CHAPTER → 1

CHEMICAL REACTIONS AND EQUATIONS

*  CHEMICAL EQUATIONS

CHEMICAL REACTIONS:- The process in which two or more substance combine with each other to form new substances with new properties is called chemical reaction.

There are two parts of a chemical reactions :-

(i) Reactants:-   The substances which take part in a chemical reaction are known as reactants.

(ii) Products:-  The new substances formed during a chemical reaction are known as products.

There are 5 ways to tell if a chemical reaction has occurred.

• Change in state.
• Change  in colour.
• Change in temperature.
• Evolution of a gas.
• Formation of precipitate.

  Chemical reaction in everyday life:-

• Digestion of food.
• Respiration.
• Rusting of iron.
• Formation of curd.
• Burning of magnesium ribbon.

Chemical Equations:-  A chemical equation is a written representation of a chemical reaction.

The representation of chemical reaction using symbols and formulae of the substances is called chemical equation.

A   +   B       →      C   +     D

Reactants                   Products   

n this equation, A and B are called reactants and C and D are called the products. The arrow shows the direction of the chemical reaction. The necessary condition such as temperature, pressure or any catalyst should be written on arrow between reactants and products.

E.g. Magnesium is burnt in air to form magnesium oxide.

(i) Word equation for above reaction would be -

 Magnesium + oxygen     →           Magnesium oxide

   ( Reactants )                                      ( Product )

Skeletal equation for above reaction would be - 

Mg +    O₂        →           MgO

BALANCING CHEMICAL EQUATIONS:-

• LAW OF CONSERVATION OF MASS :-  Mass can neither be created nor be destroyed in a chemical reaction.
• So number of elements involved in chemical reaction should remain same at reactant and products side.

For Example ,

Zn   +   H₂SO₄             →     ZnSO₄  +         H₂

(Zinc)    ( Sulphuric Acid)        (Zinc Sulphate)   ( hydrogen) 

Let us check the number of atoms of different elements on both sides of the arrow .

 

As the number of atoms of each element is same on both sides of arrow. This is   a balanced chemical equation.

 Let us take another example :-   

Fe     +    H₂O   →  Fe₃O₄   +   H₂

STEP 1 :-   Write a chemical equation.

Fe   +    H₂O  →   Fe₃O₄  +  H₂

 STEP 2:-  List the number of atoms of different elements present in the unbalanced equation.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

STEP 3 :_ Select the element which has the maximum number of atoms . now equalize the number of atoms by putting coefficient in front of it.

Fe   +      4  H₂O      →     Fe₃O₄   +     H₂

STEP 4 :-  Fe and H atoms are still not balanced choose any elements now to balance. To equalize the number of H atoms,

Fe   +     4 H₂O   →      Fe₃O₄   +     4 H₂

STEP 5 :- Now, take Fe and equalize the number of Fe atoms.

3 Fe   + 4 H₂O     →Fe₃O ₄     +    4 H₂

Now all the atoms of elements are equal on both sides.

STEP 6 :-  To make the chemical equation more information ,write the physical states of reactants and products.

Solid state = (s)

Liquid state = (l)

Gaseous State = (g)

Aqueous state = (aq)

3 Fe (S)    +   4H₂O(g)       →   Fe₃O₄S+  4H₂(g) 

STEP 7:-  Write necessary conditions of temperature pressure or catalyst on above or below arrow.

For Example:-          

 

TYPE OF CHEMICAL REACTION

There are five type of chemical reactions.

COMBINATION REACTION:-   The reaction in which two or more reaction combine to form a single product is called combination reaction They are represented by equation of the following term. 

 A  + B         →          AB      

Reactant           product

For Example :-  

Burning of coal  

 C(s)      +      o₂(g)    →      Co₂ (g)   

(carbon)     (oxygen)     (carbon dioxide)

Formation of water

2  H₂ (g)    +        O₂ (g)    →    2H₂ O (l)           

 (Hydrogen)      (oxygen)          (Water)   

Formation of slaked line

 CaO (s)     +      H₂ 0(L)     →       Ca(OH)₂ (aq)

(calcium oxide)   (water)   (calcium Hydroxide)/ slaked lime

Formation of slaked lime by the reaction of calcium oxide with water.

DECOMPOSITION REACTION :-   The reaction I which single compound breaks down to simpler product is called decomposition reaction. They are represented by equation of the following term .

  AB       →        A+B

  (i) Thermat decomposition :-   when decomposition is carried out by heating.                           

