ATOMIC MASSES AND COMPOSITION OF NUCLEUS -

COMPOSITION OF NUCLEUS

After Rutherford’s discovery of the nucleus, the scientists started realising that the nucleus itself must be made of smaller constituent particles. Several hypothesis were put forward about the possible structure of the atomic nucleus. But it was not accepted due to many drawbacks.

The discovery of the proton is credited to Ernest Rutherford, who proved that the nucleus of the H- atom is present in the nuclei of all other atoms in the year of 1917. The H- nucleus was later named ‘proton’ and recognized as one of the building blocks of the atomic nucleus.

Rutherford suggested in early 1920 that the nucleus should contain a neutral particle of approximately the same mass as the proton. This hypothesis was verified in 1932 by James Chadwick who observed emission of neutral radiation when beryllium nuclei were bombarded with α- particles. The only neutral radiation known at that time was photons. This radiation was made incident on paraffin wax, a hydrocarbon having a relatively high hydrogen content. As a result, the high energy proton ejected from the paraffin wax. This show that the neutral radiation/particle must have high energy and momentum(from conservation of momentum and energy). So definitely these neutral particles cannot be photon because photons have no sufficient momentum to eject proton from paraffin wax.

James Chadwick solve this puzzle was to assume that the neutral radiation consists of a new type of neutral particle called neutrons and its mass is very nearly the same as the mass of proton.

$\fn_cm \large m_n=1.00866u=1.6749\times10^{-27}kg$

Chadwick was awarded the Nobel prize in physics for his discovery of the neutron.

Note:

A free neutron, unlike a free proton is unstable. It decays into proton and electron and antineutrino and has a mean life of about 1000 sec.
It is stable inside the nucleus.

Finally, the nucleus has its own structure and consists of protons and neutrons.

PROTON

A proton is a +ve charged particle and has a charge equal to that of an electron and mass is about 1836 times that of the electron.

Charge of proton= $\fn_cm \large 1.6\times10^{-19}C$

Mass of proton= $\fn_cm \large 1.007825 \;a.m.u(u)=1.6726\times10{-27}kg$

NEUTRON

A neutron is a neutral particle i.e it has no charge. Its mass is almost identical to that of the proton.

Charge on neutron= $\fn_cm \large 0$

Mass of neutron= $\fn_cm \large 1.008665\;a.m.u(u)=1.6750\times10^{-27}kg$

These two constituent of a nucleus (proton and neutron) are called nucleons.

Note:

Since neutron is a neutral particle, it has high penetrating power and very low ionising power.
Except Hydrogen, the nuclei of lighter element have equal or nearly equal number of protons and neutrons.
In heavier element, the number of neutrons goes on increasing as compared to the number of protons.

ATOMIC NUMBER

The number of protons inside the nucleus of an atom is known as Atomic number. It is represented by Z. In normal condition atomic number is the same as number of electrons present inside atom.

MASS NUMBER

It is the sum of total number of protons and neutrons inside nucleus of an atom. It is represented by A.

$\fn_cm \large _{A}^{Z}\textrm{X}\rightarrow$ Chemical symbol of the element

$\fn_cm \large Z\rightarrow$ Atomic number (number of protons)

$\fn_cm \large A\rightarrow$ Mass number (Number of Proton+ Number of Neutron)

i.e number of neutrons= $\fn_cm \large =(A-Z)$

$\fn_cm \large ex-$

$\fn_cm \large _{79}^{197}\textrm{Au}\;\;\;p=79,\;\;e=79,\;\;n=(197-79)=118$

ISOTOPES

The isotopes of an element are the atom of the element which have the same atomic number but different mass number.

$\fn_cm \large ex-\;\;\;\;\;\;_{1}^{1}\textrm{H}\;\;\;\;_{1}^{2}\textrm{H}\;(deuterium)\;\;\;\;_{1}^{3}\textrm{H}\;(tritium)$

$\fn_cm \large _{6}^{10}\textrm{C}\;\;\;\;_{6}^{11}\textrm{C}\;\;\;\;_{6}^{12}\textrm{C}\;\;\;\;_{6}^{13}\textrm{C}\;\;\;\;_{6}^{14}\textrm{C}$

$\fn_cm \large _{8}^{16}\textrm{O}\;\;\;\;_{8}^{17}\textrm{O}\;\;\;\;_{8}^{18}\textrm{O}$

isotopes have the same number of protons and electrons but different numbers of neutrons.

