Chemical Combination

2.1 LAWS OF CHEMICAL COMBINATIONS
Introduction:-

Chemistry deals with the matter and the changes occurring in it, chemists are particularly interested in these changes, where one or more substances are changed into quite different substances. They had found that these chemical changes are governed by some empirical laws known as laws of chemical combinations.

These laws are: 

  1. Law of conservation of mass
  2. Law of constant composition
    (or) Law of definite proportions
  3. Law of multiple proportions
  4. Law of reciprocal proportions.
  1. Law of Conservation of Mass: Matter under goes changes. However, it has been found that in all chemical changes, there is no change in the mass of the substances being changed. For example, in iron (Fe) increase in weight on rusting is because of its combination with oxygen from the air and the increase in weight is exactly equal to the weight of oxygen combined. The French Chemist Lavosier, (1785) tried to learn about chemical changes by weighing the quantities of substances used in chemical reactions. He found that when a chemical reaction was carried out in a closed system, the total weight of the system was not changed. The most important chemical reaction that Lavosier performed was the decomposition of the red oxide of mercury to form metallic mercury and a gas; he named this gas as oxygen. Lavosier summarised his findings by formulating a law, which is known as law of conservation of mass. It states that mass is neither created nor destroyed during a chemical reaction. In other words chemical reaction the initial weight of reacting substances is equal to the final weight of the, in any products.
    The law of conservation of mass may be demonstrated by the union of hydrogen (H2) and oxygen (O2) to form water. If the H2 and O2 are weighed before they unite, it will be found that their combined weight is equal to the weight of water (H2O) formed.

Practical Verification: (Landolt Experiment)

German Chemist H. Landolt, studied about fifteen different chemical reactions with a great skill, to test the validity of the law of conservation of mass. For this, he took H-shaped tube as shown in fig. 2.2 and filled the two limbs A and B, with silver nitrate (AgNO3) in limb A and hydrochloric acid (HCl) in limb B. The tube was sealed so that the material could not escape outside. The tube was weighed initially in a vertical position so that the solutions should not intermix with each other. The reactants were mixed by inverting and shaking the tube. The tube was weighed after mixing on the formation of white precipitate of (AgCl). He observed that weight remains same.

                         AgNO3(aq) + HCl(aq)   AgCl(s) + HNO3(aq)

                                                                     white ppt:

Thus total mass of the substance before the reaction is equal to the total mass of the substances after the reaction.

In ordinary chemical changes, relatively small amount of energy released. But in nuclear changes where uranium atoms undergo fission (break up) into smaller atoms plus neutrons, the total mass of products is noticeably less than that of starting material. This clearly indicates that some mass of uranium has been converted into energy, which is evident to us as heat and radiation.

The relationship between mass that is lost and the energy that is released is given by the equation.

                                                 E = me2

Where (E) is the energy in ergs, (m) is the mass in grams and (C) is the velocity of light in centimeters per second, (3x1010 cm/sec). This relationship between mass and energy was first proposed in (1906), by the famous Physicist and Mathematician, Albert Einstein.

It follows that for every chemical change; there will be a mass change. But this mass change is too small that no one has yet been able to detect through weighing techniques.

Hence the law of conservation of mass is,' therefore, still valid from practical view point for ordinary chemical reactions i.e., "there is no detectable gain or loss of mass in a chemical reaction".

 

  1. Law of Constant Composition or Law of Definite (fixed) Proportion  By the end of Eighteenth century, chemists showed that a given compound has a definite (constant) composition. French chemist Louis Proust in (1799) summarized this result in the form of the law of definite proportion (also known as constant composition) which states, that different samples of the same compound always contain the same elements combined together in the same proportions by mass.

For instance every sample of pure water, though prepared in the Laboratory or obtained from rain, river or water pump contains one part Hydrogen (H) and (8) parts oxygen (O) by mass

          e.g.:   H2O

                    2:16
                   1:8 (parts by mass)

One of the earliest illustrations of the law of definite proportions is found in the work of Swedish chemist J. J Berzelius (1779-1848).

