3.1 Developing a Nuclear Model of the Atom

A comprehensive vocabulary review deck containing exactly 200 flashcards covering the historical development of atomic models, major scientists, landmark experiments, limitations, electromagnetic radiation, quantum theory, and spectral calculations based on the lecture study notes.

Democritus

An early Greek thinker (~$400\,\text{BC}$) who proposed that matter cannot be cut infinitely and eventually reaches a smallest indivisible piece called 'atomos'.

Atomos

The Greek word meaning 'not to be cut,' used by Democritus to describe the smallest indivisible particle of matter.

Fundamental Premise of Democritus

The philosophical idea that cutting matter smaller and smaller eventually leads to an ultimate piece that cannot be divided any further.

Democritus Model Limitation

His idea lacked experimental evidence and equipment support, and later scientists proved that atoms themselves contain smaller subatomic particles.

Democritus Test Connection

The key association linking Democritus to the word atomos, meaning 'not to be cut.'

Atomic Model Evolution Principle

The scientific process where atomic models are continually upgraded or replaced whenever new experimental evidence contradicts older models.

Observation-to-Conclusion Connection

The required testing principle of linking specific experimental observations directly to the logical atomic conclusions they demand.

John Dalton

The scientist who developed an early scientific atomic theory in $1805$, proposing that matter is made of solid, indivisible atoms.

Billiard Ball Model

John Dalton's model picturing the atom as a tiny, solid, hard, indivisible sphere.

Dalton's Postulate on Identical Atoms

Dalton's claim that all atoms of the same element are completely identical and possess the exact same mass.

Dalton's Postulate on Different Elements

Dalton's claim that atoms of different elements have different properties and different atomic masses.

Dalton Model Indivisibility Limit

Later discovery showed atoms are divisible because they contain subatomic particles (electrons, protons, and neutrons).

Dalton Model Mass Limit (Isotopes)

Later discovery showed atoms of the same element can have different masses, known as isotopes.

Dalton's Symbols (1808)

An early system created in $1808$ using marked circles to represent elements and combined atoms on paper.

Law of Definite Composition

The chemical law stating that elements in a particular compound always combine in a fixed, characteristic mass ratio.

Law of Multiple Proportions

The chemical law stating that when elements form more than one compound, the masses combine in small whole-number ratios.

Law of Conservation of Mass

The law stating that matter is neither created nor destroyed in a chemical reaction, making mass of reactants equal to mass of products.

Mass of Reactants Equation

The relationship expressing the Law of Conservation of Mass: $\text{mass of reactants} = \text{mass of products}$.

J. J. Thomson

The scientist who in $1897$ provided evidence that atoms contain smaller negatively charged particles, proving atoms are divisible.

Cathode Ray Tube (CRT)

The piece of equipment Thomson used to pass an electric current through a gas, producing rays of negative particles.

Cathode Rays

Streams of negatively charged particles produced when an electric current passes through a gas inside a CRT.

Corpuscles

The original name J. J. Thomson gave to the negatively charged subatomic particles now known as electrons.

Electron

A negatively charged subatomic particle discovered inside the atom by J. J. Thomson.

Thomson's Primary Experiment Question

If a gas atom is neutral and supposed to be indivisible, where do the emitted negatively charged particles come from?

Thomson's Atomic Divisibility Conclusion

Because negative particles came from within neutral gas atoms, particles smaller than an atom must exist, making atoms divisible.

Thomson's Positive Charge Deduction

Because original gas atoms are overall electrically neutral, a positive charge must exist in the atom to balance negative electrons.

Raisin Bun Model

Thomson's atomic model picturing negatively charged electrons embedded throughout a sphere of positively charged material.

Raisin Bun Model Positive Sphere

The positively charged 'bun' material in Thomson's model that spreads evenly to balance electron negative charges.

Raisin Bun Model Negative 'Raisins'

The negatively charged electrons stuck inside the positive sphere in Thomson's atomic model.

Limit of Thomson's Raisin Bun Model

It lacked a central concentrated nucleus, which Rutherford later proved exists via alpha particle scattering.

Thomson Test Connection

The logical sequence: $\text{CRT} \rightarrow \text{negative particles} \rightarrow \text{electrons} \rightarrow \text{atom is divisible} \rightarrow \text{Raisin Bun Model}$.

