Chemistry Basics – Study Notes for Chandigarh Police Sub‑Inspector (SI) Exam
Subject: General Awareness & Current Affairs – Chemistry Basics
Aim of the notes – Give you an easy‑to‑remember overview of the most common chemistry concepts that frequently appear in the SI exam.
How to use – Read the simple explanation, memorize the key points, practice the worked example, and then use the quick revision summary for last‑minute checks.
1. Simple Explanation
Chemistry is the science that studies matter, its properties, structure, composition, and the changes it undergoes. At the SI level we focus on:
- Basic building blocks – atoms, molecules, ions.
- How atoms combine – chemical bonds (ionic, covalent, metallic).
- The Periodic Table – organization, trends (atomic radius, electronegativity).
- Acids, bases, and neutralization – pH, common strong/weak acids & bases.
- Stoichiometry – balancing equations, converting between units (moles, grams, liters).
- Solutions & concentrations – molarity, normality, how to prepare them.
- Key industrial & environmental applications – why chemistry matters in law‑enforcement (police, forensic science, environmental monitoring).
2. Key Points (Bullet‑style)
| Topic | What to Know |
|---|---|
| Atoms & Periodic Table | • Atomic number = number of protons. • Mass number = protons + neutrons. • Electrons occupy shells; valence electrons decide reactivity. • Periodic trends: • Atomic radius ↑ down a group, ↓ across a period. • Electronegativity ↑ down a group, ↓ across a period. |
| Chemical Bonds | • Ionic (opposite charges, typically metal + non‑metal). • Covalent (sharing of electrons, usually non‑metal + non‑metal). • Metallic (sea of delocalized electrons). |
| Molecular Formula vs Empirical Formula | • Molecular = actual atoms in one molecule. • Empirical = simplest whole‑number ratio. |
| Balancing Equations | • Count atoms of each element on both sides, adjust coefficients. • Example: N₂ + 3H₂ → 2NH₃. |
| Stoichiometry | • 1 mole = 6.02 × 10²³ entities. • Mass of 1 mole = molar mass (g/mol). • Use mole ratio from balanced equation. |
| Acids & Bases | • Strong (completely dissociate): HCl, H₂SO₄, NaOH, KOH. • Weak (partial dissociation): Acetic acid, ammonia. • pH = –log[H⁺]; lower pH = more acidic. |
| Solution Concentration | • Molarity (M) = moles solute / liters solution. • Normality (N) = equivalents / liters. • Equivalent weight = molar mass / n (n = number of replaceable H⁺ or OH⁻). |
| Major Units | • Mass: grams (g), kilograms (kg). • Volume: milliliters (mL), liters (L). • Pressure: atmospheres (atm), pascals (Pa). |
| Common Forensic & Environmental Chemistry | • Detection of explosives (nitro‑compounds). • Identification of toxic gases (CO, NO₂). • Water quality tests (pH, hardness). |
| Important Chemical Reactions in SI Context | • Combustion (CₓHᵧ + O₂ → CO₂ + H₂O). • Neutralization (Acid + Base → Salt + Water). |
3. Worked Example
Problem
A forensic chemist has a 100 mL sample of an unknown solution that contains 0.2 M sodium hydroxide (NaOH).
- How many grams of NaOH is in the sample?
- If the solution is mixed with 300 mL of 0.1 M hydrochloric acid (HCl), what is the final pH of the mixture?
Step‑by‑Step Solution
1. Mass of NaOH in 100 mL
-
Molarity (M) = moles / liters
(0.2,\text{M} = \frac{n_{\text{NaOH}}}{0.1,\text{L}})
→ (n_{\text{NaOH}} = 0.2,\text{mol/L} \times 0.1,\text{L} = 0.02,\text{mol}) -
Molar mass of NaOH = 23 (Na) + 16 (O) + 1 (H) = 40 g/mol
-
Mass = moles × molar mass
(0.02,\text{mol} \times 40,\text{g/mol} = 0.8,\text{g})
Answer 1: 0.8 g of NaOH.
2. Final pH after mixing with HCl
-
Moles of NaOH: (0.02,\text{mol}) (from above)
-
Moles of HCl:
(0.1,\text{M} = \frac{n_{\text{HCl}}}{0.3,\text{L}})
→ (n_{\text{HCl}} = 0.1,\text{mol/L} \times 0.3,\text{L} = 0.03,\text{mol}) -
Reaction: NaOH + HCl → NaCl + H₂O (1:1 ratio).
-
Remaining species:
- HCl excess: (0.03 - 0.02 = 0.01,\text{mol}) HCl.
- No NaOH left.
-
**Total