Atomic Particles
* Matter is made up of atoms that further consist of sub-atomic particles.
* Sub-atomic particles are electrons, protons, and neutrons.
* Protons are positively charged, electrons are negatively charged and neutrons are neutral.
* Protons and neutrons have mass equivalent to a Hydrogen atom whereas electrons has a negligible mass.
* Protons and neutrons, together called nucleons, are present in the centre of atom surrounded by revolving electrons and other uncommon sub-atomic particles.
* In 1808, John Dalton proposed the atomic theory of matter. According to Dalton, an atom is the smallest particle of matter.
* The research done by scientists like J.J Thomson, Goldstein, Rutherford, Chadwick and Bohr in the later part of the 19th century and 20th century proved that an atom is not the smallest indivisible particle.
Atomic Models
* According to the Thomson Model of Atom, an atom is a sphere of uniform positive charge with negatively-charged electron embedded in it just like plums in a pudding.
* The gold foil experiment conducted by Rutherford disapproved Thomson’s model.
* Based on the observations from the gold foil experiment, Rutherford concluded that there is a positively charged nucleus at the centre of an atom where the entire positive charge and most of the atomic mass is concentrated.
* Electrons move in circular paths around the nucleus and are held to it by means of electrostatic force of attraction.
* Rutherford’s nuclear model fails to explain the stability of atoms.
Isotopes and Isobars
* Mass number is the sum of the number of protons and the number of neutrons. The atomic number is the number of protons.
* Isotopes have same number of protons but different number of neutron.
* As the isotopes have the same number of electrons and proton, they have similar chemical properties.
* Isobars are atoms of different elements having different atomic numbers but same mass numbers.
* Isobars have different chemical properties because they have different atomic numbers.
Theory of Electromagnetic Radiation and Planck’s Quantum Theory
* Line spectra could not be explained using the Rutherford’s nuclear model.
* In 1864, James Clark Maxwell established the wave nature of electromagnetic radiation.
* Characteristics of waves are associated with wavelength, frequency, velocity and wave number.
* Electromagnetic radiations travel with the same velocity but differ from one another in wavelength and therefore, in frequency. Their wavelengths increase from gamma rays to x-rays to ultra violet rays, visible rays, infra red rays, microwaves to radio waves.
* When these electromagnetic waves of different wavelengths are arranged in the order of increasing wavelength or decreasing frequencies, an electromagnetic spectrum is formed.
* Wave nature of electromagnetic radiation explains interference and diffraction but could not explain black body radiation and photoelectric effect.
* In 1900, Max Planck established the Planck’s Quantum Theory that states that radiant energy, such as light and heat, is propagated in the form of small packets called Quanta. The amount of energy associated with the quantum of energy is proportional to the frequency of radiation, E a v or E = hv.
Photoelectric Effect and Dual Behaviour of Electromagnetic Radiation
* In 1887, Hertz stated that when light falls on certain metals like potassium and sodium, electrons were ejected. This is called the photoelectric effect.
* Einstein explained the photoelectric effect by applying the Planck’s quantum theory of radiation.
* When a beam of light, that is, a beam of photons strikes the surface of the metal, electrons will be ejected only if the incident photons have the frequency nu greater than the threshold frequency nu nought.
* Greater is the intensity of incident light or the number of photons that strike a metal, greater is the number of electrons ejected.
* Greater is the frequency of the incident light, greater is the kinetic energy of the ejected electron.
Atomic Spectra
* The splitting of light into a series of colour bands is known as dispersion and the series of colour bands is called a spectrum.
* The spectrum of radiation emitted by a substance that has absorbed energy is called emission spectrum.
* An absorption spectrum is recorded by passing a continuum of radiation through a sample which absorbs radiation of certain wavelengths. The study of emission or absorption spectra is referred as spectroscopy.
* As the emission spectrum of each atom is unique, spectroscopy is often used to identify elements.
* The simplest emission spectrum is that of hydrogen atom.
* The hydrogen spectrum consists of multiple series of lines including Lyman, Balmer, Paschen, Brackett and Pfund series.
* The wavelengths of the series of lines in the spectrum of hydrogen can be measured by using the Rydberg’s formula.
Bohr’s Model for Hydrogen Atom
* Bohr’s model of Hydrogen atom was given by Neils Bohr in 1913.
* The first postulate of Bohr’s model states that the electron in the hydrogen atom moves around the nucleus in a fixed circular orbit of a particular radius and energy.
* The second postulate of Bohr’s model states that the energy of an electron in the orbit remains constant until the electron jumps from one orbit to another.
* The third postulate of Bohr’s model states that the frequency of the radiation emitted or absorbed by an exciting electron can be represented by the Bohr’s frequency rule.
* The fourth postulate of Bohr’s model states that any moving body that takes a circular orbit has an angular momentum equal to the product of its mass, linear velocity and radius of orbit.
* Bohr’s theory states that the stationary states for electrons can be represented by positive Integral numbers starting from 1, and are called Principle quantum numbers.
