Chemistry Podcast

Sunday, 20 February 2022

10 Class- Acids, Bases and Salts

Acids Bases and Salts

·       Natural indicator of acids and bases – Litmus, Turmeric
·       Synthetic indicator of acids and bases – Methyl Orange, Phenolphthalein
·       Olfactory indicators of acids and bases - Odor of these substance changes in acids and bases.
pH scale for acid neutral and alkaline (or base) - Acid base strength indication by H and OH ion concentration
pH scale for Acid Neutral and Alkaline (or Base) & Acid Base strength indication by H and OH ion concentration

What is Acid?

Acids are the chemicals which can donate a proton or accept an electron pair in chemical reactions having pH less than 7 and changes color of blue Litmus to red.
Example- HCl, H2SO4

What is Base?

Bases are the chemicals which can accept a proton or donate an electron pair in chemical reactions having pH more than 7 and changes color of red Litmus to blue.
Bases which dissolve in water are called as Alkali
Example- NaOH

What is Salt?

Salt is a mineral or any neutral compound containing cations (+ve ions) attached with anions (-ve ions). 
Example- NaCl

Acids and Bases Reaction with Metals:-

                   Metal react with acid to form salt.
Acid + Metal ------> Salt + Hydrogen Gas
Metal react with base to form salt.
Base + Metal -------> Salt + Hydrogen Gas

Reaction of Metal Carbonates and Metal Hydrogen-Carbonates with Acids:-

Metal carbonates and metal hydrogen carbonates react with acids to form salt, water and carbon dioxide.
Metal Carbonate + Acid ----> Salt + Water + Carbon Dioxide
Example- Na2CO3 + 2HCl ---> 2NaCl + H2O + CO2
Metal Hydrogen Carbonate + Acid --> Salt + Water + Carbon Dioxide
Example- NaHCO3 + HCl ---> NaCl + H2O + CO2

Reaction of Acids with Bases:-

Reaction of acids with bases to form salt and water is called neutralization reaction.
Acid + Base ----> Salt + Water

Reaction of Acids with Metallic Oxides:-

Acids react with Metallic Oxides to form salt and water.
Acid + Metallic Oxide ----> Salt + Water

Reaction of Bases with Non-Metallic Oxides:-

Non-Metallic Oxides are acidic in nature so these react with bases to form salt and water.
Base + Non-Metallic Oxide ----> Salt + Water

Acids or Bases in Water:-

When acids dissolve in water they produce Hydrogen Ion H+(Aq) or Hydronium Ion (H3O+)
HCl + H2O ----> H3O+ + Cl-
H+ + H2O ---> H3O+
When bases dissolve in water they produce Hydroxide Ions (OH-)
NaOH + H2O -------> Na+(Aq) + OH-(Aq)

Reactions of Acids or Bases with Water are highly exothermic. Process of mixing Acid or Base with water decrease concentration of ions per unit volume, this process is known as dilution. 

Strength of Acids or Bases:-

Strength of acids depends on number of hydrogen ions (H+) produced and strength of bases depends on number of hydroxide ions (OH-) produced. A universal indicator (present on pH paper) is used to find strength of acids or bases.

pH Scale:-

It is a scale to measure hydrogen ion concentration in solution.
pH also known by seeing the color change in litmus paper.
Meaning of ‘p’ in ‘pH’ is “potenz”, which is a German word whose meaning is “power”.
pH scale measure pH from 0 (very acidic) to 14 (very alkaline).
pH of Neutral Solution is 7
pH of Acidic Solution is less than 7
pH of Basic Solution is more than 7

How to Measure PH of Acids and Bases ?

You can measure pH of acids and bases with the help of litmus paper, pH paper and pH meter. To know pH of any substance watch below pH experiment conducted by the Chemistry Notes Info. 

What is acid rain?

If pH of rain water is below 5.6 on pH scale than, that rain is called as acid rain.

Some Naturally Occurring Acids:-

Natural Source
Acetic Acid
Citric Acid
Tartaric Acid
Oxalic Acid
Sour Milk (Curd)
Lactic Acid
Citric Acid
Ant Sting
Methanoic Acid
Nettle Sting
Methanoic Acid

Common Salt:-

Common salt is very important raw material for production of other daily use material.

