Atomic Number, Relative Atomic Masses, Isotopes & Calculations

What are atoms?

Atoms are the fundamental building blocks of matter and are composed of three subatomic particles: protons, electrons, and neutrons. These particles contribute to the overall structure, charge, and properties of an atom.

1. Protons: Protons are positively charged particles found within the nucleus of an atom. They have a relative charge of +1 and a relative mass of approximately 1 atomic mass unit (AMU). Protons play a crucial role in determining the identity of an atom. The number of protons in the nucleus defines the atomic number of an element, which distinguishes one element from another. For example, all hydrogen atoms have one proton, while all carbon atoms have six protons.

2. Electrons: Electrons are negatively charged particles that orbit the nucleus of an atom in specific energy levels called electron shells or orbitals. Electrons have a relative charge of -1 and a negligible mass compared to protons and neutrons. The number and arrangement of electrons determine the chemical behavior and reactivity of an atom. Electrons occupy different energy levels, with the innermost shell closest to the nucleus and subsequent shells further away. Each electron shell has a maximum capacity of electrons, with the first shell accommodating a maximum of two electrons, the second shell accommodating up to eight electrons, and so on.

3. Neutrons: Neutrons are neutral particles found within the nucleus of an atom. They have no charge, meaning they are electrically neutral. Neutrons have a mass similar to protons, approximately 1 AMU. The number of neutrons in an atom can vary, resulting in different isotopes of an element. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. The presence of neutrons affects the stability and nuclear properties of an atom.

The combination of protons, neutrons, and electrons determines the overall charge and mass of an atom. Since protons have a positive charge and electrons have a negative charge, atoms are electrically neutral when the number of protons equals the number of electrons. The total number of protons and neutrons in the nucleus is called the mass number, while the difference between the number of protons and the number of electrons is the atom’s overall charge.

The understanding of these subatomic particles provides the foundation for atomic theory and helps explain various phenomena such as atomic structure, chemical bonding, and nuclear reactions. The interactions between protons, electrons, and neutrons govern the behavior and properties of matter at the atomic level.

Atomic number and mass number

The atomic number of an element is the number of protons in the nucleus of its atom. The mass number or atomic mass of an element is the sum of the number of protons and neutrons in the nucleus of its atom.

Mass Number = Number of proton + Number of neutron

An element X can be represented as

ZXA

where               A= Atomic mass or mass number

Z= Atomic number

e.g.      40        mass no = 40

Ca

  • atomic no = 20

ISOTOPES

Isotopy is a phenomenon observed in atoms of elements where they have the same atomic number (number of protons) but different mass numbers (sum of protons and neutrons). This variation in mass numbers is caused by differences in the number of neutrons present within the atomic nucleus. Atoms that exhibit isotopy are known as isotopes.

Isotopes of an element possess similar chemical properties and undergo the same chemical reactions because their electron configurations remain the same. This is because the chemical behavior of an atom is primarily determined by its electron configuration, which is dictated by the number of protons and electrons. However, isotopes may have slightly different physical properties, such as variations in their atomic mass and nuclear stability.

For example, carbon has three naturally occurring isotopes: carbon-12 (12C), carbon-13 (13C), and carbon-14 (14C). All three isotopes have six protons, as they are all carbon atoms. However, they differ in the number of neutrons they possess. Carbon-12 has six neutrons, carbon-13 has seven neutrons, and carbon-14 has eight neutrons. These isotopes are denoted by their mass numbers, which are the sum of protons and neutrons.

Isotopes can be stable or unstable, depending on the balance between the forces within the atomic nucleus. Stable isotopes, such as carbon-12 and carbon-13, have a ratio of protons to neutrons that allows them to exist indefinitely without undergoing radioactive decay. On the other hand, unstable isotopes, also known as radioactive isotopes, have an imbalance in their nuclear composition, leading to the eventual decay of the nucleus over time. Carbon-14 is a radioactive isotope that undergoes radioactive decay with a half-life of approximately 5,730 years.

Isotopes find various applications in science, medicine, and industry. In fields like archaeology and geology, the ratios of isotopes in samples can be used to determine the age of artifacts or study geological processes. In medicine, isotopes are utilized in imaging techniques such as positron emission tomography (PET) scans. Isotopes also play a crucial role in nuclear power generation, where certain isotopes are used as fuel in nuclear reactors.

Overall, the existence of isotopes expands our understanding of atomic structure, contributes to research in various fields, and highlights the diverse nature of elements, allowing for a more nuanced understanding of the physical world.

Examples of atoms that exhibit isotopy are chlorine 35Cl and 37 Cl

Carbon-  12 C, 13 C  and 14 C

Potassium – 39K19   and   41K19

Oxygen –  16O16    and 18O16

CALCULATION OF RELATIVE ATOMIC MASS

The following is an example of the calculation of the relative atomic mass of an element from the percentage abundance of its isotopes.

  1. X is an element that exists as an isotopic mixture containing 90% of 39X19 and 10% of 41X19
  2. How many neutrons are present in the isotope 41X
  3. Calculate the mean relative atomic mass of X

Solution

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  1. Neutrons in 41X19

=          41-19 = 22

  1. R.A.M =          90        x          39        +          10        x          41
  • 100

=          90        x          39        +          41        x          10

100

=          3920    =          39.20

100

CALCULATIONS

  1. The following are more examples of calculations of the relative atomic masses of elements.
  2. An element Y exists in two isotopic forms 39Y18 and 40Y18 in the ratio 3:2 respectively. What is the relative atomic mass of the element?

SOLUTION

R.A.M of Y        =          3          x          39        +          2          x          40

5                      1                     5                       1

=         0.6        x         39         +         0.4        x          40

=        23.4       +       16

=       39.4

3. An element with a relative atomic mass of 16.2 contains two isotopes 16P8 with a relative abundance of 90% and mP8 with a relative abundance of 10%. What is the value of m?

SOLUTION

16.2 =  90 x 16 + 10 x m

100

16.2 = 9 x 16  +   m

10           10

16.2 = 144    m

10         10

16.2 = 144  +  m

10

16.2 x 10 = 144 + m

162   = 144 + m

162 – 144 = m

18 =   m

The value of m is 18

Read also:

IUPAC Nomenclature of Chemical Compounds

PARTICULATE NATURE OF MATTER

STANDARD SEPARATION TECHNIQUES

COMPOUNDS AND MIXTURES

ELEMENTS, SYMBOLS & VALENCY

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