Some Basic Concepts of Chemistry: common doubts, answered
The questions students ask most often about Some Basic Concepts of Chemistry, each with a short answer. For the full chapter, read the Some Basic Concepts of Chemistry notes.
About the chapter
How is Some Basic Concepts of Chemistry usually tested in NEET?
Mostly through short numericals: moving between mass, moles and number of particles, finding the limiting reagent and the amount of product, working out empirical and molecular formulas, and converting between molarity, molality and mole fraction. Conceptual questions on significant figures and the laws of combination appear too. Mole arithmetic from this chapter is reused in equilibrium, solutions and electrochemistry, so speed here pays off later.
Classifying matter
Read this section in the notes →What is the difference between a mixture and a compound?
A compound has its elements chemically joined in a fixed mass ratio and shows properties unlike those of the elements, while a mixture holds its components in any proportion and each keeps its own properties. That is why a mixture can be separated by physical means such as filtration, evaporation or distillation, whereas a compound has to undergo a chemical reaction to give back its elements.
What is the difference between a homogeneous and a heterogeneous mixture?
A homogeneous mixture has the same composition at every point, so no boundary between its components can be seen; air and sugar dissolved in water are examples. A heterogeneous mixture changes from one region to another, and its parts can often be picked out, as with salt stirred into iron filings or sand in water. Both are still mixtures and both separate by physical methods.
Properties, SI units and measurement
Read this section in the notes →What are the seven SI base units in chemistry?
They are the metre for length, kilogram for mass, second for time, ampere for electric current, kelvin for temperature, mole for amount of substance and candela for luminous intensity. Every other unit a chemist uses, such as the joule, pascal or mol L⁻¹, is built from these. Notice that the base unit of mass is the kilogram, not the gram, which matters when you check units in a formula.
How do you convert Celsius to kelvin and Fahrenheit?
Add 273.15 to a Celsius reading to get kelvin, and use °F = (9/5)(°C) + 32 for Fahrenheit, so 25 °C becomes 298.15 K and 77 °F. The kelvin scale begins at absolute zero and therefore never goes negative, unlike the other two. Any formula that contains T, such as the gas equation or Gibbs energy, needs the temperature in kelvin.
Scientific notation, precision and accuracy
Read this section in the notes →What is the difference between precision and accuracy?
Precision describes how closely repeated readings agree with one another, while accuracy describes how close they are to the true value. If a 2.00 g sample is weighed three times as 1.95, 1.94 and 1.95 g, the readings are precise but not accurate. Scattered readings that happen to average 2.00 g are accurate but not precise. Reliable data needs both qualities at once.
Significant figures and dimensional analysis
Read this section in the notes →How many significant figures are there in 0.0050600?
There are five significant figures: 5, 0, 6, 0 and 0. The zeros in front of the 5 only fix the position of the decimal point, so they are never counted. The zero between 5 and 6 sits between non-zero digits and counts, and the two zeros at the end count because the number has a decimal point, which shows they were measured.
What is the rule for significant figures in addition versus multiplication?
In multiplication and division the result keeps as many significant figures as the least precise number used; in addition and subtraction it keeps as many decimal places as the number with the fewest. So 2.5 × 1.25 is reported as 3.1, while 12.11 + 18.0 + 1.012 is reported as 31.1. A result can never claim more certainty than the weakest measurement behind it.
How does the unit factor method work in dimensional analysis?
You multiply a quantity by a fraction whose top and bottom are equal amounts in different units, so its value is one and only the unit changes. To turn 3 m into centimetres, multiply by 100 cm / 1 m and the metres cancel to leave 300 cm. Writing every unit and cancelling at each step catches a wrong conversion before it spoils the final answer.
Laws of chemical combination
Read this section in the notes →What is the difference between the law of definite proportions and the law of multiple proportions?
Definite proportions is about one compound: whatever its source, it contains the same elements in the same mass ratio. Multiple proportions is about two or more compounds made from the same pair of elements: if you hold one element's mass constant, the amounts of the second element across those compounds compare as small whole numbers. Water and hydrogen peroxide hold oxygen in a 1 : 2 ratio for the same hydrogen.
Why does mass seem to decrease when a candle burns if mass is conserved?
The mass does not disappear; it leaves as gases. Burning wax forms carbon dioxide and water vapour, which drift away from the flame, so the remaining candle weighs less. If the same burning were carried out in a sealed container, the total mass of everything inside would stay exactly the same before and after, which is what the law of conservation of mass states.