Sodium Hydroxide:-

Sodium Hydroxide is obtained by passing electricity through aqueous solution of sodium chloride (brine). Process is known as Chlor-Alkanization.   
2NaCl (aq) + 2H2O (l) ----> 2NaOH (aq) + Cl2 (g) + H2 (g)

Bleaching Powder (CaOCl2):-

Bleaching Powder is obtained by reaction between chlorine (Cl2) and dry slaked lime [Ca(OH)2].
Ca(OH)2 + Cl2 -----> CaOCl2 + H2O

Use of Bleaching Powder:-

·       For bleaching wood pulp in paper industry.
·       For bleaching washed clothes in laundry.
·       For bleaching cotton and linen in textile industry.
·       Used as oxidizing agent in chemical industry.
·       Used as disinfectant for drinking water to kill germs.

Baking Soda:-

Chemical name of baking soda is sodium hydrogen carbonate (NaHCO3). It is obtained by reaction between Sodium Chloride (NaCl), Water (H2O), Carbon Dioxide (CO2) and Ammonia (NH3).
NaCl + H2O + CO2 + NH3 ----> NH4Cl (Ammonium Chloride) + NaHCO3 (Sodium Hydrogen Carbonate)

Uses of Sodium Hydrogen Carbonate:-

·       To produce baking powder {mixture of baking soda + mild edible acid (like tartaric acid)}
·       Used to produce antacids (neutralize access acid in stomach to provide relief from acidity)
·       Used in soda acid fire extinguishers

Washing Soda:-

Washing soda is obtained by heating baking salt (Sodium Hydrogen Carbonate) and recrystallization of Sodium Carbonate produced above.
2NaHCO3 (Sodium Hydrogen Carbonate) + Heat --------> Na2CO3 (Sodium Carbonate) + H2O + CO2
Na2CO3 + 10H2O --------> Na2CO3.10H2O

Uses of Washing Soda:-

·       Used in glass, soap and paper industries
·       Used in preparation of sodium compounds like borax
·       Used as cleaning agent
·       Used for removal of permanent hardness of water

Crystal of Salt:-
Presence of fixed no. of water molecules in one formula unit of salt is called as water of crystallization.
·       Copper sulphate crystals with water molecule (CuSO4.5H2O) are blue in color, while
·        Copper sulphate crystals without water molecule (CuSO4) are white in color.

Plaster of Paris:-

Plaster of Paris is obtained by heating gypsum (CaSO4.2H2O) at 373K.
CaSO4.2H2O (Gypsum) + Heat ----> CaSO4.1/2H2O (Plaster of Paris) + 3/2H2O

Uses of Plaster of Paris:-

·       Used by orthopedic doctors for supporting fractured bones.

Tuesday, 15 February 2022

S Block Elements Class 11

S Block Elements Class 11

s block elements in periodic table

Elements in which valence or outermost electron enters in the S–orbital are called S – block elements.
The S – block elements are placed to the left hand side of long form of the periodic table. S – Orbital can accommodate a maximum of two electrons. The S–block contains two groups (vertical columns) in periodic table, namely group–1 and group-2. Elements of group -1 are called alkali metals and those of group -2 are called alkaline earth metals.



S – Block elements do not occur in nature in free state. Among the alkali metals sodium and potassium are abundant and lithium, rubidium and cesium have lower abundance. Among the alkaline earth metals magnesium and calcium are abundant in the earth crust. Radium is rarest of all.



The general electronic configuration of group – 1 element is ns1. All alkali metals have one valance electron hence form monovalent M+ ions, and are highly reactive. Group – 1 consists of following elements.

Hydrogen (H), Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (S) and Francium (Fr).


The electronic configuration of group – 1 elements is as follows –



Atomic Number

Electronic configuration in condensed form








[He] 2s1




[Ne] 3s1




[Ar] 4s1




[Kr] 5s1




[Xe] 6s1




[Rn] 7s1



Amongst alkali metals, lithium has smallest size and highest polarizing power (charge / radius ratio) due to this reason lithium behaves different from other elements of group-1 (alkali metals). It shows following anomalous properties.

    ·         Lithium is hard metal, its melting point and boiling point are highest among the group.

    ·         Lithium is least reactive among all alkali metals and it is the strongest reducing agent.

    ·         Lithium hydride is the stable amongst all the alkali metal hydrides.

    ·         LiOH is weak base while hydroxides of other alkali metals behave as strong bases.

    ·         Lithium nitrate when heated gives lithium monoxide(Li2O). While other alkali metal nitrates decompose to give corresponding nitrites.