For example     

(i) 2 Fe SO₄ (s)       →      CaO(s) +CO₂ (g)

(Ferrous sulphate )        (Ferric Oxide)                               

 Green colour                 Red brown colour

(ii)  CaCO₃ (s)                →        CaO(s)            +        CO₂ (g)

(calcium carbonate)       ( calcium oxide )   (Carbon dioxide )  Limestone quick lime

(iii) 2Pb (N0₃)₂   (s) + 2PbO(s)  →   4NO₂ (g)  +   O₂(g)    

(Lead nitrate )     (lead oxide)  (Nitrogen dioxide) oxygen   

                                                         Brown fumes

 Heating of lead nitrate and emission of nitrogen  dioxide

 ELECTRONIC DECOMPOSITION :- When decomposition is carried out by passing electricity .

 For Example:-  (i) Electrolysis of water.

 Electric current                                                                                        

 2H₂ 0       →        2H₂       +    O₂

(Water)            (Hydrogen)    (Oxygen)

Electrolysis of water is done as follows:-

When decomposition is carried out in presence of sunlight .

For Example:-

This is why silver chloride turns grey in sunlight because of the decomposition of silver chloride into silvers and chloride by light.

This reaction is used in black and write Photography.

DISPLACEMENT REACTION :-  This reaction in which more reactive element displace less reative element from its salt solution is called displacement reaction . they aare represented by equation of the following term.

A    +    BC   →       AC   +  B    

For example :- 

(i) Fe (s)  +  CuSO₄ (aq)     →       FeSO₄  (aq)   +   Cu(s)

(Iron )    (copper Sulphate)    ( iron Sulphate)      ( Copper)

Fe is more reactive than therefore iron (Fe) has displaced copper (Cu) from copper sulphate solution.

Iron nails dipped in copper sulphate solution.

(ii)  Zn(s) + CuS0₄ (aq)      →       ZnSO₄ (aq)    +  Cu (s)

(Zinc)  (copper sulphate)     ( Zinc  sulphate )   copper

Zinc is more reactive than copper , therefore it displace copper from copper sulphate solution.

(iii) Pb (s)   + CuCl₂ (aq)   →          PbCl₂ (aq)    + Cu(s)

(Lead)    (copper chloride)        ( lead chloride)    copper

Lead is more reactive elements than copper, therefore it displaces copper from copper chloride solution.

DOUBLE DISPLACEMENT REACTIVE :- The reaction in which the reactant ions exchange to form new products is called double displacement reaction . they are represented by equation of the following term.

AB    +    CD     →      AD    +        CB

For Example :-

Na₂ SO₄ (aq)   +  BaCl₂ (aq)    →    BaSO₄ (s)    +  2NaCl (aq)

(sodium sulphate) (barium chloride) (barium sulphate) (sodium chloride)

White precipitate of BaSO₄  is formed .the insoluble substance formed is known as precipitate . any reation that produces  a precipitate can be called a precipitation reaction.

OXIDATION AND REDUCTION :-

OXIDATION    →    

(i)  The addition of oxygen to reactant

(ii)   the removal of hydrogen from  reactant

 For Example  :-

  2 Cu +         O₂     →          2CuO

 (Copper)  (oxygen)       (copper oxide )

                                          Black substance

Oxidation  of copper to copper oxide

Reduction  : -      (i) the  addition of hydrogen to reactant

(ii) the removal of oxygen from a reactant.

 Redox Reactions :-  The reaction on written one substance gets oxidisied and other get reduced is known as redox reaction

  For example :-    

 CuO  +    H₂     →     Cu   +   H₂ O

In this reaction CuO is reduced to Cu and H₂  is oxidized to H₂ O . So oxidation and reduction taking place together, therefore it is a redox reaction.

 ZnO +  C    →        Zn    +   C O

    Here, C is oxidized to CO because oxygen is being added and ZnO IS REDUCED TO Zn because O is being removed.

   NOTE :→  

*  If a substance gains oxygen during a reaction , it is said to be oxidized.

*   If a substance loses oxygen during a reactant , it is said to be reduced.

ENDETHERMIC REACTION:-    Reaction in which energy is absorded are known as endothermic reaction.

For example :-  

 CaCO₃   (S)     →         CaO (s)   +   CO₂  (g)

FXOTHERMIC REACTION :-    Reaction in which heat is released along with formation of products.

For Example, 

CH₄ (q)   +  2 O₂  (g)     →    CO₂ (g) +2H₂ O(g) + Heat

EFFORTS OF OXIDATION REACTIONS IN EVERYDAY LIFE

1)  COROSION :-    When a metal is exposd to moisture ,air, acid eyc. For some time , a layer of hydrated oxide, is formed which e weaknes the metal and hence metal is said to be corroded.

Example of corrosion are:-

1. Rusting of iron .
2. Black coating on silver
3. Green coating on copper.

• Rusting of iron :-    when iron is exposed to oxygen in the presence of  moisture, reddish brown power is formed.