Same location in the periodic table
Same chemical properties (because same number and arrangement of electrons)
May not have the same nuclear and physical properties.
It was found that practically every element consists of a mixture of several isotopes. The relative abundance of different isotopes differ from element to element.
Tritium nuclei are unstable, do not occur naturally.

In the periodic table, we give the atomic weight of an element as the average value of the isotopic masses found in nature.

For example. Chlorine has two isotopes of masses/ mass number 35 and 37 which occur in nature in the ratio 3:1

i.e Atomic weight of chlorine $\fn_cm \large =\frac{(35\times3)+(37\times1)}{3+1}=35.5$

Note:

Atomic mass/mass number is a whole number because it is sum of the number of p and n. But Atomic weight may be a fraction because it is the average mass number of all isotopes of particular element.
The element gold has 32 isotopes ranging from A= 173 to 204.

ISOBARS

The atoms of different element which have the same mass number but different atomic number are called isobars.

$\fn_cm \large ex-\;\;\;\;\;\;_{11}^{23}\textrm{Na}\;\;\;\;and\;\;\;\;_{12}^{23}\textrm{Mg}$

$\fn_cm \large _{20}^{40}\textrm{Ca}\;\;\;\;and\;\;\;\;_{18}^{40}\textrm{Ar}$

$\fn_cm \large _{17}^{37}\textrm{Cl}\;\;\;\;and\;\;\;\;_{16}^{37}\textrm{S}$

It have different number of protons, different numbers of electrons and different numbers of neutrons

Chemical properties of isobars are widely different.
Their physical properties may be identical.

ISOTONES

isotones are the nuclei which have the same number of neutrons. It have both atomic number and mass number are different.

$\fn_cm \large ex-\;\;\;\;\;\;_{17}^{37}\textrm{Cl}\;\;\;\;and\;\;\;\; _{19}^{39}\textrm{K}\;\;\;\;(n=20)$

$\fn_cm \large _{8}^{16}\textrm{O}\;\;\;\;and\;\;\;\; _{6}^{14}\textrm{C}\;\;\;\;(n=8)$

$\fn_cm \large _{1}^{3}\textrm{H}\;\;\;\;and\;\;\;\; _{2}^{4}\textrm{He}\;\;\;\;(n=2)$

ATOMIC MASS UNIT (a.m.u)

It is the unit of mass which measure atomic and nuclear masses.

One atomic mass unit (amu) is defined as $(\frac{1}{12})^{th}$ of the mass of an atom of $\fn_cm \large _{6}^{12}\textrm{C}$ isotopes.

$\fn_cm \large \because$ mass of $\fn_cm \large 6.023\times10^{23}$ atoms (1 mole) of =12 gram

$\fn_cm \large \therefore$ mass of 1 atom $\fn_cm \large =\frac{12}{6.023\times10^{23}}\;gram$

By definition

$\fn_cm \large 1\;amu=\frac{1}{12}\times\frac{12}{6.023\times10^{23}}\;gram$

$\fn_cm \large \left [ 1\;amu=1.66\times10^{-27}\;kg \right ]$

This is the same as the average mass of nucleon. a.m.u is represented by u.

i.e mass of electron(e)= $\fn_cm \large 0.00055u$

mass of proton(p)= $\fn_cm \large 1.0073u$

mass of neutron(n)= $\fn_cm \large 1.0086u$

mass of H atom= $\fn_cm \large m_p+m_e=1.0078u$

RELATION BETWEEN a.m.u AND ENERGY (Mev)

We know that from mass-energy equivalence

$\fn_cm \large E=mc^2$

Here $\fn_cm \large m=1\;a.m.u=1.66\times10^{-27}kg$

$\fn_cm \large c=3\times10^{8}m/s$

Equivalence energy is given by

$\fn_cm \large E=mc^2$

$\fn_cm \large =1.66\times10^{-27}\times(3\times10^8)^2$

$\fn_cm \large \left [ E=1.49\times10^{-10}J \right ]$

We know that

$\fn_cm \large 1\;ev=1.6\times10^{-19}J$

$\fn_cm \large 1\;Mev=1.6\times10^{-19}\times10^{6}J=1.6\times10^{-13}J$

$\fn_cm \large E=\frac{1.49\times10^{-10}}{1.6\times10{-13}}Mev$

$\fn_cm \large \therefore \left [ 1 a.m.u\approx 931Mev \right ]{\color{Red} }$

In nuclear physics, the masses are often specified in terms of Mev.