Chemistry Experiments

J.J. Thomson's cathode ray tube experiment (1897) represents one of the most pivotal experiments in the history of atomic physics, leading to the discovery of the electron and fundamentally transforming our understanding of atomic structure. Conducted at the Cavendish Laboratory in Cambridge, this elegant experiment demonstrated that cathode rays consist of negatively charged particles much smaller than atoms, proving atoms are not indivisible as Dalton proposed but contain subatomic particles. The experimental apparatus consisted of a sealed glass tube with most air evacuated, containing two metal electrodes: a cathode (negative electrode) and an anode (positive electrode). When high voltage (several thousand volts) was applied, cathode rays traveled from cathode to anode, causing fluorescence where they struck the glass. Thomson introduced two crucial modifications: electric deflection plates to create a uniform electric field perpendicular to the ray path, and electromagnetic coils to produce a magnetic field also perpendicular to both ray path and electric field. By carefully balancing electric and magnetic forces on the cathode rays, Thomson determined the charge-to-mass ratio (e/m) of the particles. When only electric field was applied, rays deflected toward the positive plate, indicating negative charge. When only magnetic field was applied, rays deflected according to right-hand rule. By adjusting fields until rays showed no net deflection, Thomson calculated e/m ≈ 1.76 × 10¹¹ C/kg, about 1800 times larger than the hydrogen ion's e/m ratio, suggesting either much smaller mass or larger charge. Thomson's key observations included: cathode rays travel in straight lines (cast sharp shadows); they carry momentum (rotate paddle wheels); they produce heat when stopped (increase temperature); they cause fluorescence in certain materials; they are deflected by electric and magnetic fields; and their properties remain constant regardless of cathode material or residual gas. Most importantly, the e/m ratio remained constant, suggesting cathode rays consist of identical particles present in all matter. The experiment's implications were profound: atoms contain negatively charged particles (electrons) much smaller than atoms themselves; since atoms are electrically neutral, they must also contain positive charge; atoms are divisible, contradicting Dalton's indivisible atom concept. Thomson proposed the "plum pudding" model: atoms as spheres of uniform positive charge with embedded electrons like raisins in pudding. Modern interpretations recognize Thomson's experiment as the first to demonstrate subatomic particles and measure their properties quantitatively. While his plum pudding model was later superseded by Rutherford's nuclear model, the electron discovery remains fundamental. The experiment exemplifies elegant experimental design: controlling variables, making precise measurements, and drawing revolutionary conclusions from simple apparatus. Today, cathode ray principles find applications in television tubes, oscilloscopes, electron microscopes, and mass spectrometers. Thomson's methodology—combining electric and magnetic deflection—evolved into mass spectrometry techniques that identify compounds by mass-to-charge ratios. The experiment stands as a testament to how careful measurement and logical deduction can reveal nature's deepest secrets.

Robert Millikan's oil drop experiment (1909-1913) represents one of physics' most elegant and precise measurements, determining the elementary electric charge (e) and demonstrating charge quantization. Conducted at the University of Chicago with Harvey Fletcher, this experiment provided the first accurate measurement of the electron's charge and confirmed that electric charge exists in discrete units rather than continuous amounts, fundamentally supporting atomic theory and quantum mechanics. The experimental apparatus consisted of an atomizer spraying fine oil droplets between two horizontal metal plates connected to a variable high-voltage source. Millikan observed individual droplets through a microscope, measuring their terminal velocity under gravity alone (with electric field off) to determine mass from Stokes' law, then applying an electric field to balance gravitational force and suspend droplets motionless. By adjusting voltage until droplets hovered, Millikan calculated charge from the balance equation: qE = mg. Millikan's key innovation was observing individual droplets for extended periods, sometimes hours, as they captured or lost electrons due to natural ionizing radiation. He measured the same droplet with different charges, noting charge changes occurred in discrete jumps corresponding to integer multiples of a fundamental unit. This provided direct evidence for charge quantization—electric charge exists as integer multiples of e ≈ 1.602 × 10⁻¹⁹ coulombs. The experimental chamber was illuminated from the side, making droplets appear as bright stars against dark background. Millikan measured fall times over a fixed distance with stopwatch precision of 0.01 seconds. Temperature control was critical since viscosity depends on temperature; the apparatus included thermometers and sometimes heating coils. Air pressure was measured with barometers since Stokes' law includes viscosity term. Millikan's original 1913 paper reported e = (1.5924 ± 0.0017) × 10⁻¹⁹ C, remarkably close to modern value 1.602176634 × 10⁻¹⁹ C. He made approximately 58 series of observations over 60 consecutive days, discarding only those where he had doubts about measurements—a practice later criticized but which didn't significantly affect results. The experiment's precision stemmed from averaging many measurements and careful control of variables: temperature, pressure, electric field uniformity, and droplet selection. The experiment confirmed several important principles: electric charge is quantized; electron charge is the fundamental unit; charge conservation holds at microscopic level; Stokes' law applies to microscopic objects in air. It also provided Avogadro's number through N_A = F/e (Faraday constant divided by elementary charge), linking macroscopic electrolysis measurements with microscopic charge. Modern versions use similar principles with improved instrumentation: laser illumination, video tracking, digital timing, and computerized analysis. The experiment remains a classic physics laboratory exercise, teaching statistical analysis, error estimation, and experimental design. Millikan received the 1923 Nobel Prize in Physics for this work and for his research on the photoelectric effect. Beyond historical significance, charge quantization underpins all of electronics and solid-state physics. The experiment exemplifies how simple apparatus with clever design can answer fundamental questions. Millikan's meticulous methodology—controlling variables, making repeated measurements, applying statistical analysis—set standards for experimental physics that continue to influence scientific practice today.

