Atomic Number refers to the total number of protons in an atom of an element”. For example, aluminium with atomic number 13 has 13 protons. The number of protons in an atom of an element is very crucial since it is used to distinguish one atom from another. In any given neutral atom, the number of protons is equivalent to the number of electrons because the two have opposite charges which are used to neutralize an atom. It is fair to say, Atomic number= number of protons= number of electrons.
Results from the spectrometer showed the varied values of the RAM of the elements under study. The results contradicted with one of Dalton’s proposals on the atomic theory. Furthermore, the results that some atoms of elements had the same atomic number. These atoms, however, had different mass number due to the difference in the number of neutrons in the nucleus of an element. Isotopes refer to atoms of the same element having the same atomic number but different mass numbers. For example, neon exists naturally as Ne-11, Ne -21 and Ne-22. Isotopes can either exist naturally or artificially.
The atom of different elements that contain the same mass number but different atomic number. Isobars are the atoms or nuclides of 2 or more different chemical elements that have the same number of nucleons (protons+ neutrons). Alfred Walter Stewart has found the isotopes for the first time in the year 1918. In the Greek language- isos means equal and bar means weight. Isobars vary each other. This is because of the difference in atomic numbers but the number of neutrons makes up the difference in the number of nucleons. As they vary in their chemical elements, they exhibit different chemical properties. Some of the examples are
The atomic theory states that every matter composed of discrete, unique components called atoms. Atoms are the reason behind the matter that exists with a certain structure, shape, functions, and behaviour. The concept of atoms began as a philosophical concept in ancient Greece and revolutionized into a science in the early 19th century. Some of the most common theories of atomic models are as follows.
1. Thomson model (Plum pudding model)
2. Dalton model (Billiard ball model)
3. Lewis model (Cubical atom model)
4. Nagaoka model (Saturnian model)
5, Rutherford model (Planetary model)
6, Bohr model (Rutherford–Bohr model
7. Bohr–Sommerfeld model (Refined Bohr model)
8. Gryziński model (Free-fall model)
Thomson’s Atomic Model- Postulates and limitations
An English physicist J.J. Thomson suggested that an atom of an element contains protons which have a positive charge in a spherical model and negatively charged electrons that embed inside the orbital. But, his theory was opposed because it failed to explain how alpha particles scattered using a thin metal foil.
A. According to the postulates of Thomson’s atomic model, an atom resembles a sphere of positive charge with electrons (negatively charged particles) present inside the sphere.
B. Thomson’s model is also known as the plum-pudding model and it is called so because, the plum pudding model has electrons surrounded by a volume of positive charge, like negatively-charged "plums" embedded in a positively-charged "pudding".
C. The model has also been compared to a watermelon as the red-coloured edible part of a watermelon looked like a sphere of an atom having a positive charge and the black seeds around the area looked similar to the electrons inside the sphere.
Rutherford’s model and its limitations
Rutherford’s conducted an experiment by allowing the with α-particles to bombard against a thin sheet of gold. He observed the trajectory of these particles after their interaction with the gold foil. According to Rutherford, an atom of an element is made up of the nucleus which is concentrated in a small volume. The composition of the nucleus includes the positively charged protons and neutrons which don't have the charge. Rutherford assumed that the remaining part of the atom is empty. He further suggested that the electrons in the atom always revolve around the nucleus and both are supported by the electrostatic forces of attraction.
Rutherford has observed that:
A. A major fraction of the α-particles sheet passed through the gold sheet without any deflection, and hence most of the space in an atom is empty.
B. With a small angle, some of the α-particles were deflected by the gold so the positive charge in an atom is not uniformly distributed. The positive charge in an atom is concentrated in a very small volume.
C. Very few of the α-particles were deflected back, that is only a few α-particles had nearly 180o angle of deflection. So the volume occupied by the positively charged particles in an atom is very small as compared to the total volume of an atom
D. Rutherford had proposed that the negatively charged electrons in an atom of an element accelerate as it orbits around the nucleus emitting radiations. However, there was no information given concerning the distribution of negatively charged electrons around the nucleus and also their corresponding energies.
Bohr’s Model and its limitations
The theory was put forth by Niels Bohr in the year 1913. Bohr was Rutherford's product and came up with an advanced model of the structure of the atom. According to Bohr, the single electron in a hydrogen atom is capable of orbiting around the nucleus with uniform energy and from a fixed point which he referred to the radius. Bohr called the path moved by the electrons as energy levels. He further stated that the amount of energy of an electron orbiting around the nucleus is always constant and the atom neither loses nor gains energy during orbiting.
Limitations of Bohr’s Model
A. The theory didn't account for the splitting of the line in the spectrum in the hydrogen atom.
B. The theory was biased since it didn't explain the spectrum of atoms with two or more atoms.
C. It was unable to account for Zeeman effect (splitting of the spectral line in the presence of electronic field)
D. The theory also failed to account for stark effect (splitting of the spectral line in the presence of an electric effect)
E. Bohr didn't explain what happens during the formation of molecules through chemical bonding.