Ernest Rutherford

The scientist who tested Thomson's model in $1911$ using the gold foil experiment and developed the nuclear model of the atom.

Gold Foil Experiment

An experiment where positively charged alpha particles were fired at a thin sheet of gold foil to test positive charge distribution.

Alpha (\alpha) Particles

Positively charged particles emitted by radium, used by Rutherford to probe gold atom structure.

Gold Foil Thickness

The gold foil sheet used in Rutherford's experiment was approximately $2000\,\text{atoms}$ thick.

Rutherford's Primary Observation

Most alpha particles passed straight through the gold foil without any change in direction.

Rutherford's Secondary Observation

A small fraction of alpha particles were deflected at angles as they passed through the foil.

Rutherford's Deflection Anomaly

A very small number of alpha particles bounced straight backward, as if hitting a very dense object.

Rutherford's Tissue Paper Analogy

Rutherford compared backward alpha deflections to firing a $15\text{-inch}$ shell at tissue paper and having it bounce back at you.

Conclusion from Un-deflected Alpha Particles

Because most alpha particles passed straight through, Rutherford concluded that atoms consist mostly of empty space.

Conclusion from Back-Scattered Alpha Particles

Because rare alpha particles bounced back sharply, positive charge and mass must be concentrated in a tiny, dense centre.

Nucleus

The tiny, dense, positively charged centre of an atom discovered by Ernest Rutherford.

Protons in Rutherford's Model

Positively charged subatomic particles identified as residing inside the dense central nucleus.

Nuclear Model (Planetary Model)

Rutherford's atomic model featuring electrons orbiting around a small, dense, positively charged nucleus.

Relative Nucleus Size

The nucleus is extremely tiny compared with the overall volume of the atom.

Rutherford Model Electron Stability Problem

Classical physics predicts orbiting accelerating electrons constantly radiate energy as light and should spiral into the nucleus, causing atomic collapse.

Accelerating Charged Particles

In classical electrodynamics, any charged particle moving in a curved path accelerates and continuously emits electromagnetic radiation.

Rutherford Model Nuclear Mass Problem

The nucleus was far too heavy to contain only protons, indicating an unknown missing neutral mass.

Rutherford Model Proton Repulsion Problem

A nucleus containing only positive charges should fly apart due to intense electrostatic repulsion between like charges.

Rutherford Test Connection

Gold foil $\rightarrow$ mostly straight through = empty space; a few deflect/bounce = small dense positive nucleus.

James Chadwick

The scientist who in $1932$ discovered the neutron, solving Rutherford's missing nuclear mass discrepancy.

Chadwick's Nuclear Mass Discrepancy

The experimental finding that measured nuclear mass was greater than could be explained by known protons alone.

Neutron

A neutral subatomic particle discovered by James Chadwick located inside the atomic nucleus.

Neutron Charge

Neutral (zero electric charge).

Neutron Function in Proton Repulsion

Neutrons add mass and help stabilize the nucleus by reducing electrostatic proton-proton repulsion.

Isotopes

Atoms of the same element that have the same number of protons but different numbers of neutrons and different masses.

Chadwick's Contribution to the Nuclear Model

Improved the nuclear model by adding neutrons to the nucleus alongside protons, explaining extra mass and isotopes.

Unresolved Atomic Problem Post-Chadwick

How electrons can orbit around the nucleus without radiating energy and collapsing.

Chadwick Test Connection

Chadwick = missing nuclear mass $\rightarrow$ neutron.

Stimulated Gas Emission Observation

When gas atoms are excited by electricity or heat, they emit light of specific characteristic colors rather than a full rainbow.

Clue to Electron Energy

The distinct line pattern of light emitted from stimulated atoms provided the key evidence needed to understand electron energy levels.

Max Planck

The scientist ($1858\text{--}1947$) credited with starting the quantum revolution by analyzing blackbody radiation.

Blackbody

An ideal, perfect physical emitter and absorber of electromagnetic radiation that reflects no light and emits energy based on temperature.

Classical Prediction of Blackbody Radiation

Classical physics assumed energy was continuous and predicted that emitted intensity would grow infinitely at higher frequencies.

Blackbody Intensity Curve

An experimental bell-shaped graph plotting radiation intensity versus wavelength for a heated object.

Temperature Effect on Blackbody Glow

As a blackbody becomes hotter, its glowing color shifts from red toward white, and its total emitted energy increases.

