Chemistry: Class 11 — Chapter 1: Some Basic Concepts of Chemistry

• In ancient India, chemistry was called Rasayan Shastra, Rastantra, Ras Kriya, or Rasvidya. It included metallurgy, medicine, cosmetics, glass, and dyes.
• Systematic excavations at Mohenjodaro and Harappa show early chemical processes like baked bricks, glazed pottery, gypsum cement (containing lime, sand, and traces of $CaCO_3$), and metal forging (lead, silver, gold, copper, tin, arsenic).
• Sushruta Samhita explains the importance of alkalies. Charaka Samhita mentions the preparation of sulphuric acid, nitric acid, oxides of copper, tin, zinc, and carbonates of lead and iron.
• Rasopanishada describes the preparation of gunpowder mixtures.
• Nagarjuna was a reputed chemist, alchemist, and metallurgist whose work Rasratnakar deals with mercury compounds and extraction of metals.
• Chakrapani discovered mercury sulphide and is credited with inventing soap using mustard oil and alkalies.
• Acharya Kanda (born 600 BCE) formulated the atomic theory of matter, naming the indivisible particles Paramānu (comparable to atoms) centuries before John Dalton.

Chemistry plays a central role in science and is applicable in diverse areas such as weather patterns, brain functioning, computer operations, and chemical industries.

• Provides life-saving drugs: cisplatin and taxol (effective in cancer therapy), and AZT (Azidothymidine) used for AIDS patients.
• Assists in environmental management: Synthesizing safer alternatives to environmentally hazardous refrigerants like CFCs (chlorofluorocarbons) to protect the ozone layer.

• Solids: Have definite volume and definite shape. Particles are held very close to each other in an orderly fashion.
• Liquids: Have definite volume but no definite shape; they take the shape of the container.
• Gases: Have neither definite volume nor definite shape. They completely occupy the container.
• These states are interconvertible: $$\text{Solid} \underset{\text{cool}}{\overset{\text{heat}}{\rightleftharpoons}} \text{Liquid} \underset{\text{cool}}{\overset{\text{heat}}{\rightleftharpoons}} \text{Gas}$$

• Mixtures: Contain two or more substances in any ratio. Their composition is variable.
  • Homogeneous Mixture: Components completely mix and the composition is uniform throughout (e.g., sugar solution, air).
  • Heterogeneous Mixture: Composition is not uniform and different components are visible (e.g., mixture of salt and sugar, grains and pulses with dirt).
• Pure Substances: Have a fixed composition.
  • Elements: Consist of only one type of atoms or molecules (e.g., Sodium, Copper, Hydrogen gas $H_2$).
  • Compounds: Formed when two or more atoms of different elements combine in a fixed ratio (e.g., Water $H_2O$, Carbon dioxide $CO_2$). The properties of a compound are different from its constituent elements.

• Physical Properties: Can be measured or observed without changing the identity or composition of the substance (e.g., color, odor, melting point, boiling point, density).
• Chemical Properties: Require a chemical change to occur for measurement (e.g., acidity, basicity, combustibility).

The International System of Units (SI) defines seven base units:

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

• Mass vs Weight: Mass is the amount of matter present in a substance (constant). Weight is the force exerted by gravity on an object (variable).
• Density: Mass per unit volume. $$\text{Density} = \frac{\text{Mass}}{\text{Volume}}$$ SI unit: $kg\,m^{-3}$ (often expressed as $g\,cm^{-3}$ in laboratories).
• Temperature Scales: Celsius ($^\circ C$), Fahrenheit ($^\circ F$), and Kelvin ($K$).
  • $$^\circ F = \frac{9}{5}(^\circ C) + 32$$
  • $$K = ^\circ C + 273.15$$

Exponential notation represented as $N \times 10^n$, where $n$ is an exponent and $N$ is a digit term between $1.000...$ and $9.999...$

• Example: $232.508$ is written as $2.32508 \times 10^2$.
• Example: $0.00016$ is written as $1.6 \times 10^{-4}$.

Significant figures are meaningful digits known with certainty plus one that is estimated or uncertain.

