Atomic Spectra, Chemistry tutorial

Introduction

When a component is heated in a flame or a release of electricity is passed throughout a gas, colours of light trait of the component are seen. For instance, discharge tubes enclosing neon gas at low pressure glow red, sodium light emits yellow light and mercury provides a green glow.

Definition of Atomic Spectra

When the light produced from the element through heating or from a gas as a consequence of the passage of electricity is passed through a spectroscope it generates a spectrum identified as emission spectrum. The emission spectrum that is produced consists of a series of lines, termed line spectrum or atomic spectrum. Each element provides a unique model of spectra lines, and each line has a distinct wavelength. The atomic spectra of the elements are heavier than hydrogen and are significantly more compound.

Line Spectra of Atoms

The radiation created from elemental emission consists of a compound mixture of wavelengths that can be dispersed into its diverse individual components through means of a spectrometer. The consequence of these a measurement yields an emission spectrum which is a plot of the intensity of the emitted radiation as a function of wavelength or another related variable. The spectra of atoms in the gas phase consist of a number of sharp lines. In the case of atomic hydrogen, the spectrum is extremely easy. The occurrence of atomic in addition to molecular line spectra provides resounding evidence that the energies of atoms and molecules must be quantized rather than continuous. When the element or gas is heated to generate the spectrum, the atoms of the component or gas become more energetic. They are as a result promoted to the excited state. After a while, the added energy can be re-emitted in the form of photons as the atom return to the ground state, thereby giving increase to the examined emission spectrum. In a usual gas example made up of huge number of molecules, each atom would obtain an arbitrary amount of energy. This will form a continuous distribution of energies the consequence of that would be continuous radiation spectrum. Though, the examined spectra are made up of sharp lines, which imply that the energies of the atoms and molecules dependable for these lines must be discrete rather than continuous.

1126_The Representation of an Atom with Energy Levels that have.jpg

Fig: The Representation of an Atom with Energy Levels that have an Energy Difference (?ε)

1600_The Absorption and Emission of Energy by Electron.jpg

Fig: The Absorption and Emission of Energy by Electron

When an electron in the excited state makes a transition to ground state the energy is emitted in the form of radiation of energy. Atoms are said to possess energy levels because they can only possess discrete amount of energy. The lowest energy level is known as the ground state while the energy levels above the ground state are referred to as the excited states.

Absorption of energy makes atomic electron jump from the ground state to the excited state.  On transition back to the ground state, the energy earlier absorbed is released in form of emitting a photon. By the law of conservation of energy, the energy loss by the atom in transition from excited back to the ground state must be equal to the energy of the photon.  This gives us the fundamental   formula of spectroscopy.

Δε = εγ  = hν = hc/λ

This equation shows the relation between wavelength of the observed emitted light and the energy difference between the ground and excited states of the atom.

The Spectrum of Atomic Hydrogen

Niel's Bohr developed the concept of quantization of energy levels in electrons in atoms in his explanation of the spectrum of atomic hydrogen. Atomic hydrogen has only one electron hence the simple nature of the emission spectrum for it. It displayed several lines in the visible, infrared and the ultraviolet regions. The observed spectra lines are well fitted by the empirical Rydberg formula

v = 1/λ = [((1/n22)-(1/n12)]

where n1  and n2  are integers(n1  n2), and is the Rydberg's constant , with a value of = 1.097x105  cm-1.The spectrum of atomic hydrogen have all lines belonging to a specified series through a common value of n2 but different values of  n1. On the basis of this there are five series each having the value of n2 as shown in Table.

Table: The Series in the Spectrum of Atomic Hydrogen

1124_The Series in the Spectrum of Atomic Hydrogen.jpg

Example

Compute the wavelength of the observable line in the hydrogen spectrum that has the lengthiest wavelength. What is the energy dissimilarity between the 2 levels of the atom engaged in the transition giving increase to this line?

Visible line means that n2 = 2 of the Balmer series that lies in the visible region of the electromagnetic spectrum. Transition can only be from n2 = 2 to n1 = 3, 4, 5,.... The longest wavelength corresponds to the lowest energy difference because wavelength is inversely proportional to energy of the photon released. The lowest energy difference means that the transition should be between level n2 =2 and level n1 = 3.

1324_Energy Level Diagram of the Hydrogen Atom.jpg

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