This is process is knon as rusting of iron.

• Method to prevent corrosion are:-  

1. Galvanization
2.  Electroplating
3. By putting paints

2) RANCIDITY:-    The oxidation of fats and oils when exposed to air is known as rancidity . due to rancidity, bad smell and bad taste of food occurs

Method of prevent rancidity are :-

1. By adding antioxidants.
2. Refrigeration. 
3. Replacing air by nitrogen.
4. Keeping food in air tight containers.

•  Chips manufactures fill bag of chips with nitrogen because it is non reactive gas and it prevent the chips from getting oxidized.

Chapter→2
ACIDS BASE AND SALTS

ACIDS:-  

*  These are the substances which have sour taste

• They turn blue litmus solution  red
• They give H⁺ ions in aqueous solution

Type  of acids are as follows →

1. Strong Acid :  Hcl , H₂SO₄
2. Weak Acid : CH₃COOH, HCOOH
3. Concentrated Acid : which have more amount of acid and less amount of water.
4. Dilute Acid : which have more amount of water and less amount of acid

BASES:-    

* These are the substances which have bitter taste and soapy in touch .

• They turn red litmus solution blue
• They give OH^- ions in aqueous solution.

Type of bases are as follows :-

1. Strong Bases :-  NaOH, KOH
2. WEAK BASES :- NH₄OH
3. ALKALI :- These are the bases which are soluble in water (like NaOH,KOH)

SALTS :-  A salts is a substance produced from the reaction of an acid and a base.

FOR EXAMPLE,  NaCl, KCl.

INDICATORS :- The substances that change their colour/ smell when they are added to acidic or alkaline solutions.

Types of Indicators.

1. Natural Indicators :- Litmus , turmeric
2. Synthetic Indicators :-  Phenolphthalein, Mythyl
3. Olfactory Indicators:- Onion, vanilla essence.

CHEMICAL PROPERTIES OF ACID AND BASES.

(i)  Reaction of Metal with :-     

(a)  ACIDS 

Acid   +  Metal      →           Salt +    Hydrogen  gas                      

For Example,

2Hcl           +          Zn       →             ZnCl₂      +    H₂

(Hydrochloric acid) (zinc)       (zinc chloride ) (hydrogen)

(b)  Bases  Base + metal     →   salt + hydrogen gas

For example,

2NsoH         +          Zn    →       Na₃ZnO₂        +        H₂ ↑ 

(sodium hydroxide )  ( zinc)  ( sodium zincate)  ( hydrogen gas )

 Hydrogen gas released can be tested by bringing burning candle near gas bubbles , it burst with pop sound.

(ii)    REACTION OF METAL CARBONATES / Metal Hydrogen carbonates with 

(a)  Acids :- 

(b) Acid   +   metal carbonate/ metal hydrogen carbonate   →  salt + carbon dioxide + water.

  For example :-

1. 2 Hcl + Na₂CO₃        →          2NaCl +  CO₂    +   H₂O
2. HCl +  NaHCO₃         →         NaCl +  CO₂  +H₂O
3. Ca(OH)₂   +   CO₂     →         CaCO₃ +  H₂O

                ( Line Water)                   ( White precipitate )

On passing CO₂  Through lime water , lime water turns milky and in this way CO₂ can be tested.

And when excess CO₂ is passed , milkiness disappers and the following reaction take place.

Passing carbon dioxide gas through calcium hydroxide.

CaCo₃    +    CO₂     +     H₂O    →    Ca (HCO3)

(iii)     Reaction of acids and bases with each other  

Acid     +       base        →       Salt   +    H₂O

Neutralisation Reaction :-   The reaction between an aid and a base to give a salt and water is known as neutralization reaction.

  For Example :-    

HCl (aq)   + NaOH (aq)    →  NaCl(aq) + H₂O(l)

(iV)  Reaction of Metallic Oxides with Acids  :-

Metallic oxide +  acid     →     salt   +  water

 Metallic oxides are basic in nature

For Example :- 

CaO  + 2HCl    →   CaCl₂   +    H₂O

(v) Reaction of Non- metallic Oxide with base :-

Non- metallic oxide + base    →        salt   +    H₂O 

Non- metallic oxide are acidic in nature.

For Example :-

CO₂   +   Ca(OH)₂    →   CaCO₃  +    H₂O

CHAPTER→  3

METALS AND NON- METALS

 Physical properties of metals and non- metals.

 1     LUSTRE  -

→  Metals have shining surface.

  →  Non-metals do not have shinig surface

 * Except iodine   

 2 Hardness -

 →   Metals are generally hard.

• Except sodium lithium and potassium which are soft and can be cut with knife.

→ Non – metals are generally soft.

