Some Basic Concepts of Chemistry

Chemistry is the study of matter and the changes it undergoes. Look around your room — the table, the water in your bottle, the air you breathe — everything that has mass and occupies space is matter.

We generally see matter in three forms: solid, liquid and gas. In your kitchen, a steel dabba is a solid — it has a fixed shape and volume; water in a glass is a liquid — it has fixed volume but takes the shape of the glass; and the LPG in the cylinder is a gas — it has neither fixed shape nor fixed volume.

pure substance and mixture. A pure substance has a fixed composition. An element contains only one type of atom, like iron (Fe) or gold (Au), while a compound contains different atoms in a fixed ratio, like water (H₂O). A mixture does not have a fixed ratio — like the air around us or a spoon of sugar dissolved in tea.

To describe matter, we use two types of properties. Physical properties can be observed without changing its identity, such as colour, density, melting point and boiling point. Chemical properties are observed only when it changes into a new substance, such as burning of paper, rusting of iron, or curdling of milk due to acidity.

In daily life we never say “give me milk”, we say “give me one litre of milk”. Every measurement needs a unit to make sense. Scientists worldwide use the same system to avoid confusion, called the SI system which has seven base units.

These are length in metre (m), mass in kilogram (kg), time in second (s), temperature in kelvin (K), amount of substance in mole (mol), electric current in ampere (A) and luminous intensity in candela (cd). All other units are derived from these, like volume in m³, density in kg/m³ and velocity in m/s.

When you measure with a scale or a measuring tape, the instrument has a least count and the last digit you note is always a bit doubtful. That doubtful digit is important — it tells us about significant figures, which show how precise your measurement is. All non-zero digits are significant, zeros between non-zero digits are significant, leading zeros are not significant, and trailing zeros after a decimal point are significant. For example, 0.0025 has 2 significant figures while 2.50 has 3.

While calculating, we must respect this precision. In addition and subtraction, the answer should have the same number of decimal places as the term with the least decimal places. In multiplication and division, the answer should have the same number of significant figures as the term with the least significant figures.

Have you ever converted kilograms to grams while buying vegetables? That method is called dimensional analysis, also known as the factor-label method. You simply multiply by a conversion factor in such a way that the old unit cancels out. For example, 1 kg = 1000 g, so 2 kg × (1000 g / 1 kg) = 2000 g. Units cancel just like numbers do.

Before we knew about atoms, scientists discovered five important laws of chemical combination by weighing things carefully:

1. Law of Conservation of Mass (Lavoisier): In any chemical reaction, mass is neither created nor destroyed. Total mass of reactants = total mass of products. For example, 2 g of H₂ + 16 g of O₂ gives 18 g of H₂O.

2. Law of Definite Proportions (Proust): A pure compound always contains the same elements in the same fixed mass ratio, no matter where it comes from. Water always has H:O = 1:8 by mass — 11.2% H and 88.8% O.

3. Law of Multiple Proportions (Dalton): When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a small whole-number ratio. For CO and CO₂, with 12 g of carbon, oxygen masses are 16 g and 32 g — ratio 1:2.

4. Law of Gaseous Volumes (Gay Lussac): When gases react, their volumes are in a simple whole-number ratio at same T and P. 1 volume N₂ + 3 volumes H₂ → 2 volumes NH₃, ratio 1:3:2.

5. Avogadro's Law: Equal volumes of all gases at the same temperature and pressure contain equal number of molecules. One litre of H₂ and one litre of O₂ have the same number of molecules under the same conditions.

To explain those laws, John Dalton in 1808 gave the first atomic theory. Think of it as a set of rules for how matter is built.

His postulates were: 1. All matter is made of tiny indivisible particles called atoms. 2. All atoms of a given element have identical mass and properties, while atoms of different elements have different mass. 3. Compounds are formed when atoms combine in simple whole-number ratios. 4. A chemical reaction is just a rearrangement of atoms — atoms are neither created nor destroyed.

We now know two postulates need correction. Atoms are divisible into electrons, protons and neutrons, and atoms of the same element can have different masses — these are called isotopes. However, Dalton's core idea that reactions are rearrangements is still correct.

Atomic mass is the mass of a single atom. We use carbon-12 as the standard — one atom of C-12 is taken as exactly 12 amu. One amu is 1/12th the mass of a C-12 atom, equal to 1.66 × 10⁻²⁷ kg. Since most elements have isotopes, we use average atomic mass weighted by abundance — chlorine is 35.5 because it has both Cl-35 and Cl-37. Molecular mass is the sum of atomic masses in a molecule, like H₂O = 2×1 + 16 = 18 amu. For ionic compounds like NaCl, we say formula mass = 23 + 35.5 = 58.5 amu.

Counting atoms one by one is impossible, so chemists count them in bundles, just like we count eggs in dozens. While a dozen means 12, a mole means 6.022 × 10²³ particles. This huge number is called Avogadro’s number (Nₐ).

So, 1 mole of atoms contains 6.022 × 10²³ atoms, 1 mole of molecules contains 6.022 × 10²³ molecules, and 1 mole of ions contains 6.022 × 10²³ ions. It is simply a counting unit, but for extremely small particles.

The beauty is that the mass of one mole in grams is numerically equal to its atomic or molecular mass in amu. This is called molar mass. Carbon has atomic mass 12 amu, so 1 mole of carbon is 12 g and contains Nₐ atoms. Water has molecular mass 18 amu, so 1 mole of water is 18 g and contains Nₐ H₂O molecules.

This gives two very useful relations you will use again and again: moles = given mass / molar mass and number of particles = moles × Nₐ. With these, you can easily jump between grams, moles and actual number of atoms or molecules.

When you read on a packet that sugar is C₁₂H₂₂O₁₁, how do we know how much carbon it has? This is given by percentage composition — mass percent of an element = (mass of that element in one mole of compound / molar mass of compound) × 100. For water, H% = (2/18)×100 = 11.1% and O% = (16/18)×100 = 88.9%.

Next, we often need the formula of a compound. The empirical formula gives the simplest whole-number ratio of atoms, while the molecular formula gives the actual number. For example, CH₂O is empirical for glucose, while C₆H₁₂O₆ is its true molecular formula, where n = 6.

To find the empirical formula from percentage data, use a simple kitchen logic: assume 100 g of compound so percentage becomes grams, then convert grams → moles by dividing with atomic mass, then divide all moles by the smallest mole value to get ratio, and if needed multiply by 2 or 3 to make whole numbers. That ratio is your empirical formula.

Finally, molecular formula = n × empirical formula, where n = molar mass of compound / molar mass of empirical formula. This tells us how many times the simple ratio repeats in the real molecule.

Balancing a chemical equation is like writing a recipe — it tells you how many spoons of each ingredient you need. Stoichiometry is the calculation of amounts of reactants and products using the balanced equation, which directly gives the mole ratio.

For example, 2H₂ + O₂ → 2H₂O means 2 moles of H₂ need 1 mole of O₂ to produce 2 moles of H₂O. If you know moles of one substance, you can find the rest using this ratio.

In real kitchens and labs, one ingredient finishes first. The limiting reagent is the reactant that gets consumed first and stops the reaction; the others are left over as excess. Think of making sandwiches — if you have 10 bread slices and 5 cheese slices to make 1 bread + 1 cheese sandwich, you can make only 5 sandwiches. Cheese is limiting, 5 bread slices remain excess.

To find the limiting reagent, calculate how much product each reactant would give — whichever gives less product is the limiting one. For reactions in solution, we measure amounts using concentration and volume, but the same mole-ratio logic applies.