Structure of the Atom: A Quantum Perspective
Structure of the Atom
Dalton’s atomic theory postulated that matter is made up of indivisible atoms. This concept held for nearly a century until Becquerel’s discovery of natural radioactivity revealed alpha, beta, and gamma particles. The true nature of subatomic particles emerged with the use of discharge tubes.
The Electron and Thomson’s Atomic Model
Studies of gas conductivity at low pressure in discharge tubes revealed cathode rays, a stream of negatively charged particles. These particles traveled in straight lines, possessed high kinetic energy, and behaved like an electric current. Thomson determined the charge-to-mass ratio (q/m) of these particles to be -1.76 x 1011 C/kg.
The Proton and Rutherford’s Model
Goldstein observed canal rays, positively charged particles seemingly originating from the cathode. Rutherford and Chadwick later identified these particles as protons. Rutherford’s gold foil experiment, using alpha particles, revealed that most particles passed straight through the foil, some were deflected, and a few were repelled. This led to the nuclear model of the atom, with a dense, positively charged nucleus containing most of the mass and electrons orbiting around it.
Isotopes
Mass spectrometry revealed that atoms of the same element can have different masses. These are called isotopes. The mass of an atom is determined by the number of protons and neutrons in its nucleus.
The Neutron
Moseley’s work on X-rays allowed him to determine the atomic number (Z), the number of protons in an atom’s nucleus. The discrepancy between atomic mass and the mass of protons led to the discovery of the neutron.
Atomic Magnitudes
- Atomic Number (Z): Number of protons in the nucleus.
- Mass Number (A): Total number of protons and neutrons in the nucleus.
- Isotopic Mass: Mass of a specific isotope.
- Atomic Mass: Average mass of all isotopes of an element.
Origins of Quantum Theory
Maxwell’s work on electromagnetism laid the groundwork for quantum theory. Key wave characteristics include amplitude, wavelength, frequency, period, and velocity.
Atomic Emission Spectra
Gaseous elements emit light at specific frequencies, producing discontinuous emission spectra.
Planck’s Quantum Theory
Planck proposed that energy is emitted or absorbed in discrete packets called quanta. E = hv, where h is Planck’s constant (6.625 x 10-34 J·s).
The Photoelectric Effect
The photoelectric effect demonstrates that electrons are emitted from a metal surface when light of a certain frequency shines on it. The kinetic energy of the emitted electrons depends on the frequency of the light, not its intensity.
Limitations of Rutherford’s Model
Rutherford’s model could not explain the stability of atoms. Accelerating electrons should emit radiation and spiral into the nucleus.
Bohr’s Atomic Model
: 1913. -The energy of the electrons inside the atom is quantized.That is, the electron occupies only one position or steady states around the core values for which some energy. -The following electron moves around the circular orb nucleo.Cada of these orbits corresponds to a steady state or energy level permitted and is associated with an unnatural number £ o: 1.2 .. -The energy levels allowed when electrons are those in which the angular momentum is an integer multiple of h/2? where Planck’s constant h MVR = nx (h/2?)-Only if it absorbs or emits energy when an electron passes from one level energy to another, is called the energy of the starting level and is the energy level of arrival, the variation of energy and its corresponding frequency will be AE = EF-.. i | AE | = HV v = AE / h = | (EF-ei) / h |