• Except diamond a form of carbon which is the hardest natural substan

3 Malleability - 

 Metals can be beaten into thin sheets gold and silver are the most malleable metals.

Non- metals are non-mallerable.

4 Ductility -

 →Metals can be drawn into thin wires.

 →Non- metals can not be drawn into thin wires. They are non- ductile.

5 Conductor of heat and electricity -

→Metals are good conductor of heat and electricity of heat. Lead and mercury are poor conductor of heat.

Metals are good conductor of heat and electricity.

→Non- metals are poor conductor of heat EXCEPT graphite.

6 state -

→ The metals exist as solids. EXCEPT mercury

→The non- metals exist as solids or gaseous.EXEPT  bromine.

7 density -

→Metals have high density and high melting point.

→Except sodium and potassium.

8 Oxides -

→ Metallic oxides are basic in nature.

→Non metallic oxides are acidic in nature.

9 Sonorous -

→Metals produce a sound on strinking a hard surface.

→Non-metals are not sonorous.

CHAPTER - 4

COMPOUND AND ITS CARBON

Carbon:-  (i)Most carbon compounds are poor conductors of electricity. Therefore, the bonding in these componds does not give rise to any ions

1. They have low melting and boiling points as compared to ionic compounds
2. Forces of attraction between the molecules are not very strong.
3. The atomic number of carbonis
4. It has four elements I its outermost shell and needs to gain or lose four electrons to attain noble gas configuration.

If carbon were to gain lose electrons:-

• It  could gain four electrons forming C^4-  anion. But it would be difficult for the nucleus with six protons to hold on to ten electrons.
• It could lose four electrons forming C⁴⁺ cation. But it would require a large amount of energy to remove electrons. Carbon overcomes this problem by sharing its valences electrons with other of carbon or with atoms of other elements.

 Bonding in carbon

 Carbon form covalent bounds.  

• Covalent bound formation involves sharing of electrons between bonding atoms which may be either same or different.
• The number of lectrons contributed by an atom for sharing is known as its covalency.

       For example 

1. Molecule of hydrogen  

   

2. Molecule of oxygen

3. Molecule of Nitrogen

4. Structure of methane

Methane CH₄  is widely used as a fixed and is a major component of bio-gas and compressed natural gas ( CNG) .

• CHARACTERISTICS OF COVALENT COMPOUND :-

1. These compounds are molecula in nature  i.e. , they exist as single molecules)
2. These are insoluble in water and soluble in benzene , kerosence and petrol etc.
3. These compounds are poor conductor of electricity.

Allotropes of carbon

  The property due to which an element exists intwo or more forms, which differ in their physical and chemical properties is known as ‘Allotropes’ and the various forms are called “Allotropes”.

• carbon exists in two allotropic  form
• (i) crystalline                 (ii)  amorphous.

The crystalline forms are diamond and graphite the amorphous forms are coat , characol etc.

• In diamound , each carbon is bonded to for other carbon atoms forming rigid 3-D dtructure. Diamond is the nardest substance. * in graphite each carbon is bonded to three other carbon atom. Graphite structure is formed by the hexagonal arrays. Graphite is smooth and slippery . it is very good conductor of electricity.
• Fullerenes form another class of carbon allotropes. The first one to be identified was C-60, which has carbon atoms arranged in the shape of a football.

Chapter 5

Periodic Classification of Elements

Early classification of elements.

• Classification means identifying similar species and grouping them together.
• Lavoisier divided lements into two main types known as metals and non- metals.

Doberiner’s law of triads :- 

Doberiner tried to arrange the elements with similar properties. He showed that when the three elements in a triad were written in the order of increasing atomic masses ; the atomic masses ; the atomic mass of the middle element was roughly the average of the atomic masses of other two elements.

i.e atomic masses Li, Na and K are 7,23, and 39 respectively , tus the mean of 1^st and 3^rd elements is 23 and the atomic mass of middle element is 23.

Limitation :-  He could identify only a few such triads and so the law could not gain importance.

For example, Fe, Co, Ni, all the three elements have nearly equal atomic mass and thus does not follow this law.

Newland’s law of octaves :-

 He found that every eighth element had properties similar to that of the first. HP compared this to the octaves found in music . he called it ‘law of octaves’.

For Example. The properties of lithium and sodium were found to be the same sodium is the eighth  element after lithium.

 Limitation :-

(i) law of octaves was applicable only upto calcium as after calcium every eigth element did not possess properties similar to that of first.

(II)    According  to him , only 56 elements exist in nature and no more elements would be discovered in the future. But later on several new element  were discovered wose properties did not fit into law of octaves. 

(III)    In order to fit new elements into his table newland  adjust two elements in the same column, but put some unlike elements under the same column. 

(IV)    Thus, newland’s  classification was not accepted.