Hsslive.net provided Plus One Chemistry notes for students in their higher secondary years in two languages English Medium & Malayalam Medium. Topics- “Importance of Chemistry” that are usually covered in the first year of chemistry at the higher secondary level. valuable for both Kerala Syllabus and CBSE students with knowledge cultivated over two decades of teaching experience.
Importance of Chemistry
1. Matter and Its Classification
Matter, the substance that occupies space and possesses mass, is classified into three states: solid, liquid, and gas. For example:
- Ice (solid) – H2O in a solid state
- Water (liquid) – H2O in a liquid state
- Water vapor (gas) – H2O in a gaseous state
2. Laws of Chemical Combinations
Chemical reactions follow well-defined laws, which include:
a) Law of Conservation of Mass: The total mass of substances in a chemical reaction remains constant, indicating that atoms are neither created nor destroyed during the process.
Example: When hydrogen (H2) reacts with oxygen (O2) to form water (H2O), the total mass of reactants (H2 + O2) is equal to the total mass of products (H2O).
b) Law of Definite Proportions: Elements combine in fixed mass ratios to form compounds, leading to their unique identities.
Example: Water (H2O) always has a fixed mass ratio of hydrogen and oxygen, regardless of its source.
c) Law of Multiple Proportions: When two elements form more than one compound, the ratios of masses of one element to a fixed mass of the other element are in small whole numbers.
Example: Carbon and oxygen can form two compounds: carbon monoxide (CO) and carbon dioxide (CO2). The mass ratio of oxygen to carbon in CO2 is twice that of CO.
3. Atomic and Molecular Mass
Understanding the atomic and molecular masses is vital in quantifying substances’ amount in chemical reactions. The atomic mass unit (amu) is used to express these masses relative to the carbon-12 isotope.
Example: The atomic mass of carbon is 12.011 amu, while the molecular mass of water (H2O) is (2×1.008 amu) + 16.00 amu = 18.016 amu.
4. Mole Concept
The mole concept bridges the macroscopic and microscopic worlds by connecting mass, number of particles, and Avogadro’s number (6.022 x 10^23). It enables easy conversions between mass, moles, and particles.
Example: One mole of any substance contains Avogadro’s number of particles, which is approximately 6.022 x 10^23. Therefore, one mole of atoms of an element has a mass equal to its atomic mass in grams.
Numerical Example: Calculate the number of moles in 60 grams of carbon dioxide (CO2).
Solution: Molar mass of CO2 = (12.01 g/mol + 2×16.00 g/mol) = 44.01 g/mol Number of moles = Mass of CO2 / Molar mass of CO2 Number of moles = 60 g / 44.01 g/mol ≈ 1.36 moles