Akarshik BanerjeeClass- X l ’B’

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In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron.

Discovery Of Electron

Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electronsElectrons have so little mass that atoms must contain other particles that account for most of the mass

Conclusions Of Discovery Of

Electron

Eugen Goldstein in 1886 observed what is now called the “proton” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron)

1932 – James Chadwick confirmed the existence of the “neutron” – a particle with no charge, but a mass nearly equal to a proton

1916 – Robert Millikan determines the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge

The oil drop apparatus

Mass of ElectronMass of electron is 9.3 x 10-31

Particle Charge Mass (g) Location

Electron (e-) -1 9.11 x 10-28 Electron

cloud

Proton (p+) +1 1.67 x 10-24 Nucleus

Neutron (no) 0 1.67 x 10-24 Nucleus

Subatomic Particles

Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

THOMSON’S ATOMIC MODEL

Thomson’s model was cancelled out as it could not explain the experiments of Rutherford

Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil Particles that hit on the detecting screen (film) are recorded

Rutherford’s Experiment

Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected

His findings

The nucleus is smallThe nucleus is denseThe nucleus is positively charged

Conclusions

Based on his experimental evidence:The atom is mostly empty spaceAll the positive charge, and almost all

the mass is concentrated in a small area in the center. He called this a “nucleus”

The nucleus is composed of protons and neutrons (they make the nucleus!)

The electrons distributed around the nucleus, and occupy most of the volumeHis model was called a “nuclear model”

If the electrons revolved around nucleus in fixed orbit then it would constantly loose energy and finally collapse in the nucleus as per law of electrodynamics. Then the atoms would be highly unstable , but that is not so.. Thus Rutherford’s model was rejected.

Drawbacks:-

Bohr’s Atomic Model

Electrons revolves around the nucleus in fixed circular orbits .

During revolution there in no loss or gain of energy. Energy change takes place only when electron jumps.

In an atom only those orbits are possible in which the angular momontum of an electron is an integral multiple of nh/2 .pDrawbacks1. It was applicable to atoms containing only one electron.2. It could not explain the splitting of spectral lines under

electrical and magnetic field.3. It could not explain Heisenberg’s uncertainty principle.4. It could not explain formation of molecule through bonding.

ShellsThe principal energy levels found around the nucleus where electrons revolve.

Sub-shellsThese are the sub energy levels found within the shells.

Orbitals These are the space around the nucleus where probability of finding electron is maximum.Auf-Bau principle

It states that electrons occupy shells subshells and orbitals in increasing order of energy.

The half filled and full filled orbitals are more stable than the rest orbitals.

Hund’s RuleIt states that when two or more subshells of same energy level are available then electron is first filled one by one in them and then pairing takes place in opposite spin.

Quantum numbers describe values of conserved quantities in the dynamics of the quantum system. Perhaps the most peculiar aspect of quantum mechanics is the quantization of observable quantities. This is distinguished from classical mechanics where the values can range continuously. They often describe specifically the energies of electrons in atoms, but other possibilities include angular momentum, spin etc. Any quantum system can have one or more quantum numbers, it is thus rigorous to list all possible quantum numbers.

Quantum Numbers

•The first, describes the electron shell, or energy level.(Principal quantum number)

•The value of ranges from 1 to "n", where "n" is the shell containing the outermost electron of that atom. For example, in cesium (Cs), the outermost valence electron is in the shell with energy level 6, so an electron in cesium can have an value from 1 to 6.

•The second, describes the subshell (0 = s orbital, 1 = p orbital, 2 = d orbital, 3 = f orbital, etc.).(Azimuthal qn.)

•The value of ranges from 0 to . This is because the first p orbital (l=1) appears in the second electron shell (n=2), the first d orbital (l=2) appears in the third shell (n=3), and so on. A quantum number beginning in 3,0,... describes an electron in the s orbital of the third electron shell of an atom.

•The third, , describes the specific orbital (magnetic qn)(or "cloud") within that subshell.