Quantum Hypothesis

Max Planck's proposal that energy is not continuous, but is emitted or absorbed in small, discrete quantities called quanta.

Quantum

A small, discrete, indivisible packet of energy that represents the minimum quantity of energy change.

Quantized Oscillating Atoms

Planck's concept that atomic oscillators can only vibrate at specific allowed energy states or whole-number multiples of quanta.

Ramp vs Stairs Analogy for Energy

Continuous energy is like a smooth ramp (any value allowed); quantized energy is like a staircase (only specific steps allowed).

Money Analogy for Quantization

Currency consists of allowed coin denominations; you can combine whole coins, but cannot make arbitrary non-existent coin fractions.

Limit of Planck's Quantum Hypothesis

Planck's theory explained energy packets emitted by blackbodies, but did not propose a full structural model of the atom.

Planck Test Connection

Planck = blackbody radiation $\rightarrow$ energy is quantized.

Electromagnetic Radiation

Energy traveling through space as waves made of oscillating electric and magnetic fields.

Electromagnetic Field Perpendicularity

The electric and magnetic fields in electromagnetic radiation oscillate at right angles ($90^\circ$) to each other and to wave direction.

Wavelength (\lambda)

The shortest distance between equivalent points on a wave, such as peak to peak.

Nanometre to Metre Conversion

1 nanometre equals $1 \times 10^{-9}\,\text{m}$ ($1\,\text{nm} = 1 \times 10^{-9}\,\text{m}$).

Frequency (\nu)

The number of wave cycles that pass a given point per unit time.

Hertz (\text{Hz})

The SI unit of frequency, equivalent to one cycle per second or inverse seconds ($\text{s}^{-1}$).

Speed of Light (c)

The constant speed of electromagnetic radiation in a vacuum, equal to $3.0 \times 10^8\,\text{m/s}$.

Wavelength-Frequency Relationship

Wavelength and frequency are inversely proportional; as wavelength decreases, frequency increases.

Frequency-Energy Relationship

Frequency and photon energy are directly proportional; higher frequency means higher energy.

Key Electromagnetic Trend

$\lambda \downarrow \rightarrow \nu \uparrow \rightarrow E \uparrow$ (shorter wavelength = higher frequency = higher energy).

Infrared vs Ultraviolet Photon Energy

Ultraviolet light has shorter wavelength and higher frequency than infrared light, so a UV photon carries more energy.

Red vs Blue Light Energy Comparison

Blue light has a shorter wavelength and higher frequency than red light, so blue light carries higher energy.

Wave Speed Equation

$c = \lambda \nu$, connecting light speed $c$, wavelength $\lambda$, and frequency $\nu$.

Frequency Calculation Equation

$\nu = \frac{c}{\lambda}$, where $c = 3.0 \times 10^8\,\text{m/s}$.

Ancient Greek View of Light

Early Greek philosophers believed light was a stream of tiny particles.

Christiaan Huygens

A scientist who argued that light is a wave rather than a stream of particles.

Isaac Newton's Light Theory

Newton believed light was composed of tiny particles he called corpuscles.

James Clerk Maxwell

The scientist who formulated classical electromagnetic theory, describing light as an electromagnetic wave.

Classical Theory of Light

Maxwell's wave theory defining light as oscillating electric and magnetic fields traveling through space.

Limit of Maxwell's Light Theory

The classical wave model failed to explain experimental observations of the photoelectric effect.

Maxwell Test Connection

Maxwell = classical wave model of light; Einstein later adds the photon concept.

Heinrich Hertz

The scientist who in $1887$ first observed the photoelectric effect.

Photoelectric Effect

The emission of electrons from a metal surface when the metal absorbs electromagnetic radiation.

Photoelectric Effect Experimental Setup

Radiant light strikes a metal plate in a circuit, ejecting electrons that travel to a collector plate to create a current.

Failure of Classical Wave Theory in Photoelectric Effect

Classical wave theory predicted light intensity (brightness) would control electron ejection energy, but experiments proved frequency controls it.

Intensity vs Frequency Paradox

Dim light above a threshold frequency ejects electrons immediately, while extremely bright light below threshold frequency fails to eject electrons.

Albert Einstein

The scientist who used Planck's quantum concept to explain the photoelectric effect by describing light as photons.