Some Basic Concepts of Chemistry: NEET notes.

01 · NCERT § "Nature of Matter"

Classifying matter

• Whatever has mass and takes up room counts as matter; it exists as solids (definite shape and volume), liquids (definite volume, no definite shape) and gases (neither definite shape nor volume).
• The three states can be interconverted by changing temperature and pressure; heating a solid gives a liquid, and heating a liquid gives a gas.
• At the macroscopic level matter is split into mixtures and pure substances.
• A homogeneous mixture has the same composition throughout, such as sugar dissolved in water or air; a heterogeneous mixture does not, such as a mixture of salt and sand.
• Components of a mixture keep their own properties and can be separated by physical methods like filtration, crystallisation or distillation.
• Pure substances are either elements, made of only one kind of atom, or compounds, in which atoms of two or more elements are joined in a fixed ratio.
• A compound's properties differ from those of its constituent elements, and it can be broken into its elements only by chemical methods; water, for example, is unlike hydrogen or oxygen.

02 · NCERT § "Properties of Matter and their Measurement"

Properties, SI units and measurement

• Physical properties such as colour, melting point, boiling point and density can be observed without changing the identity of the substance; chemical properties such as acidity or combustibility show up only when a chemical change happens.
• The SI system has seven base quantities: length (metre, m), mass (kilogram, kg), time (second, s), electric current (ampere, A), thermodynamic temperature (kelvin, K), amount of substance (mole, mol) and luminous intensity (candela, cd).
• Mass is the amount of matter in a body and stays the same everywhere; weight is the force of gravity on it and can change from place to place.
• Volume has the SI unit m³; in the laboratory 1 L = 1000 mL = 1000 cm³ = 1 dm³.
• Density is mass per unit volume; its SI unit is kg m⁻³, though g cm⁻³ is more common in the laboratory.
• Temperature is reported in °C, °F or K; the kelvin scale has no negative values because 0 K is the lowest possible temperature.
• Conversions: K = °C + 273.15 and °F = (9/5)(°C) + 32.

03 · NCERT § "Uncertainty in Measurement"

Scientific notation, precision and accuracy

• Very large or very small numbers are written in scientific (exponential) notation as N × 10ⁿ, where N lies between 1.000 and 9.999 and n is an integer.
• When multiplying numbers in this form, the N parts are multiplied and the exponents are added; when dividing, the exponents are subtracted.
• For addition or subtraction, first rewrite the numbers with the same exponent, then combine the N parts.
• Precision describes how close repeated measurements of the same quantity are to one another.
• Accuracy describes how close a measurement (or the average of several) is to the true value.
• A set of readings can be precise without being accurate, for example when an instrument has a constant error.

04 · NCERT § "Uncertainty in Measurement"

Significant figures and dimensional analysis

• Significant figures are the digits known with certainty plus one last estimated digit.
• All non-zero digits count; zeros between non-zero digits count; zeros to the left of the first non-zero digit never count (0.0032 has two).
• Zeros at the right end count only when there is a decimal point in the number: 0.200 has three, while 100 written without a decimal point has one.
• Exact numbers, such as a counted 2 balls or a defined conversion factor, have an unlimited number of significant figures.
• In addition and subtraction the answer keeps as many decimal places as the entry with the fewest decimal places.
• In multiplication and division the answer keeps as many significant figures as the entry with the fewest significant figures.
• Rounding: drop the digit if it is less than 5, raise the previous digit by one if it is more than 5; if it is exactly 5, keep the previous digit if it is even and raise it by one if it is odd.
• Dimensional analysis (the unit factor method) converts units by multiplying by a ratio equal to 1, such as (1 m / 100 cm), so the unwanted units cancel.

05 · NCERT § "Laws of Chemical Combinations"

Laws of chemical combination

• Law of conservation of mass (Antoine Lavoisier): in any physical or chemical change, matter is neither created nor destroyed, so total mass stays the same.
• Law of definite proportions (Joseph Proust): a given compound always contains the same elements combined in the same proportion by mass, whatever its source.
• Law of multiple proportions (John Dalton): if the same two elements make two or more different compounds, then fixing the mass of one element and comparing the masses of the other element across those compounds gives a small whole-number ratio.
• Example of multiple proportions: 2 g of hydrogen combines with 16 g of oxygen in water and with 32 g of oxygen in hydrogen peroxide, a 1 : 2 ratio.
• Gay Lussac's law of gaseous volumes: when gases react or are produced, their volumes, measured at the same temperature and pressure, bear a simple whole-number ratio.
• Gay Lussac's law is a statement about volumes; the law of definite proportions is a statement about masses.
• Avogadro's law: at a fixed temperature and pressure, a given volume of any gas holds the same number of molecules, whatever the gas.
• Avogadro separated the ideas of atoms and molecules, which explained results like 2 volumes of hydrogen plus 1 volume of oxygen giving 2 volumes of water vapour.