•The values of range from to . The s subshell (l=0) contains only one orbital, and therefore the ml of an electron in an s subshell will always be 0. The p subshell (l=1) contains three orbitals (in some systems, depicted as three "dumbbell-shaped" clouds), so the ml of an electron in a p subshell will be -1, 0, or 1. The d subshell (l=2) contains five orbitals, with ml values of -2,-1,0,1, and 2.

•The fourth, , describes the spin of the electron within that orbital. (spin qn)

•Because an orbital never contains more than two electrons, will be either or , corresponding with "spin" and "opposite spin". It states that no two electrons of an atom can have same set of all four quantum numbers.

Pauli’s Exclusion Principle

Radiations have both wave like and particle like nature. The particle like nature can be proved by the following effects-1) Formation of Scintillations.

2) Photoelectric Effect

When monochromatic beam of radition is allowed to pass through a nicol prism and the radition is aloowed to fall at a ZnS plate then formation of Scntillations confirms then dual behaviour of radiation.

Einstine discovcred that it rasiation of a minimum Threshold frequency is allowed to fall on the surface of an active metal then ejection of electrons from the surface takes place called photoelectrons . The phenomenon is called photoelectric effect. Thus light particles had momentum and also mass.

Dual Nature of Light

The work function is the minimum energy that must be given to an electron to liberate it from the surface of a particular substance. In the photoelectric effect, electron excitation is achieved by absorption of a photon. If the photon's energy is greater than the substance's work function, photoelectric emission occurs and the electron is liberated from the surface.

Work Function

Dual Nature of Matter

Matter also show dual behavior. The wave like nature of electron can be proved through the electron diffraction experiment.

In quantum mechanics, the Heisenberg uncertainty principle states by precise inequalities that certain pairs of physical properties, such as position and momentum, cannot be simultaneously known to arbitrarily high precision. That is, the more precisely one property is measured, the less precisely the other can be measured.

Heisenberg’s Uncertainty Principle

According to Heisenberg's uncertainty principle, it is impossible to describe the exact position of an electron at a given moment in terms of position, we can speak of most probable regions where the probability of finding an electron in the space around the nucleus of an atom is high. The electron does not always remain at a fixed distance from a nucleus. It keeps moving in the whole space around the nucleus but tends to remain most of the time within a small volume around the nucleus, where the probability of locating the electron is maximum.A new atomic model, was needed to explain •Wave nature (dual character) of atoms. •The idea of uncertainty in the position of electrons in a atom. •Concept of fixed energy states. Schrodinger put the wave model or quantum mechanical model of atom forward. The behavior of an electron is defined by the mathematical representation:                                                                                                       

 

where, y = (psi) is a wave function of space coordinates 'x', 'y', 'z' and represents the amplitude of the electron wave.

m = mass of the electron E = the total permissible energy level, which the electron can have.

V = potential energy of the electron given by ze2/r. h = Planck's constant having the value 6.626 x 10-34 J s.d= (delta)stands for infinitesimal change. The wave length function y (psi) describes a number of possible states of an electron in an atom. Since a large number of solutions are possible, four quantum numbers were introduced, which describe meaningful permissible values of energy and location with respect to its nucleus.

Quantum mechanical model of atom

Electron cloud Diagram

Nodal plane

According to probability density the shapes of electron clouds can be drawn.The space around the nucleus where the probability of finding an electron is zero is called a node.

Nodal PlaneThe plane containing the node is called the nodal plane.

Rydberg’s Equation

Lyman seriesIt is the series of emission spectra obtained when an electron jumps from any higher shell to the first shell.it falls under UV region and value of n1=1.

Balmer seriesIt is the series of emission spectra obtained when an electron jumps from any higher shell to the second shell.it falls under visible region and value of n1=2.

Paschan seriesIt is the series of emission spectra obtained when an electron jumps from any higher shell to the third shell. It falls under IR region and value of n1=3.

Brackett seriesIt is the series of emission spectra obtained when an electron jumps from any higher shell to the fourth shell. It falls under IR region and value of n1=4.

Pfund SeriesIt is the series of emission spectra obtained when an electron jumps from any higher shell to the fifth shell. It falls under IR region and value of n1=5.