Isotopes of Elements

 

Isotopes of Elements

Definition

The existence of isotopes of elements was first discovered by J.J. Thomson in 1913. The name of isotope was introduced by Soddy because they have the same atomic number and hence occupied the same place in the periodic table. (Isotope is a Greek word; iso = same; topos = place). Nearly all elements found in nature are mixture of several isotopes.

 

“Isotopes are atoms of the same element having SAME ATOMIC NUMBER but DIFFERENT MASS NUMBERS (atomic masses)”. 

OR

“Isotopes are different forms of atoms of an element which have SAME NUMBER OF PROTONS (and also electrons) but DIFFERENT NUMBER OF NEUTRONS in their respective nuclei”.

For example carbon has three of isotopes namely carbon-12, carbon-13 and carbon-14.          

Different isotopes of an element have same chemical properties due to their identical electronic configuration (i.e. same number of electrons in the shells) but they have different physical properties because of their different atomic masses.

Since the proton count establishes elemental identity, chemical traits of different isotopes of the same element tend to be the same. However, isotopes differ in respects of physical properties which depend on atomic mass.

The different number of neutrons, however, affects the stability and mass of the nucleus, sometimes creating a radioactive isotope, sometimes creating a non-radioactive (or stable) isotope.

Same Properties of Isotopes	Different Properties of Isotopes
1. Atomic number (Z)	1. Mass number (A)
2. Electronic configuration	2. Atomic mass
3. Number of electrons	3. Number of neutrons
4. Number of protons	4. Physical properties which depend on mass
5. Chemical properties	5. Nuclear mass/nuclear composition
6. Same position in periodic table	6. Rate of reaction
7. Chemical identity of products of reaction

Uses of Isotopes

1.   As tracers in physical, chemical, biological, medical and metallurgical researches.

2. As therapeutic agent for diagnoses and treatments of various diseases like cancer as radioactive isotopes are easily detected and traced.

3.   Electrical power generation.

Symbolic Representation of Isotopes

The symbol for an isotope is the chemical symbol (or word showing name of element) followed by a dash and then the mass number. So C-14 is the isotope of carbon which contains 6 protons, 6 electrons and 14 – 6 = 8 neutrons. It can also be written as ¹⁴C.

In denoting particular isotopes of an element, the following notation has been internationally adopted. The symbol of the element is written with atomic at the head and atomic number at the bottom. Alternatively, the name of the element is followed by the atomic with a hyphen (-) in between. Thus the isotopes of carbon with atomic number 6 having atomic masses of 12 and 14 may be written as:

¹²C or carbon-12 or C-12

¹⁴C or carbon-14 or C-14

¹²C or carbon-12 or C-12 reads ‘carbon twelve’ meaning isotope of carbon with a mass of approximately 12 amu.

 

Kinds of Isotopes

1.   Natural non-radioactive

2.   Natural Radioactive

3.   Artificial Radioactive (made by neutron bombardment)

Classification of Elements on the Basis 

of Number of Isotopes 

Important Points

1.   Nearly all elements found in nature are mixture of several isotopes.

2.   There are 287 different isotopic species in nature.

3.   Out of 92 elements, 23 elements have no isotopes, each consisting of only one kind of atoms e.g. ₄Be, ₉F, ₁₁Na, ₁₃Al, ₁₅P, ₂₁Sc, ₂₅Mn, ₂₇Co, ₃₃As, ₃₉Y, ₄₁Nb (niobium), ₄₅Rh (rhodium), ₅₃I, ₅₅Cs, ₅₉Pr (praseodymium), ₆₅Tb (terbium), ₆₉Tm (thulium), ₆₇Ho (Holmium), ₇₉Gold (Au), ₈₃Bi, ₈₇Fr, Actinum (₈₉Ac), ₉₀Th, Protactinium (₉₁Pa). The remaining elements have 2 to 10 isotopes each.

NOTE
It is strictly improper to refer to elements that exist in only one atomic form as having “one isotope”; actually such elements like Be, F, Na, Al, P, Sc, Mn, Co, As, Y, Nb, Rh, I, Cs, Pr, Tb, Ho, Tm, Bi, Ac, Th, Pa have no isotopes i.e. they have no other atomic form that is like them in all respects except mass. The term isotopes require the existence of at least two atomic forms of an element; in the same sense that the word twin requires the existence of a pair. Recently the term mono-isotopic is evolved for elements found in nature as a single atomic (isotopic) form.

 

4.  The heavier isotopes of elements usually occur very rarely in the atomic population (e.g. 1 part in 4500 for ²H, 1 part in 140 for U-235; in the exceptional case of chlorine, the ratio of isotopes 35 and 37 is   about 3:1).

5. The isotopes with even atomic number and even atomic masses are more abundant in nature.

6.   Different isotopes of an element occur in different amount or isotopic abundances.

7.  To work out the relative atomic mass of an element, we need to find the average mass of all its   isotopes which involves two steps:

  (i) Multiply each relative isotopic mass by its isotopic abundance and add up the results.

 (ii)  Divide the result by the sum of the abundances. (If the abundances are given as             percentages, this will be 100).

 

Contrary to Dalton’s atomic theory, all atoms of a given element are not necessarily identical. In fact, most elements have been shown to be composed of two or more types of atoms mixed in a fixed proportion.

(i)   The different atoms of such an element contain equal number of protons and, therefore, have the same atomic number.

(ii)  The atoms which vary from one have different number of neutrons in the nucleus. Thus they have different atomic masses.

Applications of Isotopes

Isotopes of Hydrogen

There are three isotopes of hydrogen namely protium or ordinary hydrogen, deuterium or heavy hydrogen, tritium or radioactive hydrogen. Protium is by far the most abundant in natural hydrogen, deuterium about 0.0156 % and tritium only one out of 10, 000, 000 hydrogen atoms.  

The atomic number of three isotopes of hydrogen is 1 while their mass numbers are 1, 2 and 3 respectively. Each of the three isotopes of hydrogen has one proton in the nucleus and one extra-nuclear electron. The nucleus of protium contains one proton only while the number of neutrons in deuterium and tritium is 1 and 2 respectively. 

 

 Isotopes of Carbon

There are two stable isotopes and one radioactive isotope of carbon. The carbon-12 contain 6 proton and 6 neutron, Carbon-13 possess 6 proton and 7 neutron, carbon-14 contain 6 proton and 8 neutron. Carbon 12 is the most abundant (98.889%) isotope

 Isotopes of Oxygen

Oxygen has three isotopes; with the relative abundances of 99.759, 0.037 and 0.204 respectively.

The atomic number of each of these three isotopes of oxygen is 8 while their mass numbers are 16,17 and 18 respectively. Thus each of these isotopes has 8 extra nuclear electrons and 8 protons in the nucleus. The number of neutrons in O-16, O -17 and O -18 is 8, 9 and 10 respectively.