06 · NCERT § "Dalton's Atomic Theory"

Dalton's atomic theory

• Dalton proposed that matter is made of indivisible atoms.
• Every atom of one element is the same as every other, mass included, while atoms of two different elements have different masses.
• A compound arises when atoms of two or more elements join in a set, unchanging ratio.
• A reaction only reshuffles atoms into new partners; no atom is made or lost.
• The theory explained the laws of conservation of mass, definite proportions and multiple proportions.
• It could not explain the law of gaseous volumes, and it gave no reason why atoms combine at all.

07 · NCERT § "Atomic and Molecular Masses"

Atomic, molecular and formula masses

• Atomic masses are measured relative to the carbon-12 isotope, whose mass is fixed at exactly 12 atomic mass units.
• One atomic mass unit equals one-twelfth of a carbon-12 atom's mass, 1.66056 × 10⁻²⁴ g; it is now written u, the unified mass.
• Most elements occur as a mixture of isotopes, so the atomic mass used in calculations is an average weighted by the natural abundance of each isotope; this is why carbon is taken as 12.011 u.
• Molecular mass is the sum of the atomic masses of all atoms in a molecule, e.g. CH₄ = 12.011 + 4(1.008) = 16.043 u.
• Ionic solids such as NaCl are three-dimensional arrays of ions, not discrete molecules, so their formula mass is used instead: NaCl = 23.0 + 35.5 = 58.5 u.
• Isotope-averaging is why most tabulated atomic masses are not whole numbers.

08 · NCERT § "Mole Concept and Molar Masses"

The mole and molar mass

• The mole is the SI unit for amount of substance; one mole of anything contains 6.022 × 10²³ elementary entities (atoms, molecules, ions or formula units).
• This number is the Avogadro constant, Nᴀ = 6.022 × 10²³ mol⁻¹.
• Molar mass is how many grams one mole of a substance weighs; numerically it equals the atomic, molecular or formula mass in u (water: 18.02 u per molecule, 18.02 g mol⁻¹).
• Number of moles n = given mass / molar mass; number of particles = n × Nᴀ.
• Always check which particle is asked for: one mole of NH₃ has Nᴀ molecules but 3Nᴀ hydrogen atoms and 4Nᴀ atoms in total.
• For gases, NCERT's stoichiometry example takes one mole as 22.7 L, the molar volume of an ideal gas at 273.15 K and 1 bar; the older figure of 22.4 L belongs to 273.15 K and 1 atm.

09 · NCERT § "Percentage Composition"

Percentage composition, empirical and molecular formula

• Mass percent of an element = (its mass in one mole of the compound / molar mass of the compound) × 100.
• The empirical formula gives the simplest whole-number ratio of atoms; the molecular formula gives the actual number of each kind of atom in a molecule.
• Route from percentages to empirical formula: assume 100 g, convert each mass to moles, divide all by the smallest mole value, then multiply by a small integer if needed to reach whole numbers.
• Molecular formula = n × (empirical formula), where n = molar mass / empirical formula mass, and n is a whole number.
• Glucose (C₆H₁₂O₆) and ethanoic acid (C₂H₄O₂) both have the empirical formula CH₂O, so the empirical formula alone cannot identify a compound.
• Background, not stated in this NCERT chapter: if a problem gives vapour density relative to hydrogen, molar mass is about twice that value, a result that follows from Avogadro's law.

10 · NCERT § "Stoichiometry and Stoichiometric Calculations"

Stoichiometry and the limiting reagent

• Stoichiometry is the calculation of the masses and, for gases, volumes of reactants and products from a balanced equation.
• The coefficients in a balanced equation give mole ratios (and, for gases at the same temperature and pressure, volume ratios), not mass ratios.
• Balancing an equation means making the count of every element match on the left and right sides; this is conservation of mass written in symbols.
• The limiting reagent is the reactant that is used up first; it alone decides how much product forms, and the other reactant is left in excess.
• To find it, divide the available moles of each reactant by its coefficient; the smallest quotient marks the limiting reagent.
• When a sample is impure, first find the mass of the pure reacting substance from its percentage purity, then do the mole calculation.

