Potassium isotope

Here is an example of anion formation. The element fluorine has 9 protons and 9 electrons.

To calculate the total number of present electrons, you simply add the amount of extra charge to the atomic number. In the case of a negative ion, there are fewer protons than electrons. For example, N3- has a -3 charge; therefore, it has gained 3 electrons compared to the neutral state.

An atom contains equal numbers of protons and electrons . Since protons and electrons have equal and opposite charges , this means that atoms are neutral overall.

Where Are Electrons? Unlike protons and neutrons, which are located inside the nucleus at the center of the atom, electrons are found outside the nucleus.

Electrons belong to the first generation of the lepton particle family, and are generally thought to be elementary particles because they have no known components or substructure. The electron has a mass that is approximately 1/1836 that of the proton.

Neutrons are all identical to each other, just as protons are. Atoms of a particular element must have the same number of protons but can have different numbers of neutrons.

Negative and positive charges of equal magnitude cancel each other out. This means that the negative charge on an electron perfectly balances the positive charge on the proton. In other words, a neutral atom must have exactly one electron for every proton. If a neutral atom has 1 proton, it must have 1 electron.

Electrons are equal to protons so that positive and negative charges are balanced and the atom is electrically neutral…. Since neutrons have no charge it doesn’t matter them being equal to protons…..

Electron

Quarks, the smallest particles in the universe, are far smaller and operate at much higher energy levels than the protons and neutrons in which they are found.

Ernest Rutherford

Sir Joseph John Thomson OM

Ernest Rutherford’s most famous experiment is the gold foil experiment. A beam of alpha particles was aimed at a piece of gold foil. Most alpha particles passed through the foil, but a few were scattered backward. This showed that most of the atom is empty space surrounding a tiny nucleus.

A tiny number of alpha particles, traveling at 10% of the speed of light, hit a dense atomic center right in its middle. The collision and the repulsion cause the alpha particle to “bounce” backwards and move on a very different path. These are the reflected rays.

Conclusion of Rutherford’s scattering experiment: Most of the space inside the atom is empty because most of the α-particles passed through the gold foil without getting deflected. Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space.

The salient features of this model are as follows: (i) The atom contains a central part called nucleus which is surrounded by electrons. (ii) The nucleus of an atom is positively charged. (iii) The size of the nucleus is very small as compared to the atomic size.

Most important, he postulated the nuclear structure of the atom: experiments done in Rutherford’s laboratory showed that when alpha particles are fired into gas atoms, a few are violently deflected, which implies a dense, positively charged central region containing most of the atomic mass.

Salient features of Niels Bohr atomic model are:

• Electrons revolve around the nucleus in stable orbits without emission of radiant energy.
• An orbit or energy level is designated as K, L, M, N shells.
• An electron emits or absorbs energy when it jumps from one orbit or energy level to another.

Bohr’s model of the hydrogen atom is based on three postulates: (1) an electron moves around the nucleus in a circular orbit, (2) an electron’s angular momentum in the orbit is quantized, and (3) the change in an electron’s energy as it makes a quantum jump from one orbit to another is always accompanied by the …

Merits: 1) Bohr’s atomic model explained the stability of an atom. According to Bohr, an electron revolving in a particular orbit cannot lose energy. Therefore, emission of radiation is not possible as long as the electron remains in one of its energy levels and hence there is no cause of insatbility in his model.

In 1913, Niels Bohr proposed a theory for the hydrogen atom, based on quantum theory that some physical quantities only take discrete values. Electrons move around a nucleus, but only in prescribed orbits, and If electrons jump to a lower-energy orbit, the difference is sent out as radiation.

Rutherford randomly placed the negative electrons outside the nucleus. Bohr’s improvement of the Rutherford model was that Bohr placed the electrons in distinct energy levels. Bohr thought that electrons orbited the nucleus in quantised orbits.

The Bohr Model is very limited in terms of size. Poor spectral predictions are obtained when larger atoms are in question. It cannot predict the relative intensities of spectral lines. It does not explain the Zeeman Effect, when the spectral line is split into several components in the presence of a magnetic field.

Schrödinger used mathematical equations to describe the likelihood of finding an electron in a certain position. This atomic model is known as the quantum mechanical model of the atom. Until 1932, the atom was believed to be composed of a positively charged nucleus surrounded by negatively charged electrons.

Assuming that matter (e.g., electrons) could be regarded as both particles and waves, in 1926 Erwin Schrödinger formulated a wave equation that accurately calculated the energy levels of electrons in atoms.