Fundamental Laws of Chemistry & Nature of Matter

Fundamental Laws of Chemistry

• Three cornerstone relationships derived from 18th-century quantitative experiments:
• Law of Conservation of Mass (Lavoisier, 1774)
• Law of Definite Composition / Law of Definite Proportions (Proust)
• Law of Multiple Proportions (Dalton formalised; based on earlier data)
• Collectively, these laws forced chemists to abandon alchemy-style transmutation ideas and adopt an atom-based view of matter, laying the groundwork for modern atomic theory, stoichiometry, and the periodic table.

1. Law of Conservation of Mass

• Statement: In any chemical reaction the total mass of the products equals the total mass of the reactants.
⇒ Mass is conserved; it is neither created nor destroyed during chemical changes.
• Historical context: Antoine Lavoisier used closed vessels and precise balances to show that combustion, calcination and other reactions obey the equality of mass.
• Practical significance:
• Foundation of stoichiometric calculations.
– Balancing equations assumes the same number of atoms (hence same total mass) before & after reaction.
– Allows chemists to predict product yield or required reactant quantities.
• Introduced quantitative rigour into chemistry → transition from qualitative alchemy.
• Conceptual link: Conservation of mass later merged with conservation of energy in Einstein’s $E = mc^2$; however, for most laboratory-scale reactions the mass change from energy release is negligible.
• Late-1700s consensus on “element”: a substance that cannot be decomposed by ordinary chemical means (matches modern chemical element definition).

2. Law of Definite Composition (Proust’s Law / Law of Constant Composition)

• Statement: A given chemical compound always contains the same elements combined in the same fixed mass ratio.
• Implications:
• Reinforces the concept that compounds are built from discrete, whole-number combinations of atoms (Dalton’s later postulate).
• Distinguishes compounds from mixtures (mixtures can vary continuously in composition).
• Example – Water $\left(\text{H}<em>2\text{O}\right)$ • Atomic masses: H = $1\ \text{g·mol}^{-1}$, O = $16\ \text{g·mol}^{-1}$ • Molecular mass calculation: $M</em>{\text{H}<em>2\text{O}} = 2(1) + 1(16) = 18\ \text{g·mol}^{-1}$ • Percent composition: $\%\,\text{H} = \frac{2}{18}\times 100\% = 11\%$ $\%\,\text{O} = \frac{16}{18}\times 100\% = 89\%$ • Therefore only an 11 : 89 mass ratio yields water. Any mixture with 30 % H and 70 % O will not be water; it will be an explosive mixture of H$2$ and O$_2$ instead.
• Terminology:
• Early name “definite proportions” emphasises fixedness; modern “constant composition” highlights invariance across samples.
• Laboratory relevance: empirical-formula determination via combustion or gravimetric analysis relies on this principle.

3. Law of Multiple Proportions

• Statement: If two elements form more than one compound, the different masses of one element that combine with a fixed mass of the other are in very small whole-number ratios (1 : 1, 2 : 1, 3 : 2, …).
– John Dalton (1803) made this observation central to his atomic theory.
• Demonstration – Oxides of Nitrogen (N fixed at $14\ \text{g·mol}^{-1}$):
• $\text{NO}:\ 14\text{ g N} : 16\text{ g O}\Rightarrow 1:1$
• $\text{NO}<em>2: 14 : 32\;\Rightarrow 1:2$ • $\text{NO}</em>3: 14 : 48\;\Rightarrow 1:3$
• $\text{NO}<em>4: 14 : 64\;\Rightarrow 1:4$ • Sample Problem (given in transcript): – Two carbon oxides contain 42.9 % C and 27.3 % C respectively. Show they satisfy the law. – Choose $100\ \text{g}$ samples. • Compound A: $m</em>\text{C}=42.9\ \text{g},\; m<em>\text{O}=57.1\ \text{g}$ $\text{O per g C}=\frac{57.1}{42.9}=1.331$ • Compound B: $m</em>\text{C}=27.3\ \text{g},\; m_\text{O}=72.7\ \text{g}$
$\text{O per g C}=\frac{72.7}{27.3}=2.663$
• Ratio of O masses with same C mass:
$\frac{2.663}{1.331}\approx 2.00:1$ (a small whole number), confirming the law.
• Importance:
• Provides experimental evidence for atoms having integral combining capacities (“valences”).
• Basis for Dalton’s integer atomic weight scale and for chemical formula deduction.

Nature and States of Matter

• Matter exists principally in three common macroscopic phases: solid, liquid, gas (plasma & Bose–Einstein condensates discussed in advanced courses).
• Learning objectives (from slide):

1. Recognise that substances are made of tiny particles (atoms, molecules, ions).

2. Describe arrangement, spacing, motion of particles in each phase.

3. Identify phase changes (melting, freezing, vaporisation, condensation, sublimation, deposition).

Particle-Level Description

• Solids: closely packed particles in fixed positions; vibrational motion only; definite shape & volume.
• Liquids: particles close but can slide; definite volume, no definite shape; moderate KE.
• Gases: particles far apart; rapid, random motion; no definite shape or volume; compressible.
• Phase changes occur when energy (heat) is added or removed, altering particle kinetic energy and intermolecular forces.
• Example: melting requires energy input to overcome lattice energy; freezing releases energy.

Ancient Theories of Matter (Pre-Scientific Era)

Early Greek Philosophers & Elemental Ideas

• Anaximenes (c. 545 BCE): primal substance = air.
• Thales of Miletus (6th century BCE): primal substance = water.
• Heraclitus (c. 540–480 BCE): primal substance = fire.
• Empedocles (c. 490–430 BCE): combined prior views → four fundamental elements: earth, air, fire, water.
• Explained material diversity as varying proportions of these four.
• Limitation: upon decomposition, the four elements were not literally recovered, hinting at flaws.
• Indirectly foreshadowed the Law of Constant Proportion by stressing composition ratios.

Atomism – Leucippus & Democritus (~400 BCE)

• Thought experiment with endlessly cutting gold led to concept of atomos – indivisible units.
• Five key postulates:

1. Matter composed of tiny, indivisible atoms; variety of matter arises from atom combinations.

2. Atoms move constantly in the void (empty space).

3. Atoms are completely solid (billiard-ball model).

4. Atoms are uniform with no internal structure.

5. Atoms differ in size & shape.

Epicurus (341–270 BCE)

• Popularised atomism; added:
• Atoms possess different weights.
• All atoms travel at equal speeds regardless of size (conceptual, not physical accuracy).

Aristotelian Opposition & Continuous Theory

• Aristotle rejected atomism; embraced Empedoclean four-element model + fifth element Aether (heavenly matter).
• Argued elements could transmute (foundation of later alchemy).
• Championed Continuous Theory of Matter – substances divisible ad infinitum without changing nature.

Synthesis & Modern Perspective

• Early theories were philosophical, based on logic rather than experiment.
• Lack of instruments prevented verification → doctrines persisted for centuries.
• Rise of scientific method (16th–18th centuries): quantitative balances, gas collection apparatus, calorimetry, etc. → experimental scrutiny.
• Converging evidence (conservation, definite & multiple proportions) led Dalton (1803) to formalise Modern Atomic Theory.
• Postulate of indivisible atoms later modified (discovery of electrons, protons, neutrons, quarks).
• Present view: atoms are not indestructible; they contain subatomic particles and can undergo nuclear reactions, but chemical reactions obey the classical laws almost exactly.
• Impact: understanding atomic structure underpins modern chemistry, materials science, nanotechnology, pharmaceuticals, and emerging quantum technologies.