints	More than AM	Less than AKM
05.	Density	Greater	Lesser
06.	Basic character	Less basic	More basic
07.	Size	Smaller than AM	Larger than AKM
08.	Reducing property	Lower than AM	Higher than AKM
09.	Oxides and Hydroxides	More covalent in nature	Less covalent in nature
10.	Reaction with water	Ca,Sr & Ba react less vigorously	All AM react more vigorously
11.	Solubility in water	Sulphates, carbonates and phosphates of AKM are insoluble in water	AM are soluble in water except Li2CO3 and Li3PO4 which are insoluble
12.	Polarizing capacity	Greater than AM	Lower than AKM

Uses of alkaline earth metals

Magnesium (Mg):

It is the second most common cation in intracellular fluid.

FUNCTION:

• Magnesium is needed for DNA repair.
• It moderates cellular differentiation and proliferation and improves tissue sensitivity to circulating insulin.
• It is a component of enzymes required for the synthesis of adenosine triphosphate (ATP) and the release of energy from ATP.
• It is also a component of enzymes involved in muscle contraction and protein synthesis.
• Decreased magnesium in the blood is called hypomagnesemia. Clinically, it is accompanied by increased neuromuscular irritability.

1. MgO

MgO is used

• As antacid
• In Mg deficiency
• As an universal antidote
• As an ingredient of zinc phosphate cellulose.

Antacid An agent that neutralizes acidity, esp. in the stomach and duodenum. Examples are aluminium hydroxide and magnesium oxide.

Antidote A substance that neutralizes poisons or their effects.

Zinc phosphate cellulose The oldest of the dental cements, composed of a powder (zinc oxide and magnesium oxide) and a liquid (phosphoric acid and water).

2. Mg(OH)2

Mg(OH)2 is used

• In high dose as a laxative.
• Usually as a non-systemic antacid.
• Prophylactically in stress induced ulcer.
• As soap suspending and emulsifying agent.

3. Magnesium tri-silicate

It is used

• As gastric antacid.
• In drug toxicity.

M/A: Magnesium tri-silicate is made of large particle having hardness. In addition it has laxative properties. Thus it absorbs the toxic particles in the 
stomach and excretes with the stool reducing the toxic effect.  

4. Magnesium tri-silicate hydrated [Mg2Si3O8.2H2O]

It acts as

• Suspending agent without the help of any other compound.
• Thickening agent in the production of chocolate

5. 7-8.5% Mg(OH)2 is called milk of magnesia and is used as laxative.

6. MgSO4

• In the treatment of eclampsia as an intravenous bolus and a continuous infusion of magnesium sulfate are administered.
• A soluble salt of magnesium, is used parenterally in clinical medicine to treat, e.g., hypomagnesemia and torsade de pointes.

7. MgCO3

It is used

• As an antacid and mid laxative.
• As an additive in tooth & face powder.

Calcium (Ca)

Ca is the major component of bone marrow. Bone matrixis made of calcium carbonate, calcium phosphate, and collagen fibres.

FUNCTION:

• Calcium is important for blood clotting, enzyme activation, and acid-base balance;
• It gives firmness and rigidity to bones and teeth; and
• It is essential for lactation, the function of nerves and muscles including heart muscle, and maintenance of membrane permeability.

EXCESS: Hypercalcemia can cause constipation, renal stones, cardiac arrhythmias, cardiac arrest, and depressed brain function (e.g., lethargy or coma). High serum 
calcium levels are usually the result of either hyperparathyroidism or metastatic cancer and may be reduced with hydration, diuresis, corticosteroids, or bisphosphonate 
drugs like pamidronate.

1. Calcium chloride (CaCl2.2H2O): It is used to raise the calcium content of the blood in disorders such as in hypocalcemic tetany, calcium channel 
   blocker, or beta blocker overdose.

2. Calcium hydroxide (Ca(OH)2): It is used as an astringent applied to the skin and mucous membranes.

3. Calcium sulphate (CaSO4): It is used

• As diluent in tablet.
• As plaster cast in the plaster of broken bone.

4. Calcium disodium edetate: It is used to treat poisoning caused by those metals.

5. Calcium gluconate: It is used to treat hypocalcemia, calcium channel blocker, or beta blocker overdose.

6. Calcium glycerophosphate: It is used as a dietary supplement and in formulating drugs.

7. Calcium tungstate: A fluorescent material used for radiological imaging.

Halogens

Group VII A of the periodic table consists of five elements viz. Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At). The elements F, Cl, Br 
and I are collectively called Halogens.

The term halogen is derived from the Greek halos (= salt) and genes (= born) meaning salt producers because most of these elements exist in sea water, 
notably in the form of their sodium compounds.

