Arrangement of Electrons in Atoms The Development of a New Atomic Model

Light Before 1900, scientists thought that light behaved only as wave discovered that also has particle-like characteristics

Light is electromagnetic radiation. form of energy that acts as a wave as it travels includes: gamma, X rays, ultraviolet (UV), visible and infrared (IR) light, microwaves, and radio waves All forms are combined to form electromagnetic spectrum

Light as a Wave

c = λf Wave Properties: all form of EM radiation travel at a speed (c) of 3.0 x 108 m/s in a vacuum wavelength: (λ) distance between points on adjacent waves; in nm (109nm = 1m) frequency: (f) number of waves that passes a point in a second, in waves/second Inversely proportional! c = λf

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

E = hf Quantum Theory Max Planck (1900) German physicist Observed - emission of light from hot objects Concluded –energy is emitted in small, specific amounts (quanta) Quantum - minimum amount of energy change Max Planck Planck’s constant (h) = 6.626 x 10-34 J*s 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. When studying elements, scientists found that when certain elements were heated in a flame those elements emitted varying colors of light. Light carries energy Something in the atom must be carrying energy in order to emit light E = hf Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Photoelectric Effect Albert Einstein (1905) Observed - photoelectric effect 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

Photoelectric Effect No electrons are emitted Electrons are emitted Bright red light infrared rays Dim blue light ultraviolet rays or or Metal plate Metal plate

Photoelectric Effect when light is shone on a piece of metal, electrons can be emitted no electrons were emitted if the light’s frequency was below a certain value scientists could not explain this with their classical theories of light

Solar Calculator Solar Panel

Photoelectric Effect Einstein added on to Planck’s theory in 1905 suggested that light can be viewed as stream of particles photon- particle of EM radiation having no mass and carrying one quantum of energy energy of photon depends on frequency

Photoelectric Effect EM radiation can only be absorbed by matter in whole numbers of photons when metal is hit by light, an electron must absorb a certain minimum amount of energy to knock the electron loose this minimum energy is created by a minimum frequency since electrons in different metal atoms are bound more or less tightly, then they require more or less energy

H Line-Emission Spectrum Why had hydrogen atoms only given off specific frequencies of light? current Quantum Theory attempts to explain this using a new theory of atom

H Line-Emission Spectrum ground state- lowest energy state of an atom excited state- when an atom has higher potential energy than it has at ground state (excited by heat or electricity) line-emission spectrum- pattern of wavelengths of light created when visible light from excited atoms is shined through a prism

Line-Emission Spectrum pattern of wavelengths of light created when visible light from excited atoms is shined through a prism excited state Wavelength (nm) 410 nm 486 nm 656 nm 434 nm ENERGY IN PHOTON OUT Prism Slits ground state Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

H Line-Emission Spectrum scientists using classical theory expected atoms to be excited by whatever energy they absorbed continuous spectrum- emission of continuous range of frequencies of EM radiation

H Line-Emission Spectrum when an excited atom falls back to ground state, it emits photon of radiation the photon is equal to the difference in energy of the original and final states of atom since only certain frequencies are emitted, the differences between the states must be constant

Bohr Model created by Niels Bohr (Danish physicist) in 1913 linked atom’s electron with emission spectrum electron can circle nucleus in certain paths, in which it has a certain amount of energy

Bohr Model Can gain energy by moving to a higher rung on ladder Can lose energy by moving to lower rung on ladder Cannot gain or lose while on same rung of ladder

Bohr Model a photon is released that has an energy equal to the difference between the initial and final energy orbits

Bohr Model Energy of photon depends on the difference in energy levels 6 Energy of photon depends on the difference in energy levels Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 5 4 3 2 1 nucleus Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

http://www.teachersdomain.org/asset/phy03_vid_quantum/ http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

Other Elements Each element has a unique bright-line emission spectrum. i.e. “Atomic Fingerprint” Helium Bohr’s calculations only worked for hydrogen!  Didn’t explain chemical behavior of atoms. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

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Electrons as Waves In 1924, Louis de Broglie (French scientist) suggested the way quantized electrons orbit the nucleus is similar to behavior of wave electrons can be seen as waves confined to the space around a nucleus waves could only be certain frequencies since electrons can only have certain amounts of energy

Uncertainty Principle In 1927 by Werner Heisenberg (German theoretical physicist) electrons can only be detected by their interaction with photons any attempt to locate a specific electron with a photon knocks the electron off course Heisenberg Uncertainty Principle- it is impossible to know both the position and velocity of an electron

Electrons as Waves shows that anything with both mass and velocity has a corresponding wavelength

Schrödinger Wave Equation In 1926, Erwin Schrödinger (Austrian physicist) his equation proved that electron energies are quantized only waves of specific energies provided solutions to his equation solutions to his equation are called wave functions

Teacher, may I be excused? My brain is full!