Some Basic Concepts of Chemistry

1. Mixtures

Contain two or more substances in any ratio.

• Homogeneous: Uniform (e.g., air, sugar solution)
• Heterogeneous: Non-uniform (e.g., sand & salt)

2. Pure Substances

Have a fixed composition.

• Elements: One type of particle (Na, O₂)
• Compounds: Two or more elements in fixed ratio (H₂O)

Physical: Measured without changing identity (color, melting pt).
Chemical: Require chemical change (acidity, combustibility).

7 Base SI Units

• Length: metre (m)
• Mass: kilogram (kg)
• Time: second (s)
• Electric current: ampere (A)
• Temperature: kelvin (K)
• Amount of substance: mole (mol)
• Luminous intensity: candela (cd)

Temperature Conversion

Significant Figures Rules:

• All non-zero digits are significant.
• Leading zeros are not significant (0.003 has 1).
• Zeros between non-zeros are significant.
• Trailing zeros are significant if after decimal (0.200 has 3).
• Exact numbers have infinite sig figs.

Accuracy vs Precision

Precision: Closeness of various measurements to each other.
Accuracy: Closeness of a measurement to the true value.

• Law of Conservation of Mass: Matter is neither created nor destroyed. Mass of reactants = Mass of products.
• Law of Definite Proportions: A compound always contains exactly the same proportion of elements by weight.
• Law of Multiple Proportions: When two elements form $>1$ compound, the masses of one combining with a fixed mass of the other are in simple whole number ratios.
• Gay Lussac’s Law: Gases combine/are produced in simple volume ratios at same T & P.
• Avogadro’s Law: Equal volumes of gases at same T & P contain equal number of molecules.

• Matter consists of indivisible atoms.
• All atoms of a given element have identical properties (including identical mass). Atoms of different elements differ in mass.
• Compounds are formed when atoms of different elements combine in a fixed ratio.
• Chemical reactions involve reorganization of atoms. These are neither created nor destroyed.

Atomic Mass Unit (amu)

Defined as a mass exactly equal to 1/12th the mass of one $^{12}\text{C}$ atom.

Average Atomic Mass

Calculated by taking into account fractional abundance of isotopes.

Vapour Density

One Mole is the amount of a substance containing as many particles as there are atoms in exactly $12\text{ g}$ of $^{12}\text{C}$.

Key Formulae

• Given mass ($w$): $n = \frac{w}{M}$
• Number of particles ($N$): $n = \frac{N}{N_A}$
• Gases at STP: $n = \frac{V(\text{in L})}{22.4}$

Empirical & Molecular Formula

• Empirical: Simplest whole number ratio of atoms.
• Molecular: Exact number of different types of atoms.
• Relation: Molecular Formula = $n \times$ (Empirical Formula)

Limiting Reagent (LR)

Reactant consumed first. Dictates product amount.
Tip: Divide given moles by stoichiometric coefficient. Lowest value = LR.

POAC

Principle of Atomic Conservation: Moles of atoms of an element remain conserved.

Percentage Yield

• Molarity ($M$): Moles of solute / Vol of sol in L (Temp dependent)
• Molality ($m$): Moles of solute / Mass of solvent in kg (Temp independent)
• Mole Fraction ($x$): $x_A = \frac{n_A}{n_A + n_B}$
• Normality ($N$): Gram eq. of solute / Vol of sol in L
• ppm: $\frac{\text{Mass of component}}{\text{Total mass}} \times 10^6$

Master Formula