Astatine is an unstable element of radioactive origin and is not called halogen.

Electronic structure of halogen:

The alkaline earth metals have inert gas configuration in the core and ns2 electron in the outermost shell.

Element	Atomic number	Electronic structure	Valence electron
F	9	[He] 2s2p5	2s22p5
Cl	17	[Ne] 3s23p5	3s23p5
Br	35	[Ar] 3d104s24p5	4s24p5
I	53	[Kr] 4d105s25p5	5s25p5
At	85	[Xe4f14 5d106s26p5	6s26p5

These configurations show that all the halogens have seven electrons (ns2p5) in their outer most shell.

The chemical behaviours of halogens, therefore, show a tendency to achieve the inert gas configuration by accepting an electron from other atoms either by 
transfer or by sharing.

At the ordinary temperatures the halogens exist as diatomic molecules, F2, Cl2, Br2 and I2 for the same reason, because diatomic halogens have achieved stability by sharing an electron
 from each atom in an electron pair bond (covalent molecules) :

The Physical Properties

Some physical properties are discussed below

1. Occurrence.

These elements do not occur in the free state. But in the combined state they occur fairly abundantly in nature.

2. Colour.

All the halogens are coloured. Their colour is due to the fact that their molecules absorb light in the visible region by means of which the outer electrons are 
excited to higher energy levels.

Fluorine →pale yellow	Bromine →reddish brown
Chlorine →greenish yellow	Iodine →deep violet

3. Odour. All the halogens have very pungent and unpleasant odours. If inhaled even in very small quantities, they cause inflammation of the mucous 
membranes of the throat and nose and serious irritation of respiratory duets. Large quantities may have fatal effect.

4. Physical State and Density:  

There is a change of state from gas to solid and hence density increases as we move from F to I.

Here both the number of the electron and size of atom increases from top to bottom. But the number of electron increases more than the size thereby increases the attractive forces and 
ultimately the density from top to bottom increases.

Elements	F	Cl	Br	I
Physical State	Gas	Gas	Liquid	Solid
Density (gm/c.c.)	1.3	1.9	3.4	4.93

5. Atomic and Ionic Radii: These parameters increase as we descend the group form F to I. This is due to the fact that as the number of shells 
increases the distantness between the outermost shell and nucleus also increases resulting in an increase in the radii.

Elements	F	Cl	Br	I
Covalent radii (Aº)	0.72	1.00	1.14	1.35
Ionic radii (Aº)	1.36	1.81	1.16	2.16

6. Electronegativity.

Halogens have large values of electronegativity. These values decrease as we proceed from F to I in the group. Large electronegativity values of halogen 
atoms indicate that X atoms have a strong tendency to form X– ions.

(Here, from top to bottom with the increase of an extra electron shell, the size of the atom increases and the attractive forces between the electron and nucleus decreases.)

Elements	F	Cl	Br	I
Electronegativity	4	3	2.8	2.5

7. Melting and Boiling Points.

The melting and boiling points of the halogens regularly increase form F to I. This happens as the attractive forces (van der Waals forces) between molecules
 progressively increases as the molecules increase in size and mass.

Elements	F	Cl	Br	I
Melting point (ºC)	- 223	-102	- 7	114
Boiling Point (ºC)	-188	- 34	59	185

8. Ionization Potentials.

The ionisation potentials of halogens are very high. As they have tendency to accept electron to become negative ion so they have a little tendency to lose 
an electron. These values progressively decrease from F to I, as the atomic size of these elements increases in the same order.

Elements	F	Cl	Br	I
Ionization Potentials (ev)	17.42	13.01	11.84	10.45

9. Electron Affinity.

Electron affinity values decrease from Cl to I.

Elements	F	Cl	Br	I
Electron Affinity (ev)	3.70	4.00	3.80	3.40

Why the electron affinity value of F is less than that of Cl?

Electron affinity is the amount of energy released when an electron is added to a neutral isolated gaseous atom. The atoms with high electron affinity will 
form the anions quite easily.

The F has smaller size and the electrons are more densely packed to each other. When other electrons come, it shows repulsion which gives it a lower EA 
value than Cl.

10. Solubility.

The halogens are slightly soluble in water and their solubility decreases from Cl to I.
F cannot be dissolved in H2O, since it decomposes H2O liberating O2 and some O3. The halogens dissolve much better than H2Oin many organic solvents
 like carbon disulphide, chloroform, ether etc.