The formation of sodium chloride (NaCl) from its constituent elements represents one of chemistry's most fundamental and illustrative examples of ionic bonding. This experiment demonstrates how the transfer of electrons from sodium atoms to chlorine atoms creates ions that attract through electrostatic forces, forming the characteristic crystal lattice structure of table salt. Understanding this process provides key insights into ionic bonding, lattice energy, and the properties of ionic compounds that form the basis of much inorganic chemistry. The reaction involves elemental sodium (a soft, silvery-white metal) reacting violently with chlorine gas (a greenish-yellow, toxic gas) to produce sodium chloride (a white crystalline solid): 2Na(s) + Cl₂(g) → 2NaCl(s). This highly exothermic reaction releases approximately 411 kJ/mol of energy, manifested as bright yellow flame and heat. The dramatic nature of this reaction makes it an excellent demonstration of chemical bonding principles, though it must be conducted with extreme caution due to the reactivity of both reactants. At the atomic level, the process involves: sodium atoms (electron configuration [Ne]3s¹) lose their single valence electron to achieve stable noble gas configuration ([Ne]), forming Na⁺ ions; chlorine atoms (electron configuration [Ne]3s²3p⁵) gain one electron to achieve stable noble gas configuration ([Ar]), forming Cl⁻ ions. This electron transfer creates ions with opposite charges that attract through Coulombic forces, assembling into a face-centered cubic crystal lattice where each Na⁺ is surrounded by six Cl⁻ ions and vice versa. The experiment can be demonstrated in several ways: dropping small pieces of sodium into chlorine gas (most dramatic but hazardous), reacting sodium with hydrochloric acid (safer classroom demonstration), or electrolysis of molten sodium chloride (showing reverse process). In classroom settings, the reaction is often simulated or shown via video due to safety concerns with handling reactive sodium and toxic chlorine gas. Key observations include: sodium metal melts due to heat of reaction (melting point 97.8°C), bright yellow flame characteristic of sodium emission, white smoke/solid of sodium chloride forming, and eventual cooling to white crystalline solid. The product can be tested for ionic properties: high melting point (801°C), solubility in water, electrical conductivity in molten or aqueous state, and characteristic cubic crystal structure visible under microscope. Theoretical calculations illustrate the energetics: ionization energy of Na (496 kJ/mol, endothermic), electron affinity of Cl (-349 kJ/mol, exothermic), lattice energy of NaCl (-788 kJ/mol, highly exothermic). The net exothermicity (-411 kJ/mol) explains the reaction's spontaneity. Born-Haber cycle analysis reconciles these energy terms, providing quantitative understanding of ionic compound stability. Modern computational chemistry can simulate this reaction at quantum mechanical level, showing electron density changes, charge transfer, and lattice formation. Molecular dynamics simulations can model crystal growth and defect formation. These computational approaches complement experimental observations, providing deeper understanding of bonding dynamics. Beyond academic interest, sodium chloride formation has industrial significance: it's produced on massive scale for food seasoning, preservation, chemical feedstock (chlor-alkali industry), water softening, and de-icing. Understanding its formation informs material science applications like solid electrolytes, optical materials (sodium chloride windows for IR spectroscopy), and geochemical processes (salt dome formation). The experiment teaches multiple concepts: periodic trends (sodium's low ionization energy, chlorine's high electron affinity), octet rule, ionic radii, crystal structures, lattice energy, and Born-Haber cycles. It illustrates how simple electron transfer leads to dramatic property changes: from reactive metal and toxic gas to stable, edible compound. This transformation exemplifies chemistry's power to create useful materials from elemental components through understanding of bonding principles.