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Structure of atom

Subatomic Particles : Discovery and Characteristics

Subatomic Particles 

Electrons, protons and neutrons are the three main subatomic particles that form an atom.

Discovery of Electron (Michael Faraday’s Cathode Ray Discharge Tube Experiment)

          Experimental Setup: 

Glass tube is partially evacuated (low pressure inside the tube).

Very high voltage is applied across the electrodes. Observation:

      Stream of particles move from the cathode (−ve) to the anode (+ ve). These particles are known as cathode ray particles.

          Results:

Cathode rays move from the cathode to the anode.

Cathode rays are not visible; they can be observed with the help of phosphorescent or fluorescent materials (such as zinc sulphide).

These rays travel in a straight line in the absence of an electric or magnetic field.

The behaviour of cathode rays is similar to that of the negatively charged particles (electrons) in the presence of an electrical or magnetic field.

Characteristics of cathode rays do not depend upon the material of the electrodes and the nature of the gas present in the tube.   Conclusions:

Cathode rays consist of electrons.

Electrons are the basic units of all atoms.

Charge to Mass Ratio of Electrons (J. J. Thomson’s Experiment)

J. J. Thomson measured the ratio of charge (e) to the mass of an electron (me) by using the following apparatus.

He determined  by applying electric and magnetic fields perpendicular to each other as well as to the path of the electrons.

The amount of deviation of the particles from their path in the presence of an electric or magnetic field depends upon: 1. the magnitude of the negative charge on the particle (greater the magnitude on the particle, greater the deflection) 2. the mass of the particle (lighter the particle, greater the deflection) 3. the strength of the electric or magnetic field (stronger the electric or magnetic field, greater the deflection) Observations:

When only electric field is applied, the electrons deviate to point A (as shown in the figure).

When only magnetic field is applied, the electrons strike point C (as shown in the figure).

On balancing the strength of electric and magnetic fields, the electrons hit the screen at point B (as shown in the figure) as in the absence of an electric or magnetic field. Result:

To test your knowledge of this concept, solve the following puzzle.

Charge on Electron (Millikan’s Oil-Drop Experiment)

Millikan's Oil-Drop Apparatus

Atomiser forms oil droplets.

The mass of the droplets is ascertained by calculating their falling rate.

X-ray beam ionises the air.

Oil droplets acquire charge by colliding with gaseous ions on passing through the ionised air.

The falling rate of droplets can be controlled by controlling the voltage across the plate.

Careful observation of the effects of electric field strength on the motion of droplets leads to the conclusion that q = ne. Here, q is the magnitude of electrical charge on the droplets, e is the electrical charge and n is 1, 2, 3,… Results:

Charge on an electron = −1.6022 × 10−19 C

Mass of an electron

Discovery of Proton

Electric discharge carried out in a modified cathode ray tube led to the discovery of particles carrying positive charge; these are known as canal rays.

These positively charged particles depend upon the nature of gas present in them.

The behaviour of these positively charged particles is opposite to that of the electrons or cathode rays in the presence of an electric or magnetic field.

The smallest and lightest positive ion is called a proton (obtained from hydrogen).

Discovery of Neutron

Neutrons are electrically neutral.

They were discovered by Chadwick, by bombarding a thin sheet of beryllium with alpha particles.

The given table lists the properties of these fundamental particles.

Name

Symbol

Absolute Charge/C

Relative Charge

Mass/kg

Mass/u

Approx. Mass/u

Electron

e

−1.6022 × 10−19

−1

9.10939 × 10−31

0.00054

0

Proton

p

+1.6022 × 10−19

+1

1.67262 × 10−27

1.00727

1

Neutron

n

0

0

1.67493 × 10−27

1.00867

1

Macroscopic objects have particle character, so their motion can be described in terms of classical mechanics, based on Newton’s laws of motion.

Microscopic objects, such as electrons, have both wave-like and particle-like behaviour, so they cannot be described in terms of classical mechanics. To do so, a new branch of science called quantum mechanics was developed.

Quantum mechanics was developed independently by Werner Heisenberg and Erwin Schrodinger in 1926.

Quantum mechanics takes into account the dual nature (particle and wave) of matter.

On the basis of quantum mechanics, a new model known as quantum mechanical model was developed.

In the quantum mechanical model, the behaviour of microscopic particles (electrons) in a system (atom) is described by an equation known as Schrodinger equation, which is given below:

Where,

= Mathematical operator known as Hamiltonian operator

ψ = Wave function (amplitude of the electron wave)

E = Total energy of the system (includes all sub-atomic particles such as electrons, nuclei)

The solutions of Schrodinger equation are called wave functions.

Hydrogen atom and Schrodinger equation

After solving Schrodinger equation for hydrogen atom, certain solutions are obtained which are permissible.

Each permitted solution corresponds to a definite energy state, and each definite energy state is called an orbital. In the case of an atom, it is called atomic orbital, and in the case of a molecule, it is called a molecular orbital.

Each orbital is characterised by a set of the following three quantum numbers:

Principal quantum number (n)

Azimuthal quantum number (l)

Magnetic quantum number (ml)

For a multi-electron atom, Schrodinger equation cannot be solved exactly.

Important Features of the Quantum Mechanical Model of an Atom

The energy of electrons in an atom is quantised (i.e., electrons can only have certain specific values of energy).

The existence of quantised electronic energy states is a direct result of the wave-like property of electrons.

The exact position and the exact velocity of an electron in an atom cannot be determined simultaneously (Heisenberg uncertainty principle).

An atomic orbital is represented by the wave function ψ, for an electron in an atom, and is associated with a certain amount of energy.

There can be many orbitals in an atom, but an orbital cannot contain more than two electrons.

The orbital wave function ψ gives all the information about an electron.

|ψ|2 is known as probability density, and from its value at different points within an atom, the probable region for finding an electron around the nucleus can be predicted.

Orbitals and Quantum Numbers

Smaller the size of an orbital, greater is the chance of finding an electron near the nucleus.

Each orbital is characterised by a set of the following three quantum numbers:

The principal quantum number (n)

Positive integers (n = 1, 2, 3,………) Determines the size and energy of the orbital Identifies the shell n = 1, 2, 3, 4, ……..

Shell = K, L, M, N, ……..

With an increase in the value of n, there is an increase in the number of allowed orbitals (n2), the size of an orbital and the energy of an orbital.

The Azimuthal quantum number (l)

Also known as orbital angular momentum or subsidiary quantum number Defines the three-dimensional shape of an orbital For a given value of n, l can have n values, ranging from 0 to n − 1.

For n = 1, l = 0

For n = 2, l = 0, 1

For n = 3, l = 0, 1, 2

For n = 4, l = 0, 1, 2, 3, ………and so on

Each shell consists of one or more sub-she

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