Laws of Chemical Combination

LAWS OF CHEMICAL COMBINATION A multitude of objects in our surroundings are visible and tangible. Matter is defined as anything that occupies space and has mass. John Dalton is credited with formulating the theory on the structure of matter, which is now recognized as Dalton's atomic theory. To comprehend the composition of different compounds, it's essential to possess a theory that accommodates both qualitative and quantitative observations during chemical transformations. Antoine Lavoisier, John Dalton, and other scientists formulated specific laws related to the composition of matter and chemical reactions. These principles are collectively referred to as the laws of chemical combination. 1. The Law of Conservation of Mass This principle, attributed to Lavoisier, is also known as the law of the indestructibility of matter. According to this law, the total mass of a system remains constant in all chemical changes; that is, mass is neither created nor destroyed during a chemical change. In chemical changes, the conservation of total mass holds true. This means that the mass before the reaction is always equal to the mass after the reaction. All chemical reactions adhere to this law, making it the foundation for all quantitative work in chemistry. Therefore, in a chemical change, the total masses of reactants are equal to the total masses of products. This relationship is valid when reactants are completely transformed into products. However, if the reacting materials are not entirely consumed, the relationship becomes: Total masses of reactants = Total masses of products + Masses of unreacted reactants. In summary, the conservation of total mass is a fundamental aspect of chemical changes, ensuring that the mass before the reaction is always equal to the mass after the reaction. H2 + 1/202 → H20 (λ) (g) (g) 1 mole mole 1 mole mass before the reaction = 1 x2+1/2×32=18 gm mass after the reaction = 1 x 18 = 18 gm Ex. A 15.9g sample of sodium carbonate is added to a solution of acetic acid weighing 20.0g. The two substances react, releasing carbon dioxide gas to the atmosphere. After reaction, the contents of the reaction vessel weigh 29.3g. What is the mass of carbon dioxide given off during the reaction? Sol. The total mass of reactants taken 15.9 +20.0 35.9 gm. From the conservation of mass, the final mass of the contents of the vessel should also be 35.9 gm. But it is only 29.3 gm. The difference is due to the mass of released carbon dioxide gas. Hence, the mass of carbon dioxide gas released = 35.9-29.36.6 gm Ex. H2(g)+02 (g) Before reaction initially 1 mole H20 (λ) mole After the reaction 0 0 1 mole mass before reaction = mass of 1 mole H2 (g) + mole O2 (g) 2+16=18 g mass after reaction = mass of 1 mole water = 18 g 2. Law Of Constant or Definite Proportion Formulated by Proust in 1799, this law can be expressed as follows: A chemical compound consistently comprises the same elements combined in fixed proportions by mass. In other words, a chemical compound maintains a fixed composition that remains independent of its preparation method or the source from which it is derived. All chemical compounds exhibit a consistent composition, regardless of how they are prepared or their origin. Ex. In water (H2O), Hydrogen and Oxygen combine in 2: 1 molar ratio, the ratio remains constant whether it is tap water, river water or sea water or produced by any chemical reaction. Ex. 1.80 g of a certain metal burnt in oxygen gave 3.0 g of its oxide. 1.50 g of the same metal heated in steam gave 2.50 g of its oxide. Show that these results illustrate the law of constant proportion. Sol. In the first sample of the oxide, wt. of metal 1.80 g, wt. of oxygen (3.0-1.80) g = 1.2 g wt.of metal 1.80 81.5 wt.of oxygen 1.2 g In the second sample of the oxide, wt. of metal 1.50 g. wt. of oxygen (2.50-1.50) g=1g wt. of metal 1.50 g A 1.5 wt. of oxygen 1 g Thus, in both samples of the oxide the proportions of the weights of the metal and oxygen are fixed. Hence, the results follow the law of constant proportion. 3. The Law of Multiple Proportion Proposed by Dalton in 1808, Dalton's Law states that if two elements combine to create more than one compound, the distinct masses of one element combining with a fixed mass of the other element exhibit a simple ratio to each other. For instance, hydrogen and oxygen form two compounds: H₂O (water) and H₂O₂ (hydrogen peroxide). When an element combines with another element to produce multiple different compounds, the mass of one element, combining with a constant mass of the other, bears a simple ratio to each other. Note: In this context, "simple ratio" refers to the ratio between small natural numbers, like 1:1, 1:2, 1:3. Subsequently, this straightforward ratio evolves into the element's valency and eventually its oxidation state. Ex: Carbon and Oxygen when combine, can form two oxides viz CO (carbon monoxide), CO2 (Carbon dioxides) In CO2, 12 g carbon combined with 32 g of oxygen. In CO, 12 g carbon combined with 16 g of oxygen. Thus, we can see the mass of oxygen which combine with a constant mass of carbon (12g) bear simple ratio of 16:32 or 1: 2.