Some Basic Concepts in Chemistry

Chemistry is the study of matter, its properties, composition, and the changes it undergoes during chemical reactions. Basic concepts in chemistry lay the foundation for understanding atomic and molecular behavior.

1.0Important Terminology and Formula

1. Atomic Weight: Atomic weight (or atomic mass) is the weighted average mass of an element’s atoms, considering all isotopes. It is usually expressed in atomic mass units (amu).

$Atomicweight(RelativeAtomicweight)=\frac{ActualmassofoneatomofElement}{1/12thmassofoneatomofCarbon-12}$

2. Molecular Weight: Molecular weight (or molecular mass) is the sum of the atomic weights of all atoms in a molecule. It helps determine how much of a substance is needed for reactions and calculations.

Example:

Calculate the actual mass of a molecule of the compound C₆₀H₂₂ is:  

Solution - The relative molecular weight of C₆₀H₂₂ = 742

The actual mass of one molecule of the given compound will be = 742 amu.

= 742 x 1.67 x 10^–24 gram

= 1.24 × 10^–21 gram

3. Mole Concept:

• The mole concept is a method to quantify atoms, ions, and molecules. One mole is defined as 6.022 × 10²³ entities (Avogadro's number), providing a bridge between atomic and macroscopic scales.
• A mole is the amount of a substance that contains as many entities (atoms, molecules or other particles) as there are atoms in exactly 0.012 kg (or 12g) of the carbon - 12 isotope.

4. Gram Atomic Mass and Gram Molecular Mass: Gram atomic mass is the mass of one mole of an element’s atoms in grams, equal to the atomic weight. Gram molecular mass is the mass of one mole of a compound in grams, equal to the molecular weight.

2.0Relationships in Mole Concept Calculations

Percentage Composition, Empirical Formula & Molecular Formula:

• Percentage Composition: Indicates the mass percentage of each element in a compound.
• Empirical Formula: Represents the simplest whole-number ratio of atoms in a compound.

Examples- 

• Molecular Formula: Shows the exact number of atoms of each element in a compound, which can be a multiple of the empirical formula.

Example-

In a compound, X is 75.8% and Y is 24.2% by weight. If atomic weight of X and Y are 24 and 16 respectively, then, calculate the empirical formula of the compound.

Solution-

Empirical formula = X₂Y

3.0Laws of Chemical Combination

The Laws of Chemical Combination describe the fundamental principles that govern how elements combine to form compounds. Let’s learn in brief-

1. Law of Mass Conservation (Law of Indestructibility of Matter): According to this law, Mass is conserved in a chemical reaction. It was given by Lavoisier and tested by Landolt.

Total mass of reactants = Total mass of products

2. Law of Definite Proportions/ Law of Constant Composition: According to this law, Compounds always contain elements in fixed proportions by mass. It was given by Proust.

Ex. In water (H₂O), hydrogen and oxygen are always combined in a fixed ratio of approximately 1:8 by mass, meaning that water from any source will always have this same composition.

3. Law of Multiple Proportions: It was given by John Dalton.  According to this law, When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios.

Ex. Nitrogen and oxygen combine to form five oxides, which are: Nitrous oxide (N₂O), nitric oxide (NO), nitrogen trioxide (N₂O₃), nitrogen tetraoxide (N₂O₄) and nitrogen pentaoxide (N₂O₅).

4. Law of Gaseous Volume (Gay-Lussac's Law)

Gaseous reactants and products combine in simple whole-number volume ratios at constant temperature and pressure.

Example: 1 volume of N₂ + 3 volumes of H₂ → 2 volumes of NH₃ (ratio 1:3:2).

5. Avogadro's Law: Equal gas volumes at the same temperature and pressure have an equal number of molecules. Example: H₂ + Cl₂ → 2HCl, where 1 volume of H₂ and Cl₂ each gives 2 volumes of HCl with equal molecule counts.
6. Equivalent Weight: Equivalent weight is the mass of a substance that reacts with or replaces one mole of hydrogen atoms or combines with one mole of hydroxide ions. It is useful in stoichiometric calculations, especially in titrations and redox reactions.

Equivalent Weight =  Molecular weight/Valency Factor

Ex. Calculate the Equivalent weight of SO₄^–2.

Solution:

Equivalent weight = 96/2 = 48

4.0Concentration Terms

Concentration Terms refer to ways of expressing the amount of solute present in a solution. Each term is suited to specific applications based on the nature of the solution and the measurement needs. Here are the common concentration terms:

1. Molarity (M): The number of moles of solute per litre of solution. It’s widely used in laboratories to prepare solutions for reactions.

$Molarity=\frac{numberofmolesofsolute}{Volumeofsolution(L)}=\frac{n}{V_{(L)}}$

2. Molality (m): The number of moles of solute per kilogram of solvent. It’s useful in cases where temperature changes are involved, as molality is not affected by temperature.

$Molalityofasolution=\frac{Numberofmolesofsolute}{weightofsolvent(kg)}=\frac{Numberofmolesofsolute\times 1000}{weightofsolvent(g)}$

3. Normality (N): The number of gram equivalents of solute per liter of solution. It’s often used in titrations, particularly in acid-base and redox reactions.

$Normality(N)=\frac{Numberofgramequivalentsofsolute(geq)}{Volumeofsolution(L)}$

$Normality(N)=\frac{Massofsolute(g)}{Equivalentmass(E)\times Volumeofsolution(L)}$

4. Mole Fraction (X): The ratio of the moles of one component to the total moles of all components in the solution. It’s used in studies involving partial pressures and colligative properties.​

$Molefractionofsolute{X}_B=\frac{molesofsolute(n)}{Molesofsolute(n)+molesofsolvent(N)}$

$Molefractionofsolute{X}_A=\frac{molesofsolute(n)}{Molesofsolute(n)+molesofsolvent(N)}$

Where, X_A + X_B = 1

5. Parts Per Million (ppm): A way to express very dilute concentrations, indicating the parts of solute per million parts of solution. It’s common in environmental science, for measuring pollutant levels.