Chemical properties of Halogens

1. Reactivity. All the halogens are the most reactive elements as a family.

The main reasons for their extreme reactivity are: (a) Dissociation energies required converting the diatomic halogen into mono atomic and (b) electron affinity required to be an anion.
 Halogen molecules(X2) have low dissociation energies values and hence it is much easier for these molecules to dissociate into X atoms and react with other substances more. Because
 of high electron affinity values of halogen atoms, halogens have a strong tendency to add on electron to form X– ion. The reactivity of halogen atoms decreases as we move down the 
group from F to I.

2. Oxidising Property.

The halogens, due to high electron affinity values, have a great tendency to accept an electron and hence act as strong oxidising agent. The oxidising 
property of a halogen molecule, X2 is represented by

-1/2 X2 (s, l or g) + e —+ aq →X— (aq)

3. Reaction with water:

Fluorine decomposes H2O vigorously even in dark, chlorine decomposes H2O in sunlight, bromine also decomposes H2Oin sunlight but very slowly while 
iodine does not decompose water at all.

2F2 + 2 H2O→4HF + O2	I2 + H2O →No reaction
Cl2 + H2O →HClO + HCl	Br2 + H2O →HBrO + HBr

4. Formation of Hydracids.

All the elements of this group combine with hydrogen and form hydrogen halides (HX) which are called hydracids or halogen acids.

The affinity of these elements for hydrogen decreases from fluorine to iodine. Thus F2 combines with H2 explosively and even in the dark, Cl2reacts with H2 
 in diffused sunlight, Br2, reacts with H2 only on heating above 500ºC while I2 reacts with H2 only on heating and in presence of Pt as a catalyst.

X2 + H2→2HX

5.  Reaction with oxygen: Halogens form a number of oxides by indirect reactions and arc.

Halogens do not combine directly with O2. However, binary halogen-oxygen compounds can be prepared by indirect methods. Examples of such compounds
 are OF2, O2F2, Cl2O, ClO2, Br2O, BrO2, I2O4, I2O5 etc.

6. Formation of Oxy-acids.

Excepting F2, all other halogens form oxy.-acids. These are mainly of four types
viz. HXO (e.g. HClO, HBrO, HIO); HXO2 (e.g. HClO2); HXO3 (e.g. HClO3, HBrO3, HIO3) and HXO4 (e.g. HClO4, HBrO4, HIO4).

7. Reactions with metals:

Almost all metals react with halogens at various conditions and with different rates. They normally form halides of the metals with high oxidations state.

nX2 + 2M→2MXn

8. Formation of Interhalogen Compounds.

Two different halogen atoms, due to the difference in their electronegativity values, combine with one another to form a number of covalent compounds 
which are called interhalogen compounds. These are of four types namely AB (e.g. ClF, BrF etc); AB3 (e.g. ClF3, BrF3, ICl3); AB5(e.g. BrF5 , IF5 ) and
AB7 (e.g. IF7 ) where A and B are two different halogen atoms.

9. Reactions with organic compounds:

Hydrocarbons burn in Cl2 gas to form HX and free carbon. Unsaturated organic compounds give important reaction’s with halogens.

Uses of halogens

1. Chloride

FUNCTION:

Chloride is the major extracellular anion and contributes to many body functions including

• The maintenance of osmotic pressure,
• Acid base balance,
• Muscular activity, and
• The movement of water between fluid compartments.
• It is associated with sodium in the blood.

2. Chlorine

• Chlorine is an active bleaching agent and germicide, owing to its oxidizing powers.
• It is used extensively used to disinfect water supplies and treat sewage.
• The sodium and chlorine ions are important in maintaining the proper electrolyte balance in body fluids.

3. Bromide

• It is a central nervous system depressant, and over dosage can cause serious mental disturbance.

4. Iodine

FUNCTION:

• Iodine is part of the hormones triiodothyronine (T3) and thyroxine (T4), and prevents goiter by enabling the thyroid gland to function normally.
• Iodine is a commonly used antiseptic.
• It is used in the Gram staining.

5. Radioactive iodine (131I)

It is used

• To treat hyperthyroidism.  
• In diagnosis and treatment of thyroid disorders and Graves’ disease.
• In the treatment of toxic goiter and thyroid carcinoma.

DEFICIENCY SYMPTOMS: Iodine deficiency in the diet may lead to simple goiter characterized by thyroid enlargement and hypothyroidism. In young 
children, this deficiency may result in retardation of physical, sexual, and mental development, a condition called cretinism.

6. Potassium iodide (KI)

• This form of potassium is recommended for use following exposure to radioactive iodides downwind from a nuclear reactor accident.

7. Povidone-iodine

• A complex of iodine with povidone is used in dilute concentration as a surgical scrub, in aerosol spray, in vaginal douche solutions, and in ointments 
  and gels.

8. Iodized salt

• Its use prevents goiter due to iodine deficiency.

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