* The radius of each stationary state of a hydrogen atom can be calculated using the equation 52.9 multiplied by n squarepm.
* The energy of each stationary state of a hydrogen atom can be calculated using the equation En is equal to minus RH multiplied by (one divided by square), where n is the principal quantum number of the orbit and RH is Rydberg’s constant.
* The magnitude of velocity of electron increases with increase of positive charge on the nucleus and decreases with increase of principal quantum number.
* Bohr’s model explains the stability of an atom as well as the line spectrum of hydrogen atom. Thereby overcoming the limitations of Rutherford’s atomic model.
* Although Bohr’s model helps find answers to many questions, it has its set of limitations.
Development Leading to Quantum Mechanical model of The Atom
* The dual behavior of matter was proposed by de Broglie in 1924.
* According to de Broglie, the relation between the wavelength and momentum of a material particle is given by the equation: lambda is equal to h divided by p.
* Based on the dual behavior of matter and radiation, Werner Heisenberg gave uncertainty principle in 1927.
* It is impossible to measure simultaneously and accurately the exact position and momentum of a small particle like an electron.
* The uncertainty in the position and velocity decreases as the mass of the object increases.
* As the uncertainty in position reduces the uncertainty in velocity increases and vice versa.
* Because of the inverse relationship between the uncertainty in position and velocity and the mass of the particle, Heisenberg’s uncertainty principle only affects the moving microscopic objects.
Bohr’s model not only ignores dual nature of matter but also contradicts the Heisenberg’s Uncertainty Principle.
Quantum Mechanical Model of the Atom
* Quantum mechanics was developed by Heisenberg and Ervin Schrodinger in 1926.
* Schrodinger wave equation when solved for hydrogen or hydrogen-like species gives the values of quantised energy levels and their wave functions.
* The square of wave function gives the probability density.
* Finding the probability density at different points in an atom helps to determine the atomic orbitals.
* Atomic orbitals are distinguished from each other based on shape, size and orientation that are expressed in terms of three quantum numbers: principal, azimuthal and magnetic quantum numbers.
* The principal quantum number helps to identify the shells in which the electrons reside, the number of electrons in shells, the energies of shells, and the size of orbitals.
* The azimuthal quantum number tells about the number of subshells in a shell, their energies and the shape of the orbitals.
* The magnetic quantum number helps to find the number of allowed orientations and the number of orbitals present in a subshell.
* A fourth quantum number, called spin quantum number represents the spin of the electron about its axis, which in turn, helps to determine the magnetic behavior of an atom.
Atomic Orbitals: Shapes and Energies
* The shape of an orbital is depicted by its probability density function.
* The region where the probability density function reduces to zero is called nodes, the total number of nodes for a orbital is given by the expression (n-1).
* Probability density of electrons can be depicted through boundary surface diagrams, in which case the shape of the diagram gives the shape of the orbital.
* The shape of s-orbital is symmetric spherical with only one possible orientation.
* The shape of p-orbitals is dumb-bell with three possible orientations along the three axis.
* Four out of five d – orbitals have the same shape, but different orientations; the fifth orbital has a different shape and orientation.
* For single-electron atoms like hydrogen, the energies of all the orbitals in a shell is the same and the size and energy of orbitals increase with increase in the value of n.
* The sub-shells belonging to a principal shell of multi-electron species have different energies owing to the mutual repulsion among the electrons present in these sub-shells.
* s – orbitals have lower energies than p – orbitals that in turn have lower energies than d – orbitals.
* The energies of the orbitals in the same sub-shell decrease with increase in the atomic number.
* The lower the value of (n + l ) for an orbital, the lower is its energy.
* If two orbitals have the same value of (n + l ), the orbital with lower value of n will have the lower energy.
Atomic Orbitals: Filling of Electrons
* According to the Aufbau principle, the filling up of electrons into various orbitals takes place according to the increasing order of their energy.
* According to the (n + ) rule, the orbital with the lower value of (n + ) has the lower energy and hence is filled up first.
* Pauli’s exclusion principle states that an orbital can have maximum two electrons and these must have opposite spins.
* The Hund’s rule of maximum multiplicity states that in a set of degenerate orbitals, the orbitals are singly occupied first and then pairing of electrons takes place.
Electronic Configuration
* The three main principles that guide us on the filling of orbitals in an atom are the Aufbau Principle, Pauli’s Exclusion Principle, and the Hund’s Rule of Maximum Multiplicity.
* Electronic Configuration can be defined as the distribution of electrons into various orbitals of an atom.
* We can represent the electronic configuration of atoms in two ways.
• Orbital diagram method
• (nlx) method
* We can also write the electronic configurations of elements whose atomic numbers are between 11 and 17, in terms of the electronic configuration of Neon.
* In some cases the actual configuration differs slightly from the expected ones – example, chromium and copper. This is because of the extra stability of half-filled and completely filled sub shell configuration.
* The cause of this extra stability has been attributed to the symmetry distribution and exchange energy.