Quantum Theory and Atomic structures Flashcards
The Wave and Particle Nature of Light
Limitations of the Planetary Model (Rutherford)
Two significant shortcomings:
1. Unlike planets orbiting a sun, electrons are charged particles. An *electric charge is known to *. An orbiting charge should steadily * energy and * toward the nucleus, colliding with it in a small fraction of a second.
2. The planetary model could not explain the * and
* that were observed.
accelerating, emit electromagnetic waves
lose, spiral
highly peaked emission, absorption spectra of atoms
- is a form of energy propagated as mutually * . It includes visible light, infrared,
ultraviolet, X ray, and radio waves.
Electromagnetic radiation
perpendicular electric and magnetic fields
A * is a disturbance that transmits energy through a medium.
wave
- is the height of the crest of a wave above the center line of the wave.
Amplitude
The * is the distance between successive crests or troughs of a wave motion. Designated by lambda.
wavelength
The frequency of a wave motion is the * or troughs that
* through a given * in a unit of * . It is expressed by the unit time−1 (e.g., s−1, also called a hertz, Hz).
number of wave crests, pass, point, time
distinctive feature of electromagnetic radiation is its constant * of in a
vacuum, often referred to as the speed of light.
speed c = frequency × wavelength
In phase - *
Out of phase - *
Constructive Interference
Destructive Interference
Blackbody Radiation and Classical Physics
A * is one that absorbs and emits all frequencies.
• As temperature T increases, the * of emitted light also increases.
• Spectral distribution of the intensity of blackbody radiation as a function of * . Increase in temperature shifts the maximum to * frequencies.
• Classical physics was unable to explain the distribution of frequencies.
blackbody
frequency
frequency
higher
The first person to offer a successful explanation of blackbody radiation was *.
• He assumed that the * emitted by the body was due to the * of the electrons in the constituent particles in the material body. Similar to classical
physics.
* He then assumed that the * of the oscillators were proportional to an * of the *
Max Planck (1900) radiation, oscillations
energies
integral multiple n, frequency
Planck’s law:
E = nhv
h - 6.626 * 10^-34 J-s
In 1888,* discovered that * light caused electrons to be * from a * surface. This is called the *.
Heinrich Hertz
ultraviolet , emitted, metallic
photoelectric effect
Two experimental observations of the photoelectric effect are in stark contrast with classical wave theory of light.
1. The * of the electrons is independent of the * of the incident radiation.
• Classical physics predicts that the kinetic energy of the emitted electrons should be proportional to the intensity of the incident radiation.
2. There is a * characteristic of the metallic surface, below which no electrons are rejected.
• Classical physics predicts that the photoelectric effect should occur for any frequency of light with sufficient intensity.
kinetic energy, intensity
threshold frequency
Photoelectric Effect and Albert Einstein
In order to explain these results, Albert Einstein used Planck’s hypothesis, and
extended it in an important way.
• Planck applied his quantization concept to the * of electronic oscillators. Once the light was emitted, it behaved like a classical wave.
• Einstein proposed that the radiation itself existed in small *, e = hv, also known as *.
emission and absorption mechanism
packets of energy
photons
Einstein showed that the kinetic energy of the ejected electron is related to the incident radiation and the work function of the metal (f, similar to the
ionization energy)
KE = E(hv) of incident radiation - work function
Balmer Series for Hydrogen Atoms
Emission spectrum of Hydrogen
Johannes Rydberg accounted for all the lines in the hydrogen atomic spectrum by
generalizing the Balmer formula to:
𝟏/𝝀 = R (𝟏/(𝒏𝟏)^𝟐 − 𝟏/(𝒏𝟐)^2)
In 1913, Niels Bohr explained the * by assuming Planck’s idea of energy being quantized.
line spectra
Niels Bohr assumed that the * of the electron in a hydrogen atom is quantized.
• Classical physics predicts that an electron revolving around a nucleus will lose energy and spiral into the nucleus, and so a stable orbit of the electron is
classically forbidden.
• Bohr first assumed the existence of stationary electron orbitals.
• He then specified these orbits by invoking a quantization condition, and assumed that the angular momentum of the electron must be *
angular momentum
quantized