From Rutherford to Schrodinger Discovery of the Atom From Rutherford to Schrodinger

Evolution of the atomic model

Rutherford Model Derived from the gold foil experiment

Gold Foil Results

Milikan’s Oil Drop Experiment Mass of oil drop determined by the rate of fall in absence of electric field Charge adjusted to exactly hold oil drop in place. Charge always multiple of -1.60e-19 Coulombs Use e/m ratio, mass of electron 9.110 x 10-28 g.

Waves Low frequency High frequency long wavelength l Amplitude Low frequency short wavelength l Amplitude High frequency 6

Wave characteristics of light Frequency and wavelength are inversely related c = λν

Electromagnetic Spectrum

Electromagnetic Spectrum

Visible Spectrum of Light Waves 1/33,000” long Waves 1/70,000” long Red Orange Yellow Green Blue Indigo Violet PRISM Slit Ray of White Light All light is bent (refracted) passing through a prism; violet is bent most and red least. A beam of sunlight produces a continuous band of rainbow colors showing that light is a mixture of colors. 10

Double Slit Experiment

Emission Spectra for Barium 1870- Kirchoff & Einstein Each element emits a different line spectrum Can be used to identify an element Elements emit only certain wavelengths of light which correspond to particular color “Fingerprint” of element Emission Spectra for Barium

End of 1900 century: Believed that matter and energy were distinct Matter = particles Energy = light in the form of a wave Max Planck (1900) Found that matter could not absorb or emit any quantity of energy Energy is emitted in small, specific amounts (quanta) Quantum = fixed amount Max Planck Why are different ranges of wavelengths emitted at different temperatures? Why do different elements exhibit different colors when heated, or characteristic colors for each element? In 1900, Max Planck explained the “ultraviolet catastrophe” by assuming that the energy of electromagnetic waves is quantized rather than continuous—energy could be gained or lost only in integral multiples of some smallest unit of energy, a quantum. • Classical physics had assumed that energy increased or decreased in a smooth, continuous manner. • Planck postulated that the energy of a particular quantum of radiant energy could be described by the equation E = h, where h is the Planck’s constant and is equal to 6.626 x 10-34 joule•second (J•s). • As the frequency of electromagnetic radiation increases, the magnitude of the associated quantum of radiant energy increases. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 13

Continuous vs. Quantized A B Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 330 14

Quantum Theory Albert Einstein (1905) Used Planck’s theory to explain the photoelectric effect Electrons are ejected from the surface of a metal when light shines on the metal Albert Einstein Planck’s quantization hypothesis was used to explain a second phenomenon that conflicted with classical physics. • When certain metals are exposed to light, electrons are ejected from their surface. – Classical physics predicted that the number of electrons emitted and their kinetic energy should depend only on the intensity of light, not on its frequency. – However, each metal was found to have a characteristic threshold frequency of light — below that frequency, no electrons are emitted regardless of the light’s intensity, above the threshold frequency, the number of electrons emitted was found to be proportional to the intensity of light and their kinetic energy proportional to its frequency, a phenomenon called the photoelectric effect. “The free, unhampered exchange of ideas and scientific conclusions is necessary for the sound development of science, as it is in all spheres of cultural life. ... We must not conceal from ourselves that no improvement in the present depressing situation is possible without a severe struggle; for the handful of those who are really determined to do something is minute in comparison with the mass of the lukewarm and the misguided. ... Humanity is going to need a substantially new way of thinking if it is to survive!" (Albert Einstein) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 15

Photoelectric Effect

Photoelectric Effect No electrons are emitted Electrons are emitted Bright red light infrared rays Dim blue light ultraviolet rays or or When a photon strikes the surface of a metal, it transfers its energy to an electron in a metal atom. The electron either “Swallows” the entire photon or does not swallow any of it. If the energy is too small, it will not be able to escape. The energy (frequency) matters, not the number of photons (intensity). X-rays have high frequencies and high energy (can damage living organisms). Radio waves have low frequencies and low energy and pose no hazard. Metal plate Metal plate

Photoelectric effect Light with a frequency below the threshold frequency will not cause the photoelectric effect Light with a frequency higher than the threshold frequency is linearly related with the energy of the emitted electron All metals had the same results, but had different threshold frequencies E= hv

