arrangement of but, there was one major question electrons...
TRANSCRIPT
Arrangement of Electrons in Atoms
CHAPTER 4
The Development of a New
Atomic Model
• The Rutherford model was a great improvement over the Thomson model of the atom.
• But, there was one major question that needed to be answered.
• If the electrons are negatively charged and the nucleus is positively charged, what prevents the electrons from being drawn into the nucleus of an atom?
Properties of Light
• Prior to 1900, most scientists believed that light behaved as waves.
• Later, we found that light actually behaves as light and as particles.
• Most of light’s behavior is due to its wave-like behavior.
Electromagnetic Radiation
• Electromagnetic radiation is any form of energy that exhibits wavelike behavior as it travels through space.
• Examples include X-rays, UV light, microwaves, visible light, and radio waves.
• All forms of EMR travel at the same speed, 3.0 x 108 m/s. This is called the speed of light.
Electromagnetic Spectrum Light as Waves• Light is repetitive in nature like waves.
• Most of the measurable properties of waves are wavelength and frequency.
• Wavelength (!) is the distance between corresponding points on adjacent waves.
• Wavelength is measured mostly in nanometers (nm).
• (1nm = 1 x 10-9 m)
Light as Waves
(continued)
• Frequency (") is defined as the number of waves that pass a given point in a specific time, usually one second.
• Frequency is measured in hertz (Hz).
• 1 hertz = 1 wave/second
Relating Wavelength and
Frequency
• We can write a mathematical expression that relates frequency and wavelength.
• c = !"
• c is the speed of light, ! is the wavelength, and " is the frequency.
• Since the speed of light is the same for all forms of EMR, the product of wavelength and frequency is constant.
• Wavelength and frequency are inversely proportional.
The Photoelectric Effect
• Wave theory could not explain everything involving the interactions between light and matter.
• One particular phenomenon that scientists were perplexed about was the photoelectric effect.
• The photoelectric effect refers to the emission of electrons from a metal when light shines on a metal.
The Photoelectric Effect
• The central question surrounding the photoelectric effect focused on the frequency of the light that hits a metal.
• Particle theorists believed that there was a minimum frequency of light needed to remove electrons.
• Wave theorists believed that any frequency of light would knock loose electrons.
Solution to the
Photoelectric Effect Problem
• 1900: German physicist Max Planck proposes a partial solution to the photoelectric effect.
• Planck believes that hot objects give off EMR in small packets called quanta.
• A quantum is the minimum amount of energy that can be gained or lost by an atom.
Max Planck 1858 - 1947
Planck’s Constant
• Planck finds a relationship that exists between quanta and the frequency of radiation.
• E = h"
• E (energy in joules), " (frequency), and h (Planck’s constant)
• h = 6.626 x 10-34 J・s
Einstein’s Solution
• 1905: Albert Einstein begins to expand on Planck’s idea about quanta.
• Einstein proposes that EMR has a wave-particle duality.
• He thought that since light and other forms of EMR can be thought of as waves, then EMR can also be thought of a a stream of particles.
• These particles were called photons.
Albert Einstein 1879 - 1955
Photons, Photons, Photons
• Photons are particles of EMR that have zero rest mass and a quantum of energy.
• Planck’s equation can be rewritten in terms of a relationship between the energy of a photon and its frequency.
• Ephoton = h"
• Einstein concluded that in order for an electron to be ejected from a metal, the electron must be struck by a photon with a minimum amount of energy.
•
Light as Particles
• The minimum amount of energy needed is tied to the minimum frequency of the light needed.
• Different elements require different minimum frequencies to exhibit the photoelectric effect.
• Einstein wins a Nobel Prize in Physics in 1924 due to his work on the photoelectric effect.
The Hydrogen-Atom Line-
Emission Spectrum
• Electrons can gain or lose energy.
• Electrons can be in the ground state or the excited state.
• The ground state of an electron is the lowest energy state for an atom and is the most stable.
• The excited state of an electron is any energy state higher in potential energy than the ground state.
The Hydrogen-Atom Line-
Emission Spectrum (continued)
• When scientists passed an electric current through a vacuum containing hydrogen gas at low pressure, the excited hydrogen atoms had a pinkish glow.
• When this light was passed through a prism, the light split into specific color bands.
• This separation is called the line emission spectrum of hydrogen.
Line Emission Spectrum
Ground State vs. Excited State
• In order for an atom (or electron) to reach an excited state from the ground state, energy must be added.
• Once an atom (or electron) reaches an excited state and begins to return to the ground state or lower energy state, the atom releases a photon of energy.
• This photon has an energy that is equal to the energy difference between the two energy states.
• Ephoton = E2 -E1 or Einitial - Efinal
Ground State vs. Excited State
Bohr Model of the Atom
• 1913: Danish physicist Niels Bohr proposes a model of the atom based on electrons and photon emission.
• Main idea: electrons revolve around the nucleus in circular paths called orbits.Niels Bohr
1885 - 1962
Bohr Model of the Atom
• The atom and electrons are in the lowest energy state (ground state) when the electrons are in orbits closest to the nucleus.
• The energy of the electron increases as the distance between the orbit and nucleus increases.
• In order to move from orbit to orbit, a photon must be released or absorbed.
The Quantum Model of the Atom
• 1924: French scientist Louis de Broglie suggests that electrons be considered as waves confined to the space around an atomic nucleus
Louis de Broglie 1892 - 1987
The Quantum Model of the Atom
• 1927: German physicist Werner Heisenberg tries to detect electrons by using their interactions with photons.
• Since photons have about the same energy as electrons, finding a specific electron with a photon would cause the electron to veer off course.Werner Heisenberg
1901 - 1976
The Quantum Model of the Atom
• Due to this fact, Heisenberg deduces that there is always an uncertainty in attempting to find an electron.
• Heisenberg Uncertainty Principle (HUP): It is impossible to know both the position and velocity of an electron.
The Quantum Model of the Atom
• 1926: Austrian physicist Erwin Schrödinger uses the wave-particle duality to write an equation that treats electrons as waves.
• Pairing of the Schrödinger Wave Equation (SWE) with the HUP, leads to the foundation of modern quantum theory.
Erwin Schrödinger 1887 - 1961
The Quantum Model of the Atom
• When the SWE is solved, the results are called wave functions.
• Wave functions can only give the probability of finding an electron at a given point.
• Due to wave functions, we know that electrons do not travel in circular orbits. Instead, electrons reside in certain regions called orbitals (3-D regions about the nucleus that indicate the probable location of an electron).