    ·         Lithium combines with ammonia to form lithium amide (Li2NH) while other alkali metals form amides of the general formula MNH2 (where M = Na, K, Rb, Cs).



The element lithium of second row shows similar properties with its diagonally opposite member magnesium of third row. This is known as diagonal relationship. The diagonal relationship is due to the similarity in ionic sizes and polarizing power (i.e. charge/radius ratio) of lithium and magnesium. Also electronegativity of Li (1.00) and Mg (1.20) are not much different.

s block elements class 11 chemistry notes


    1.    Atomic and ionic radius of S Block Elements

Alkali metals have the largest atomic and ionic radius in their respective periods. On moving down the group, the size goes on increasing due to the presence of extra shells hence the atomic and ionic radius also increases.

    2.      Ionization enthalpy or energy of S Block Elements

Ionization enthalpies of alkali metals are low and decrease down the group from Li to Cs. This is due to increase in atomic number, size, nuclear charge and increase in screening effect.


    3.      Electronegativity of S Block Elements

The electronegativity of the alkali metals is very low due to their electropositive character. The electronegativity decreases from Li to Cs as the electropositive character increases.


    4.      Melting and boiling point of S Block Elements

Alkali metals have very low melting and boiling points due to presence of weak intermetallic bonds. It decreases on moving down the group.


    5.      Oxidation States of S Block Elements

Due to the presence of only one electron in the valence shell, they exhibit only +1 oxidation state.



The alkali metals are highly reactive due to their large size and low ionization enthalpy.

    1.      Reactions of S-block elements with oxygen (air)

Alkali metals burns vigorously in oxygen to form oxides. Lithium forms monoxide (Li2O) and sodium forms peroxide, (Na2O2).

4Li   +   O2       2Li2O      (Oxide)

2Na   +   O2      Na2O2      (Peroxide)

The other elements form super-oxides.

M   +   O2       MO2    (Superoxide)

   (where, M = K, Rb, Cs)


    2.      Reaction of S-block elements with water

The alkali metals react with water to form corresponding hydroxides and evolve hydrogen gas (Dihydrogen). (M = Li, Na, K, Rb, Cs)

2M   +   2H2O      2MOH   +   H2


   3. Reaction of S-block elements with hydrogen (Dihydrogen)

The alkali metals react with dry hydrogen at high temperature (673K) to form corresponding hydrides.

2M   +   H2      2MH


    4.      S-block elements Reaction with halogens

All the alkali metals react vigorously with halogen to form their respective ionic crystalline halides with general formula M+X- where M = Na, K, Rb, Cs and X = Cl, Br, I and F.

2M   +   X2     Δ    2M+X-





1] Sodium carbonate (washing soda, Na2CO3)

 Sodium carbonate is usually prepared by a process known as the ammonia–soda process or Solvay process. Raw materials are NaCl, NH3 and limestone (for CO2).

The process involves the following steps

Step – I

CO2 gas bubbled through a solution of NH3 to form NH4HCO3 i.e. ammonium hydrogen carbonate.

NH3   +  H2O   +  CO2      NH4HCO3

Step – II

Sodium hydrogen carbonate (NaHCO3) precipitates out because of the common ion effect caused due to the presence of excess NaCl.

NH4HCO3   +   NaCl        NaHCO3    +    NH4Cl

Step III

The precipitate NaHCO3 is then filtered of and ignited to get sodium carbonate (Na2CO3)

2NaHCO3      Δ     Na2CO3   + CO2    +   H2O


Properties of Na2CO3

Sodium carbonate crystallizes from water as a decahydrate (sodium carbonate decahydrate, Na2CO3.10H2O, also known as washing soda).Above 373K, the monohydrate from becomes completely anhydrous and change to a white powder called soda ash.

Na2CO3.H2O     above 373K   Na2CO3   +   H2O



Uses of Na2CO3

                    i.            It is used in the manufacture of glass soap, borax and caustic soda.

                  ii.            It is used in water softening in laundry and cleaning.

                iii.            It is used in paper, paints and textile industries.



2] Sodium hydroxide (caustic soda, NaOH)

Commercially sodium hydroxide is prepared by the electrolysis of sodium chloride in Costner – Kellner cell. It is also prepared by adding calcium hydroxide to the solution of sodium carbonate.

