12.7 - SUMMARY

Atoms are made up of two types of particles: protons (which are positively charged) and electrons (which are negatively charged). For an atom to be neutral, the number of protons must equal the number of electrons. This balance ensures that the overall charge of the atom is zero. If the number of protons and electrons is not equal, the atom becomes an ion, which can be either positively charged (if it loses electrons) or negatively charged (if it gains electrons).

J.J. Thomson proposed the 'plum pudding' model of the atom in which he described the atom as a sphere of positive charge with negatively charged electrons scattered within it, much like raisins in a pudding. This model suggested that the positive charge was spread out across the entire atom. However, later experiments led by Ernest Rutherford showed that this model was flawed by demonstrating that the positive charge is concentrated in a small nucleus, with electrons orbiting around it.

Ernest Rutherford's experiments indicated that atoms have a small, dense nucleus at their center that holds most of the atom's mass and positive charge. The electrons orbit this nucleus much like planets orbit the sun. This was a major advancement in understanding atomic structure, suggesting that atoms are mostly empty space, with a tiny nucleus at the center where protons and neutrons reside.

Rutherford's model faced significant challenges, particularly in explaining atomic stability. According to classical physics, electrons in orbit would be accelerating and thus would emit energy in the form of electromagnetic radiation, causing them to lose energy and spiral into the nucleus, which would lead to atomic collapse. Moreover, this model could not account for the distinct line spectra that different elements emit when energized, which suggested that electrons occupy specific energy levels.

Niels Bohr introduced a new model of the atom that incorporated quantum theory to resolve the issues presented by Rutherford's model. His model was specifically designed for hydrogen and similar atoms with a single electron. He proposed that electrons can only exist in specific stable orbits where they do not emit energy, thus preventing them from spiraling into the nucleus. This concept was groundbreaking, as it combined ideas of quantized energy levels with the structure of the atom.

Bohr's first postulate states that an electron revolves around the nucleus only in prescribed, stable orbits without losing energy. His second postulate introduces the concept of quantized angular momentum, indicating that the electron's angular momentum can only take specific values, which leads to discrete energy levels. The third postulate explains that energy can be emitted or absorbed when an electron transitions between these orbits, leading to the observed spectral lines of elements.

Bohr calculated the energy levels of the hydrogen atom, showing they are quantized. For every orbit, there is a quantized energy value, with the energy decreasing as the electron gets closer to the nucleus. The most stable state (ground state) for hydrogen has an energy of -13.6 eV. Higher energy levels correspond to excited states, where the electron is further from the nucleus and requires energy input to reach.

De Broglie's hypothesis introduced the idea that particles, such as electrons, also have wave properties. This suggested that electrons in Bohr's stable orbits can be viewed not just as particles moving in a path but also as waves. When the circumference of the electron's orbit is an integer multiple of its wavelength, standing waves are formed, leading to the quantized orbits.

While Bohr's model effectively describes hydrogenic atoms, it doesn't account for the complexities of multielectron systems, like helium or larger. In these systems, each electron interacts not just with the nucleus but also with every other electron, complicating energy states and making Bohr's simple model insufficient. This led to the development of quantum mechanics to handle these interactions properly.