Berzelius us heated 10g of lead (Pb) with various amounts of sulphur (S). He got exactly 11.56g of lead sulphide and the excess of sulphur was left over, when he used 18g of lead (Pb) with 1.56g of sulphur (S), he got exactly 11.56g of lead sulphide (PbS) and the 8g of lead (Pb) remained unused.

  1. Law of Multiple Proportions: The fact that the same element, can combine in more than one ratio to form different compounds was published by, John Dalton, (1803) in the form of law of multiple proportion: "It states that if two elements combine to form more than one compounds. The masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers or simple multiple ratio."

For example: Carbon (C) forms two stable compounds with oxygen (O) namely carbon monoxide (CO) and carbon dioxide (CO2).

 

Compound

Mass of Carbon (C)

Mass of Oxygen (O)

Ratio of Oxygen (O)

Carbon monoxide CO Carbon dioxide CO2

12

12

16

32  

1

2

The different masses of oxygen 16 and 32 which combine with the fixed mass of C (12g) are in ratio of [16:32], i.e. 1:2, which is simple whole number ratio, and obeys the law of multiple proportion.

Another illustration of this law is the formation of water (H2O) and (H2O2) from hydrogen and oxygen.

Compound

Mass of Hydrogen (H)

Mass of Oxygen (O)

Ratio of Oxygen (O)

Water H2O

Hydrogen peroxide (H2O2)

       2

       2

     16

     32

       1

       2

The different masses of oxygen 16:32, which combine with the fixed mass of hydrogen (2g) are in ratio of 16:32 i.e. 1:2 which is again in a ratio of simple whole numbers. The excellent illustration of law of multiple proportion is furnished, when the elements nitrogen (N) and oxygen (O) combine together to form a series of five oxides of nitrogen, in which these two elements are present in different proportions.

 

S.No.

 

Name of Oxides

Mass of

   (N)

Mass of

    (Q)

Fixed mass of

   (N)

Variable mass of

   (O)

Ratio of (O)

  1.

Nitrous oxide (N2O)

28

16

14

8(1x8)

1

  2.

Nitric oxide (NO) ,

14

16

14

16(2x8)

2

  3.

Nitrogen trioxide

(N2O3)

28

48

14

24(3x8)

3

  4.

 

Nitrogen tetra oxide

(N2O4)

28

64

 

14

32(4x8)

4

   5.

Nitrogen penta oxide(N2O5)

28

80

14

40(5x8)

5

By fixing the mass of (N), the mass of (O) in different oxides varies.

i.e.  8 : 16 : 24 : 32 :  40                                    

       1  :  2:   3 :  4   :   5

These figures (in multiple ratios), are according to the law of multiple proportion.

4.      Law of Reciprocal Proportion: This law was enunciated by Ritcher in (1792-94). It states that "when two different elopements separately combine with the fixed mass of third element, the proportions in which they combine with one another shall be either in the same ratio or some simple multiple of it".

For instance, when two elements C and O separately combine with if to form methane (CH4) and water (H2O) respectively it is very’ clear, that in methane 3g of C combine with 1g of hydrogen, and in water (H2O) 8g of t combine with the same (fixed) mass i.e. (1g) of H, now when C and O combine with each other to form carbon dioxide (CO2), they do so in the same proportion i.e. 12:32 = 3:8 parts by mass.

Another illustration of law of reciprocal proportion is provided when, 12g of C combine with 32g of O to form carbon dioxide (COJ and 32g of sulphur (S) combine with the same (fixed) mass of oxygen (O) i. e. 32 g to form sulphur dioxide.

The above example shows that the mass of C and S that combine with the same mass of O are in the proportion of 12:64 i.e. 3:16.

 According to the statement of law of reciprocal proportion, that the proportion in which C and S combine with one another shall be either in the same ratio (3:8) or some simple multiple of it i. e (3:16).

It is very clear that in the formation of carbon disulphide (CS2), C and S combine in the ratio of (12:64) i.e. (3:16) which is simple multiple of (3:8).

 
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