Atomic Structure

The atom is the basic building block of matter. An atom is, in fact, the smallest unit of an element.  It is comprised of subatomic particles: protons, neutrons, and electrons. Chemical changes may not break an atom of an element down any farther.  Ernest Rutherford found in 1911 that an atom has a dense, positively charged nucleus that takes up a very small volume of the entire atom. Protons and neutrons are known as nucleons; they are found in the nucleus. Electrons are present outside the nucleus in regions of space known as orbitals. Atoms of the same element exhibit similar chemical properties and characteristics.

A Summary of the Subatomic Particles

Electrons

Electrons are subatomic particles that carry negative charges (of approximately -1.602·10-19C). An electron is prevalently believed to be an elementary particle of negative charge and has a mass of 9.11·10-31kg, approximately 1/1836 that of the proton. Their masses are thus negligible when calculating atomic mass. The valence electrons of an atom (the ones farthest away from the nucleus on the outer shell) determine the specified atom’s chemical reactivity. The electron was identified as a particle in 1897 by J. J. Thomson (1856-1940).

Protons

The proton is a subatomic particle with an electric charge of +1.602·10-19C. It is found in the nucleus of each atom, along with neutrons, but is also stable by itself. In fact, a proton is essentially the hydrogen ion, H+. The mass of a proton is approximately 1.67·10-27kg.

The Concept of Atomic Mass

Atomic mass units, also known as amus,  are used to compare the relative masses of atoms (the atomic masses). One amu is about the size of a proton or neutron (1.67·10-27kg). One amu is defined as one-twelth of the mass of a carbon-12 atom, which retains 6 neutrons and 6 protons. Thus, carbon-12 has an atomic mass of exactly 12amu.

The atomic mass of an element is also equivalent to the molar mass of the element. The molar mass is the mass of one mole of, or an Avogadro’s Number (6.022·1023kg) of, atoms of that element.

The Bohr Model of the Hydrogen Atom

Niels Bohr (1885-1962) postulated that an electron in a hydrogen atom can only exist in fixed energy states, or the energy of an atom is quantized. The Bohr Model views electrons as circling the nucleus. The greater the radius of the orbit of the electron around the nucleus, the greater the energy state of the electron, and vice versa. At ground state, the electron is at its lowest energy level. The atomic emission spectrum can thus be explained from the Bohr Model; when electrons fall from one energy level to a lower one, energy is released in the form of light. Since electrons exist in fixed energy states, only certain quantities of energy can be released, and thus, only certain wavelengths of light can be released.

The Quantum Model

Bohr’s model did effectively explain the structure of the hydrogen atom with one electron, but failed to rationally explain the structures of atoms with many electrons; he did not take into account the repulsions between the electrons. In the quantum model of the atom, electrons do not “circle” around the nucleus in perfect orbits, as Bohr had stated. Rather electrons exist in regions of space of high probability for their presence called orbitals. According to the Heisenberg uncertainty principle, one cannot perfectly determine the momentum and position of an electron concurrently.

Quantum Numbers

An electron can be described by four quantum numbers, as explained below. According to the Pauli Exclusion Principle, no two electrons can have the same set of four quantum numbers.

The principal quantum number (symbolized as n) desribes the energy level of the electron. The energy of each orbital increases as the distance from the nucleus increases. Sets of orbitals with the same n-value are often referred to as electron shells or energy levels.

The Azimuthal quantum number (or angular momentum quantum number), symbolized as ℓ, is a quantum number for an atomic orbital that determines its orbital angular momentum (essentially the type of orbital of the electron). For any given n, the value of ℓ can range from 0 to n-1, corresponding to the type of orbital. ℓ=0, 1, 2, 3 correspond respectively to s, p, d, and f orbitals. The maximum number of electrons in the various subshells is:

subshell maximum number of electrons
s 2
p 6
d 10
f 14

The magnetic quantum number, symbolized as m, notes which orbital the electron is in. m can retain any value from -ℓ to +ℓ.

The spin quantum number, symbolized as s, is a quantum number that states the intrinsic angular momentum (or simply spin) of a given particle. s can be either +½ or -½.

Magnetism

If an electron has unpaired electrons, it is paramagnetic and be attracted to an electric field. On the other hand, if all of the electrons are paired in an atom, it is diamagnetic and be repelled by an electric field.

 

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