Atomic Spectra
Atomic spectra
Since an electron can exist only at certain energy levels, it can emit photons of certain frequencies. These specific frequencies of light are then observed as spectral lines. When an electron falls from one energy level in an atom to a lower energy level, it emits a photon of a particular wavelength and energy. When many electrons emit the same wavelength of photons it will result in a spike in the spectrum at this particular wavelength, resulting in the banding pattern seen in atomic emission spectra. Likewise, if a photon has the exact wavelength (and energy) needed by the electron to be able to jump up the energy levels, it gets absorbed. Then the number of photons of this wavelength will drop and it will show up as a dark band (bunch of atomic absorption lines) in this part of the spectrum.
When atoms are excited they emit light of certain wavelengths which correspond to
different colors. The emitted light can be observed as a series of colored lines with dark
spaces in between; this series of colored lines is called a line or atomic spectra. Each
element produces a unique set of spectral lines. Since no two elements emit the same
spectral lines, elements can be identified by their line spectrum. The examples are neon
signs, sodium and mercury vapour lamps.
Electromagnetic Radiation and the Wave Particle Duality
Electromagnetic radiation (energy) can travel through vacuum or through materials as a transverse wave. The spectrum of electromagnetic radiation has all possible wavelengths
Electromagnetic Radiation and the Wave Particle Duality
Electromagnetic radiation (energy) can travel through vacuum or through materials as a transverse wave. The spectrum of electromagnetic radiation has all possible wavelengths
and frequencies including visible light. According to the wave particle duality concept, the
electromagnetic radiation is a wave, and it also behaves like a particle. While trying to
explain the black body radiation, Max Planck discovered that energy was limited to certain
values and was not continuous. Thus when the energy increases, it makes tiny jumps
called quanta. Later Einstein proposed that energy was bundled into packets, which were
subsequently named as photons. So, when the energy increases, it jumps up in steps of
quanta and it does not increase smoothly as assumed in classical physics.
Electrons can only exist in certain areas around the nucleus called shells (orbit). Each orbit corresponds to a specific energy level which is designated by a quantum number n. Since electrons cannot exist between these energy levels, the quantum number n is always an integer value (n=1,2,3,4...). The electron with the lowest energy level (n=1) is the closest to the nucleus. An electron occupying its lowest energy level is said to be in the ground state. The energy of an electron in a certain energy level can be found by the equation: En =RH /n2 ,whereRH isaconstantequalto2.179x10-18 Jandnisnumber of the orbit in which the electron exists.
When light is incident on an atom, its electrons absorb photons, gain energy and jump up to higher energy levels. Similarly, an electron can jump down the energy levels by emitting a photon. The energy of the photon emitted or gained by an electron is:
Electrons can only exist in certain areas around the nucleus called shells (orbit). Each orbit corresponds to a specific energy level which is designated by a quantum number n. Since electrons cannot exist between these energy levels, the quantum number n is always an integer value (n=1,2,3,4...). The electron with the lowest energy level (n=1) is the closest to the nucleus. An electron occupying its lowest energy level is said to be in the ground state. The energy of an electron in a certain energy level can be found by the equation: En =RH /n2 ,whereRH isaconstantequalto2.179x10-18 Jandnisnumber of the orbit in which the electron exists.
When light is incident on an atom, its electrons absorb photons, gain energy and jump up to higher energy levels. Similarly, an electron can jump down the energy levels by emitting a photon. The energy of the photon emitted or gained by an electron is:
Ephoton = RH [ 1 / ni2 – 1 / nf2 ]
where ni is the initial energy level of the electron and nf is the final energy level of the
electron. The frequency of the photon emitted when an electron descends energy levels
can be found using the formula: νphoton = (Ei – Ef), where Ei and Ef are the initial and final
energy of the electron.
Since an electron can exist only at certain energy levels, it can emit photons of certain frequencies. These specific frequencies of light are then observed as spectral lines. When an electron falls from one energy level in an atom to a lower energy level, it emits a photon of a particular wavelength and energy. When many electrons emit the same wavelength of photons it will result in a spike in the spectrum at this particular wavelength, resulting in the banding pattern seen in atomic emission spectra. Likewise, if a photon has the exact wavelength (and energy) needed by the electron to be able to jump up the energy levels, it gets absorbed. Then the number of photons of this wavelength will drop and it will show up as a dark band (bunch of atomic absorption lines) in this part of the spectrum.
If hydrogen atoms inside a lamp are excited by an electric current and the light coming
from the lamp is diffracted into its different frequencies. The frequencies of light
correspond to certain energy levels (n). It is therefore possible to predict the frequencies
of the spectral lines of Hydrogen using the equation ν = 3.2881 x 1015
s−1(1/22−1/n2),where n must be a number greater than 2. These lines are in the visible
light and some longer wavelengths of ultraviolet (Balmer’s lines). There are several other
series in the Hydrogen atom which correspond to different parts of the electromagnetic
spectrum. The Lyman series (n=1) for example, extends into the ultraviolet.
Comments
Post a Comment