Na2CO3   +    Ca(OH)2       2NaOH    +    CaCO3 


Properties of NaOH

Sodium hydroxide is a deliquescent, white crystalline solid. It is readily soluble in water and forms a strong alkaline solution.

The solution of sodium hydroxide at the surface reacts with CO2 in the atmosphere to form Na2CO3.

2NaOH   +     CO2       Na2CO3   +   H2O


Uses of NaOH

            i.            It is useful in purification of bauxite and to manufacture of soap, paper, artificial silk and a number of chemicals.

                  ii.            It is useful in petroleum refining and textile industries.

                iii.            It is important as laboratory reagent and in preparation of pure fats and oils.



3] Sodium Chloride (NaCl)

It is also known as common salt or table salt. It is mainly obtained by the evaporation of sea water. Crude sodium chloride thus obtained contains CaSO4, CaCl2, MgCl2 as impurities and is deliquescent. Pure NaCl is obtained by passing HCl gas through saturated solution of crude NaCl. Due to common ion effect NaCl gets precipitated.


Properties of Sodium Chloride (NaCl)

It is white crystalline solid. It has melting point 1081K and boiling point 1713K. It has solubility of 36gram per 100gram of water at 273K. NaCl when heated with concentrated H2SO4 and MnO2 it liberates chlorine gas.

2NaCl  +  MnO2  +  2H2SO4    MnSO4  +  Na2SO4  +  2H2O   +   Cl2 


Uses of Sodium Chloride (NaCl)

                i.            It is used as a common salt or table salt for domestic purpose and preservative for meat, fish etc.

                          ii.            It is used for the preparation of Na2O2, NaOH, Na2CO3 etc.



4] Sodium hydrogen carbonate (NaHCO3)

It is also known as sodium bicarbonate and baking soda. It is made by saturating a solution of sodium carbonate with CO2.

Na2CO3   +   H2O    +    CO2        2NaHCO3

Properties Sodium hydrogen carbonate (NaHCO3)

It is white crystalline solid, sparingly soluble in water. On heating, it gives Na2CO3 and CO2.

2NaHCO3     373K     Na2CO3   +   CO2   +   H2O

It’s aqueous solution is alkaline due to hydrolysis.

NaHCO3   +   H2O      NaOH   +   H2CO3


Uses Sodium hydrogen carbonate (NaHCO3)

                    i.            It is used in fire extinguisher and for baking of cakes breads etc.

                    ii.            It is also used as medicine to minimize the acidity of stomach i.e. antacid.




Sodium and potassium are very important and play a vital role in biological system. Our bodies obtain them from fruits and vegetable. Common salt is the most important source of sodium in the diet. Their ions maintain the sensitivity of nerves and control muscles. 

Deficiency of sodium shows reduction in fat deposit, atrophy, lung infection retarted bone growth, reduces blood pressure etc. Deficiency of potassium reduces heart beats. Hypertrophy of kidneys and paralysis of muscle.





The general electronic configuration of group–2 elements is ns2. All alkaline earth metals have two valance electrons, hence form divalent M2+ ions, and are highly reactive like alkaline metals. Group – 2 consists of following elements.

Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), Radium (Ra).


The electronic configuration of group – 2 elements is an follows –



Atomic Number

Electronic Configuration




[He] 2s2




[Ne] 3s2




[Ar] 4s2




[Kr] 5s2




[Xe] 6s2




[Rn] 7s2




Beryllium, the first member of the group – 2 metals, shows anomalous properties as compared to magnesium and rest of the members of group-2. It is due to its small size and highest polarizing power. It shows following anomalous properties.

    ·         Beryllium is hard metal while other alkaline earth metals are soft.

    ·         The melting, boiling points and ionization enthalpy of Be are highest of all the alkaline earth metals.

    ·         Be forms covalent compounds while other members form ionic compound.

    ·         Be does not liberate H2 from acids while other metals liberate H2.




The element beryllium is the first member of group – 2 which shows similar properties with its diagonals opposite element aluminum of group – 13. This diagonal relationship is due to similarity in ionic sizes and polarizing power.

S Block Elements Notes


    1.      Atomic and ionic radius

The atomic and ionic radii of the alkaline earth elements are smaller than those of the corresponding alkaline metals of the same periods. The atomic and ionic radius increases down the group with increase in atomic number.


    2.      Ionization enthalpy or energy

Ionization enthalpies of alkaline earth metals are higher than those of alkali metals. Ionization enthalpy decreases down the group.


