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Learn about the behavior of light as it travels through space and how electrons occupy the space surrounding the nucleus in atoms. Explore concepts such as wavelength, frequency, quantum theory, and the atomic emission spectrum.
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Chapter 5 Electrons in Atoms
Section 5.1—Light & Quantized Energy • In chapter 4, we discussed Rutherford’s wonderful scientific discovery, the nuclear model of the atom. It lacked one thing—details of how electrons occupy the space surrounding the nucleus. • Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space.
Section 5.1—Light & Quantized Energy • Wavelength () is measured from one crest to the next crest, or from trough to trough. Wavelength is usually expressed in meters, centimeters, or nanometers (1nm = 1x10-9 m). • Frequency () is the number of waves that pass a point per second. One hertz (Hz) is the SI unit of frequency and is expressed as s-1.
Section 5.1—Light & Quantized Energy • The amplitude is the wave’s height from the origin to a crest, or from origin to a trough.
The speed of light (c) in a vacuum is 3.00 x 108 m/s. It has the formula: c = • Wavelength & frequency are inversely related, which means that as one increases, the other decreases.
Section 5.1—Light & Quantized Energy • When white light is passed through a prism, it is separated into a continuous spectrum of colors. • These visible colors only make-up a small portion of the electromagnetic spectrum (EM Spectrum)—includes all the forms of electromagnetic radiation, noting their differences in wavelength & frequency.
Electromagnetic spectrum • ***TRANSPARENCY
Section 5.1—Light & Quantized Energy • The sequence of the visible colors can be remembered by the character Roy G. Biv.
In 1900, a German physicist named Max Planck concluded that matter can gain or lose energy only in small, specific amounts called quanta—Quantum concept.
Section 5.1—Light & Quantized Energy • That is, a quantum is the minimum amount of energy that can be gained or lost by an atom. This is demonstrated by the equation: Equantum = h • Where E is energy, h is Planck’s constant (6.626 x 10-34 J.s), & is frequency.
Planck’s Constant • The frequency of an e- is proportional to its energy.
Section 5.1—Light & Quantized Energy • In the photoelectric effect, electrons are emitted from a metal’s surface when light of a certain frequency shines on the surface. • What does this sound like? • A calculator powered by photoelectric cells converts energy from incident light into electrical energy.
Section 5.1—Light & Quantized Energy • In 1905, Einstein said that EM radiation has both wavelike & particle-like properties. That is, light can also be described as a stream of particles called photons. • Photons are radiation with no mass that carries a quantum of energy. Einstein calculated that a photon’s energy depends on its frequency: Ephoton = h
Section 5.1—Light & Quantized Energy • The atomic emission spectrum of an element is the set of frequencies of the EM waves emitted by atoms of the element. An atom’s emission spectrum can be used to identify that particular element.
Section 5.2—Quantum Theory & the Atom • When an atom is in the lowest energy state, it is in the ground state. • When it gains energy, it is in an excited state. • The smaller the electron’s orbit, the lower the atom’s energy level.
Section 5.2—Quantum Theory & the Atom • Niels Bohr, a Danish physicist who worked alongside Rutherford, in 1913, assigned a quantum number (n) to each orbit & calculated its radius. • Bohr’s model explained the spectral lines of hydrogen. • Bohr’s idea of quantized energy levels laid the groundwork for atomic models to come, but was later disproved because it did not fully account for the chemical behavior of atoms.
Section 5.2—Quantum Theory & the Atom • The movements of electrons in atoms are not completely understood even now; however, substantial evidence indicates that electrons do not move around the nucleus in circular orbits. • In 1924, a French physicist named Louis de Broglie derived an equation that predicts that all moving particles have wave characteristics. = h/m
Section 5.2—Quantum Theory & the Atom • In 1926, Austrian physicist Erwin Schrodinger stated that electrons act as waves which led to the quantum mechanical model of the atom. Very similar to Bohr’s model, but is different in that it makes no attempt to describe the electron’s path around the nucleus. • The quantum mechanical model of the atom predicts a 3-dimensional region around the nucleus called an atomic orbital describes the electron’s probable location.
Section 5.2—Quantum Theory & the Atom • Because the atomic orbital is “fuzzy,” it does not have an exactly defined size. The quantum mechanical model assigns principal quantum numbers (n) that indicate the relative sizes & energies of atomic orbitals.
Therefore, n specifies the atom’s major energy levels called principal energy levels. Principal energy levels contain energy sublevels named s, p, d, or f.
Section 5.2—Quantum Theory & the Atom • All s sublevels are spherical in shape and all p sublevels are dumbbell shaped. Not all d & f orbitals have the same shape. • Each orbital may have a maximum of 2 electrons. So the maximum number of electrons related to each principal energy level equals 2n2.