• Rules:
  1. All non-zero digits are significant.
  2. Zeros preceding the first non-zero digit are not significant (e.g., $0.03$ has one significant figure).
  3. Zeros between two non-zero digits are significant (e.g., $2.005$ has four).
  4. Zeros at the end or right of a number are significant if they are on the right side of the decimal point (e.g., $0.200\,g$ has three). Terminal zeros without a decimal point are not significant (e.g., $100$ has one).
  5. Exact numbers have an infinite number of significant figures (e.g., $2$ balls = $2.000000...$).
• Rounding Off Rules:
  • If the rightmost digit to be removed is $> 5$, the preceding number is increased by $1$.
  • If it is $< 5$, the preceding number is unchanged.
  • If it is $= 5$, the preceding number is unchanged if even, and increased by $1$ if odd (e.g., $6.35 \rightarrow 6.4$; $6.25 \rightarrow 6.2$).

• Precision refers to the closeness of various measurements for the same quantity.
• Accuracy is the agreement of a particular value to the true value of the result.

1. Law of Conservation of Mass: Formulated by Antoine Lavoisier (1789). Matter can neither be created nor destroyed during physical or chemical changes.
2. Law of Definite Proportions: Proposed by Joseph Proust. A given compound always contains exactly the same proportion of elements by weight.
3. Law of Multiple Proportions: Proposed by John Dalton (1803). If two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers.
4. Gay Lussac’s Law of Gaseous Volumes: Proposed in 1808. When gases combine or are produced in a chemical reaction, they do so in a simple ratio by volume, provided all gases are at the same temperature and pressure.
5. Avogadro’s Law: Proposed in 1811. Equal volumes of all gases at the same temperature and pressure should contain an equal number of molecules.

• One atomic mass unit (amu) (now unified mass, u) is defined as a mass exactly equal to $1/12^{\text{th}}$ of the mass of one carbon-12 atom. $$1\,\text{amu} = 1.66056 \times 10^{-24}\,g$$
• Average Atomic Mass is calculated by taking into account the natural abundance of different isotopes of an element. For Carbon: $$(0.98892)(12\,u) + (0.01108)(13.00335\,u) + (2 \times 10^{-12})(14.00317\,u) = 12.011\,u$$

• Molecular Mass is the sum of atomic masses of the elements present in a molecule. For methane ($CH_4$): $$12.011\,u + 4(1.008\,u) = 16.043\,u$$
• Formula Mass is used for ionic substances (like $NaCl$) that do not contain discrete molecules. For $NaCl$: $$23.0\,u + 35.5\,u = 58.5\,u$$

$$\text{Mass % of an element} = \frac{\text{Mass of that element in the compound} \times 100}{\text{Molar mass of the compound}}$$

• An empirical formula represents the simplest whole-number ratio of various atoms present in a compound.
• A molecular formula shows the exact number of different types of atoms present in a molecule of a compound.
• $$\text{Molecular Formula} = n \times \text{Empirical Formula}$$ where $$n = \frac{\text{Molar Mass}}{\text{Empirical Formula Mass}}$$

Stoichiometry deals with the calculation of masses (and volumes) of reactants and products involved in a chemical reaction. Example balanced equation:

CH4 (g) + 2O2 (g) -> CO2 (g) + 2H2O (g)

The reactant that is completely consumed first in a reaction, limiting the amount of product formed, is called the limiting reagent.

The concentration of a solution can be expressed in several ways:

1. Mass Percent: $$\text{Mass %} = \frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100$$
2. Mole Fraction ($x$): Ratio of the number of moles of a particular component to the total number of moles of the solution. $$x_A = \frac{n_A}{n_A + n_B}$$
3. Molarity ($M$): Number of moles of solute in $1$ litre of solution. $$\text{Molarity (M)} = \frac{\text{No. of moles of solute}}{\text{Volume of solution in litres}}$$
4. Molality ($m$): Number of moles of solute present in $1\,kg$ of solvent. $$\text{Molality (m)} = \frac{\text{No. of moles of solute}}{\text{Mass of solvent in kg}}$$