    3.      Electronegativity

The electronegativity of group–2 elements decreases down the group. Their electronegativity values are higher than those of alkali metals.


    4.      Melting and boiling point

The melting and boiling points are not regular in the group, mainly due to different crystal structure of the metals. However, these are higher than those of alkali metals.


    5.      Oxidation states

Due to the presence of only two electrons in the valence shell, they exhibit only +2 oxidation state.



Due to low ionization enthalpy, alkaline earth metals are fairly reactive. However, their chemical reactivity is lower than those of alkali metals.

    1.      Reaction with oxygen

With pure oxygen, Be, Mg and Ca form oxides whereas Ba, Sr and Ra form peroxides.

2M   +    O2        Δ       2MO     (M = Be, Mg or Ca)

                                Metal oxide

M   +    O2        Δ       MO2     (M = Ba, Sr or Ra)

                                Metal peroxide

The affinity of metals towards oxygen increases down the group.


    2.      Reaction with water

Alkaline earth metals react with water to form metal hydroxides and evolve hydrogen gas. The chemical reactivity with water increases from Mg to Ba.

M   +   2H2O          M(OH)2    +    H2       (M = Mg, Ca, Sr or Ba)


    3.      Reaction with hydrogen

Except Be all the alkaline earth metals react with hydrogen on heating to form metal hydride of the general formula, MH2.

M   +    H2        Δ       MH2     (M = Mg, Ca, Sr or Ba)


    4.      Reaction with halogens

All the alkaline earth metals react with halogens at high temperature to form their halides.

M   +    X2          MX2     (X = F, Cl, Br, I)





1] Calcium oxide or quick lime (CaO)

It is prepared commercially by heating limestone (CaCO3) in a reverberatory kiln at 1070 to 1270K.

CaCO3            Δ           CaO    +     CO2


Properties Calcium oxide or quick lime (CaO)

Calcium oxide is a white amorphous solid. It has a melting point of 2870K. When it is exposed to atmosphere, it absorbs moisture and carbon dioxide.

CaO   +    H2O        Ca(OH)2

CaO   +   CO2          CaCO3


Uses Calcium oxide or quick lime (CaO)

                    i.            It is used in the manufacture of sodium carbonate from caustic soda.

                  ii.            It is an important primary material for large scale production of cement.



2] Calcium carbonate (CaCO3)

Calcium carbonate occurs in nature in several forms like chalk, lime stone, marble state, calcite etc. It is prepared by passing carbon dioxide through slaked lime.

Ca(OH)2    +    CO2        CaCO3    +   H2O

It can also be prepared by adding sodium carbonate to calcium chloride.

CaCl2    +    Na2CO3         CaCO3    +     2NaCl


Properties Calcium carbonate (CaCO3)

It is white fluffy powder. It is almost insoluble in water. When CaCO3 is heated to 1200K, it decomposes to evolve CO2.

CaCO3          1200K         CaO     +    CO2

It reacts with dilute acid to liberate carbon dioxide.

CaCO3    +     H2SO4         CaSO4    +    H2O    +    CO2


Uses Calcium carbonate (CaCO3)

                    i.            It is used in the manufacture of quick lime and in building material in the form of marble.

                  ii.            It is used as antacid, mild abrasive of tooth paste, constituent of chewing gum.



3] Calcium hydroxide or slaked lime (Ca(OH)2)

It is prepared by adding water to quick lime.

CaO   +   2H2O        Ca(OH)2    +     H2 

It is also prepared by treating calcium chloride with caustic soda.

CaCl2   +    2NaOH        Ca(OH)2    +    2NaCl


Properties Calcium carbonate (CaCO3)

It is white amorphous powder, sparingly soluble in water and the aqueous solution is called lime water. When CO2 is passed through lime water, it turns milky

Ca(OH)2   +   CO2        CaCO3    +    H2O


The milkiness disappears on passing CO2 gas in excess.

CaCO3   +   CO2    +    H2O        Ca(HCO3)2   

                                        Soluble milkiness disappears


Uses Calcium carbonate (CaCO3)

                    i.            It is used to prepare mortar which is an important building material.

                  ii.            It is used glass and tanning industry also to prepare bleaching powder and for purification of sugar.



Hope you enjoy learning s block elements in periodic table. You can also learn with other chemistry notes chapter on other elements in periodic table from given below topics-


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