11 · NCERT § "Reactions in Solutions"

Reactions in solutions: concentration terms

• Mass percent = (mass of solute / mass of solution) × 100.
• Mole fraction of a component = its moles / total moles of all components; the mole fractions of all components add up to 1.
• Molarity (M) = moles of solute per litre of solution; it changes with temperature because volume changes.
• Molality (m) = moles of solute per kilogram of solvent; it does not depend on temperature because it uses masses only.
• Dilution keeps the moles of solute fixed, so M₁V₁ = M₂V₂.
• Converting molarity to molality needs the solution's density: find the mass of 1 L of solution, subtract the solute's mass to get the solvent's mass, then divide the moles of solute by that mass in kg.
• Converting molality to molarity goes the other way: mass of solution = mass of solvent + mass of solute, then use density to get the volume.

Revise

Must-know facts

1. 01Seven SI base units: m, kg, s, A, K, mol, cd.
2. 02K = °C + 273.15; °F = (9/5)(°C) + 32.
3. 031 L = 1000 mL = 1000 cm³ = 1 dm³.
4. 04Leading zeros are never significant; trailing zeros are significant only when a decimal point is present.
5. 05Multiplication/division: fewest significant figures; addition/subtraction: fewest decimal places.
6. 06Conservation of mass: Lavoisier; definite proportions: Proust; multiple proportions: Dalton; gaseous volumes: Gay Lussac; equal volumes, equal molecules: Avogadro.
7. 07Dalton's theory could not explain Gay Lussac's law of gaseous volumes.
8. 081 u = 1/12 of the mass of a carbon-12 atom = 1.66056 × 10⁻²⁴ g.
9. 09Nᴀ = 6.022 × 10²³ mol⁻¹.
10. 10Molar mass in g mol⁻¹ is numerically equal to the molecular or formula mass in u.
11. 11Ionic compounds have formula mass, not molecular mass (NaCl = 58.5 u).
12. 12One mole of an ideal gas occupies 22.7 L at 273.15 K and 1 bar.
13. 13Empirical formula of glucose is CH₂O.
14. 14Molecular formula = n × empirical formula, with n = molar mass / empirical formula mass.
15. 15The limiting reagent decides the amount of product, not the reactant present in larger mass.
16. 16Molality is temperature-independent; molarity is not.
17. 17Sum of mole fractions of all components = 1.
18. 18Dilution: M₁V₁ = M₂V₂.

Where marks go

Common traps

Reference

Formulas

Kelvin and Fahrenheit conversion

K = °C + 273.15 ; °F = (9/5)(°C) + 32

No negative temperatures exist on the kelvin scale.

Density

d = m / V

SI unit kg m⁻³; commonly g cm⁻³ or g mL⁻¹.

Number of moles

n = mass / molar mass = N / Nᴀ

Nᴀ = 6.022 × 10²³ mol⁻¹; molar mass in g mol⁻¹.

Moles of a gas at STP

n = V / 22.7 L

STP taken as 273.15 K and 1 bar.

Average atomic mass

Aᵥ = Σ (fractional abundance × isotopic mass)

Explains non-integer atomic masses such as 12.011 u for carbon.

Mass percent of an element

% element = (mass of element in 1 mol compound / molar mass) × 100

Percentages of all elements add up to 100.

Molecular formula multiplier

n = molar mass / empirical formula mass

Background, not in this NCERT chapter: molar mass ≈ 2 × vapour density (relative to H₂).

Mole fraction

x_A = n_A / (n_A + n_B + ...)

Dimensionless; all fractions sum to 1.

Molarity

M = moles of solute / volume of solution in L

Unit mol L⁻¹; temperature dependent.

Molality

m = moles of solute / mass of solvent in kg

Unit mol kg⁻¹; temperature independent.

Dilution

M₁V₁ = M₂V₂

Moles of solute stay constant on dilution.

Glossary

Key terms

Homogeneous mixture

A mixture whose composition is uniform throughout, like a salt solution.

Compound

A pure substance made of two or more elements in a fixed mass ratio, separable only by chemical means.

Precision

Agreement of repeated measurements with each other.

Accuracy

Agreement of a measurement with the true value.

Significant figures

The certain digits of a measurement plus one estimated digit.

Unified mass (u)

One-twelfth of the mass of a carbon-12 atom, used to express atomic and molecular masses.

Formula mass

The sum of atomic masses in the formula unit of an ionic compound.

Mole

The amount of substance holding 6.022 × 10²³ specified entities.

Molar mass

Mass in grams of one mole of a substance.

Empirical formula

Formula showing the simplest whole-number atom ratio.

Limiting reagent

The reactant consumed first, which caps the amount of product.

Molality

Moles of solute per kilogram of solvent.

Molarity

Moles of solute per litre of solution.