Dalton's atomic theory
Read this section in the notes →Why could Dalton's atomic theory not explain Gay Lussac's law of gaseous volumes?
Dalton had no idea of molecules as distinct from atoms, so he could not explain why 2 volumes of hydrogen and 1 volume of oxygen give 2 volumes of water vapour. Avogadro resolved it by proposing that equal volumes of gases at the same temperature and pressure hold equal numbers of molecules, and that hydrogen and oxygen exist as two-atom molecules that split during the reaction.
Atomic, molecular and formula masses
Read this section in the notes →Why do we say formula mass for NaCl instead of molecular mass?
Solid sodium chloride contains no separate molecules. It is a continuous lattice in which every Na⁺ ion is surrounded by Cl⁻ ions and every Cl⁻ by Na⁺ ions, so NaCl only records the simplest ratio of ions. The mass of that unit, 58.5 u, is therefore called the formula mass, and molecular mass is reserved for substances built from discrete molecules such as H₂O.
The mole and molar mass
Read this section in the notes →Why is the atomic mass in u numerically equal to the molar mass in g per mol?
Because the Avogadro constant links the two units exactly. One atomic mass unit is 1.66056 × 10⁻²⁴ g, and 6.022 × 10²³ of them add up to one gram. So an oxygen atom of 16 u gives a mole of oxygen atoms weighing 16 g. The number is the same; only the meaning changes, from the mass of one particle to the mass of a mole of particles.
How do you find the number of atoms rather than molecules in a mole problem?
Convert to moles of the substance, multiply by the Avogadro constant to get molecules, then multiply by how many of that atom each formula carries. For example, 0.5 mol of NH₃ holds 3.011 × 10²³ molecules but 1.5 mol, or about 9.03 × 10²³, hydrogen atoms. Most errors here come from stopping at molecules when the question asks about one element.
Percentage composition, empirical and molecular formula
Read this section in the notes →How do you calculate an empirical formula from percentage composition?
Take 100 g of the compound so each percentage becomes a mass in grams, divide each by its atomic mass to get moles, then divide every mole value by the smallest one. If a ratio comes out near 1.5 or 1.33, multiply all of them by 2 or 3 to reach whole numbers. Rounding 1.5 up to 2 is the usual mistake that gives a wrong formula.
How do you get the molecular formula from the empirical formula?
Divide the molar mass by the empirical formula mass to find a whole number n, then multiply every subscript in the empirical formula by n. Glucose has the empirical formula CH₂O with a mass of 30 g mol⁻¹; its molar mass is 180 g mol⁻¹, so n = 6 and the molecular formula is C₆H₁₂O₆. Without the molar mass you cannot go beyond the empirical formula.
Stoichiometry and the limiting reagent
Read this section in the notes →How do you find the limiting reagent in a chemical reaction?
Divide the moles of each reactant by its coefficient in the balanced equation; the reactant with the smallest value is the limiting reagent. It is consumed first and decides how much product can form, while the others remain partly unused. Comparing masses directly misleads, because a heavier sample may contain fewer moles and the equation may need several moles of one reactant for each mole of another.
Why does NCERT take 22.7 L as the molar volume of a gas instead of 22.4 L?
The two values belong to different standard pressures. At 273.15 K and 1 bar, the convention NCERT follows, one mole of an ideal gas fills 22.7 L; at 273.15 K and 1 atm, which is slightly higher pressure, it shrinks to 22.4 L. Neither number is wrong, so check the conditions stated in a question and use the volume that matches them.
Reactions in solutions: concentration terms
Read this section in the notes →What is the difference between molarity and molality?
Molarity is moles of solute per litre of solution, while molality is moles of solute per kilogram of solvent. Since molarity depends on a volume, it shifts slightly as the solution expands or contracts with temperature; molality uses only masses, so it stays constant. This is why molality is preferred when a solution is heated or cooled, as in boiling and freezing point problems.
How do you convert molarity into molality using density?
Consider one litre of solution. Its mass is density × 1000 g; subtract the solute's mass, which is molarity × molar mass, to get the mass of solvent. Then divide the moles of solute by that solvent mass in kilograms. The common error is dividing by the mass of the whole solution, but molality counts only the solvent, so the solute's mass has to be removed first.
What is mole fraction and why do all the mole fractions add up to one?
The mole fraction of a component is its number of moles divided by the total moles of every component present. Since the parts together make up the whole, the fractions must sum to 1. A solution of 1 mol solute in 9 mol water has a solute mole fraction of 0.1 and water 0.9. Mole fraction has no unit and does not change with temperature.
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