Introduction
Chemistry, often dubbed the "central science," bridges other natural sciences, including physics, geology, and biology. The study of chemistry is essential for understanding the world around us and is a fundamental part of the NEET Chemistry syllabus. This study note document will cover the key concepts of chemistry necessary for NEET preparation, breaking down complex ideas into digestible sections.
Matter and Its Nature
Definition of Matter
Matter is anything that has mass and occupies space. It can exist in three primary states: solid, liquid, and gas.
Classification of Matter
Matter can be classified into:
- Pure Substances
- Elements: Substances that cannot be broken down into simpler substances by chemical means. E.g., Oxygen ($O_2$), Hydrogen ($H_2$).
- Compounds: Substances composed of two or more elements chemically combined in a fixed ratio. E.g., Water ($H_2O$), Carbon Dioxide ($CO_2$).
- Mixtures
- Homogeneous Mixtures: Mixtures that have a uniform composition throughout. E.g., Saltwater.
- Heterogeneous Mixtures: Mixtures that do not have a uniform composition. E.g., Sand in water.
Remember, compounds have fixed ratios of elements, while mixtures can have variable compositions.
Atomic and Molecular Masses
Atomic Mass Unit (amu)
The atomic mass unit is defined as one-twelfth the mass of a carbon-12 atom. It is used to express atomic and molecular masses.
$$ 1 , \text{amu} = \frac{1}{12} , \text{mass of one carbon-12 atom} $$
Calculation of Molecular Mass
The molecular mass of a substance is the sum of the atomic masses of all atoms in a molecule.
ExampleFor example, the molecular mass of water ($H_2O$) is calculated as: $$ \text{Molecular mass of } H_2O = 2 \times \text{atomic mass of H} + 1 \times \text{atomic mass of O} $$ $$ = 2 \times 1.008 , \text{amu} + 16.00 , \text{amu} $$ $$ = 18.016 , \text{amu} $$
Mole Concept and Molar Mass
Definition of Mole
A mole is a unit that measures the amount of substance. One mole contains exactly $6.022 \times 10^{23}$ entities (Avogadro's number).
Molar Mass
The molar mass of a substance is the mass of one mole of that substance, typically expressed in grams per mole (g/mol).
ExampleFor instance, the molar mass of carbon dioxide ($CO_2$) is: $$ \text{Molar mass of } CO_2 = 12.01 , \text{g/mol (C)} + 2 \times 16.00 , \text{g/mol (O)} $$ $$ = 44.01 , \text{g/mol} $$
NoteThe molar mass in grams per mole is numerically equal to the molecular mass in amu.
Stoichiometry
Balancing Chemical Equations
Chemical equations must be balanced to obey the Law of Conservation of Mass. This means the number of atoms of each element must be the same on both sides of the equation.
ExampleFor example, to balance the combustion of methane: $$ CH_4 + O_2 \rightarrow CO_2 + H_2O $$ Balanced equation: $$ CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O $$
Stoichiometric Calculations
Stoichiometry involves calculating the amounts of reactants and products in a chemical reaction using the balanced equation.
ExampleFor example, in the reaction $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$, if we start with 16 g of $CH_4$ (1 mole), we need 2 moles of $O_2$ (64 g) to completely react with it.
Common MistakeEnsure to convert all masses to moles before using stoichiometric ratios.
Concentration Terms
Molarity (M)
Molarity is the number of moles of solute per liter of solution.
$$ M = \frac{\text{moles of solute}}{\text{volume of solution in liters}} $$
ExampleTo prepare 1 M solution of NaCl, dissolve 58.44 g of NaCl (1 mole) in enough water to make 1 liter of solution.
Molality (m)
Molality is the number of moles of solute per kilogram of solvent.
$$ m = \frac{\text{moles of solute}}{\text{mass of solvent in kg}} $$
Normality (N)
Normality is the number of gram equivalents of solute per liter of solution.
$$ N = \frac{\text{gram equivalents of solute}}{\text{volume of solution in liters}} $$
NoteNormality is often used in acid-base and redox reactions.
Laws of Chemical Combination
Law of Conservation of Mass
Mass is neither created nor destroyed in a chemical reaction.
Law of Definite Proportions
A given compound always contains exactly the same proportion of elements by mass.
Law of Multiple Proportions
If two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in the ratio of small whole numbers.
ExampleFor example, carbon forms two oxides: $CO$ and $CO_2$. The mass ratio of oxygen in $CO$ and $CO_2$ is 1:2.
Empirical and Molecular Formulas
Empirical Formula
The empirical formula gives the simplest whole-number ratio of atoms in a compound.
Molecular Formula
The molecular formula gives the actual number of atoms of each element in a molecule of the compound.
ExampleFor example, the empirical formula of glucose is $CH_2O$, while the molecular formula is $C_6H_{12}O_6$.
TipTo find the molecular formula, divide the molar mass by the empirical formula mass and multiply the subscripts in the empirical formula by this number.
Conclusion
Understanding these basic concepts of chemistry is crucial for mastering the subject and performing well in the NEET exam. By breaking down complex ideas into smaller, manageable sections and using examples, we can grasp these fundamental principles more effectively. Remember to practice stoichiometric calculations, balance chemical equations, and familiarize yourself with concentration terms to excel in NEET Chemistry.