Quantum Theory The energy of a photon is proportional to its frequency. E = h We are unaware of quantum effects in the world around us because quanta of energy are extremely small (hence the small Planck’s constant). E: energy (J, joules) h: Planck’s constant (6.6262  10-34 J·s) : frequency (Hz) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Implications of Photoelectric Effect Light exists as a quantized, or discrete, packet of energy called a photon Amount of energy dependent on the frequency of light. Therefore light has a dual nature. It behaves both as a wave and as a particle

Dual Nature of Light Light has properties of both waves and particles Wave- Light has a wavelength and frequency Einstein‘s photons of light were individual packets of energy that had many characteristics of particles. • Einstein’s hypothesis that energy is concentrated in localized bundles was in sharp contrast to the classical notion that energy is spread out uniformly in a wave. Particle- Light carries a packet of energy called a photon Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 325 21

Waves Low frequency High frequency long wavelength l Amplitude Low frequency 60 photons 162 photons low energy short wavelength l Amplitude Einstein‘s photons of light were individual packets of energy that had many characteristics of particles. • Einstein’s hypothesis that energy is concentrated in localized bundles was in sharp contrast to the classical notion that energy is spread out uniformly in a wave. High frequency high energy 22

Electromagnetic Spectrum

The New Model of the Atom Must explain the photoelectric effect and line spectra 1915 Niels Bohr – decided to start with hydrogen 1 electron in the atom Line Spectrum for Hydrogen

Bohr model of the atom

Bohr’s Model Nucleus Electron Orbit Energy Levels

Bohr Model of Hydrogen 5 4 3 2 1 nucleus Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Bohr model of the atom How does the Bohr model explain the photoelectric effect and spectral lines? Bohr Model

An Excited Lithium Electron Excited Li electron Energy Photon of red light emitted Lithium electron lower energy state Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 326

Excitation of Hydrogen Atoms Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 328

Return to Ground State

How does the Bohr model fit in with the Periodic Table?

Problems with the Bohr Model The energies of the electrons calculated by Bohr only worked for one electron systems

Problems with the Bohr Model DeBroglie λ = h mv All objects with mass and in motion travel as a wave Therefore, electrons cannot move around the nucleus in a fixed circular path

Problems with the Bohr Model Heisenberg Uncertainty Principle Impossible to know both the velocity and position of an electron at the same time Werner Heisenberg ~1926 g Microscope Werner Heisenberg (1901 - 1976) The uncertainty principle: a free electron moves into the focus of a hypothetical microscope and is struck by a photon of light; the photon transfers momentum to the electron. The reflected photon is seen in the microscope, but the electron has moved out of focus. The electron is not where it appears to be. Electron 36

Problems with the Bohr Model Heisenberg Uncertainty Principle Therefore, the electron cannot be moving around the nucleus in a fixed circular path at a fixed distance from the nucleus. A new model of the atom was proposed

1926- Schrodinger Energy in electrons is quantized Electrons move in a wave-like behavior Uncertainty Principle Believed that the atom contains energy levels Needed to explain line spectra of elements

New Model: Quantum Mechanical Model Based on Probability

The Schrodinger model of the atom Based on probability Electron Clouds

A closer look at s and p orbitals

D orbitals

Results of Schroedinger Calculations

The Schrodinger Model Energy Diagrams

Many Electron System

Maximum number of electrons Orbital Type Number of orbitals Maximum number of electrons s 1 2 p 3 6 d 5 10 f 7 14

Many Electron System

Quantum Numbers n l (n-1 and 0) Orbital Designation ml (-l to l) # of Orbitals 1 2 3 4 1s 2s 2p 3s 3p 3d -1. 0, 1 -1, 0, 1 -2, -1, 0, 1, 2 1 ** 3 ** 5

Atomic Radius Atomic Radius Trends--Table

Ionization Energy Definition of ionization energy Ionization energy: the energy required to remove an electron from an isolated atom or molecule in the gaseous phase

Graph of Ionization Energies

Electron Affinity Definition of Electron Affinity The electron affinity is a measure of the energy change when an electron is added to a neutral atom to form a negative ion

Graph of Electron Affinities

Trends in Electronegativity Definition of electronegativity: a property of an atom which increases with its tendency to attract the electrons of a bond.

Trends in Electronegativity