Niels Bohr, who was interested in explaining the spectrum of discrete lines that is observed when light is emitted by different elements. Many models of the atom have been proposed based on experimental data, including J.J. Thomson's discovery of the electron and Ernest Rutherford's discovery of the nucleus. Rutherford explained that in an atom, the nucleus is positively charged and surrounded by electrons in his Rutherford Model (negatively charged particles). The electrons follow a predetermined course known as orbits. Rutherford's concept, in which electrons travel around in fixed orbital shells, was amended by Bohr.
He also stated that each orbital shell has a set of energy levels. As a result, Rutherford essentially explained an atom's nucleus, whereas Bohr advanced the model. He went over electrons and the various energy levels linked with them. He came to the conclusion that electrons will have more energy if they are placed away from the nucleus, while electrons will have less energy if they are located near the nucleus.
Bohr’s model of the hydrogen atom was the very first model of atomic structure that explained the spectrum of hydrogen model. Bohr’s theory explains the emission and absorption spectra of atomic hydrogen and hydrogen-like ions with low atomic number. Bohr’s model of hydrogen atom is based on three postulates (Lee. J. D (1996)):
- An electron moves around the nucleus in a circular orbit.
- An electron’s angular momentum of an orbit is quantized and
- The change in the energy of electron makes to jump from one orbit to another accompanied by either by radiation and absorption of energy.
The Bohr model was important step in development of atomic theory but there arises several problems or limitation with his concept.
Limitations of the Bohr model
- It cannot applied be applied to multielectron atom even for helium atom.
- It violets the Heisenberg Uncertainity Principle as in Bohr’s atomic model theory consider electrons to have both known radius and orbit i.e know the position and momentum at the same time, which is impossible according to Heisenberg.
- It failed to explain to account for the effect of magnetic field (Zeeman Effect) or electric field (Stark Effect) when the spectral lines split into several
- components which could not be explained on the basis of Bohr’s model.
Development of atomic theory
Modern atomic theory
Neil Bohr, a Rutherford student, created a new model of the atom in 1913. Electrons are said to be grouped in concentric circular orbits around the nucleus, according to him. This model, known as the planetary model, is based on the solar system. . The four principles that make up the Bohr model are as follows:
- Only a few orbits around the nucleus are occupied by electrons. Those orbits are known as 'stationary' orbits because they are stable.
- Each orbit has a certain amount of energy attached to it. E1 is the energy of the orbit closest to the nucleus, E2 is the energy of the orbit next to it, and so on.
- When an electron leaps from a lower orbit to a higher one, energy is absorbed, and when an electron falls from a higher orbit to a lower one, energy is emitted
- The difference between the two orbital energies can be used to compute the energy and frequency of light emitted or absorbed.
Erwin Schrodinger, an Austrian scientist, advanced the Bohr atom model in 1926. To describe the possibility of detecting an electron in a specific position, Schrödinger employed mathematical equations. The quantum mechanical model of the atom is the name given to this atomic model. Unlike the Bohr model, the quantum mechanical model does not specify an electron's exact path; rather, it forecasts the probability of the electron's position. The nucleus in this model is surrounded by an electron cloud. The probability of locating the electron is greatest where the cloud is most dense, and the electron is less likely to be in a less dense portion of the cloud. The probability of locating the electron is greatest where the cloud is most dense, and the electron is less likely to be in a less dense portion of the cloud. As a result, the concept of sub-energy levels was incorporated in this model.
The atom was thought to be made up of a positively charged nucleus surrounded by negatively charged electrons until 1932. James Chadwick injected alpha particles into beryllium atoms in 1932. There was a radiation that was not identified. This radiation, according to Chadwick, is made up of particles with a neutral electrical charge and a proton-like mass. The neutron was created from this particle. Chemists now have a good model of the atom, thanks to the discovery of the neutron. Many more particles have been discovered in the atom since 1932 as a result of continuing experimentation. Existing nuclei have also been bombarded with various subatomic particles to produce new elements. The idea that protons and neutrons are made up of even smaller units called quarks has added to the atomic theory. The quarks themselves are made up of energy strings that vibrate. The study of the atom's composition is still an intriguing and continuous endeavor.
Periodic properties of elements – Reasons, significance and measurements.
The elements of the periodic table are arranged in order of increasing atomic number. All of these elements exhibit several different trends, and we can use the periodic law and tabulation to predict their chemical, physical, and atomic properties.
- 1. Atomic radius: is the distance from atomic nucleus and to the valence electron in an atom.
Variation in period: The atomic radius tends to decrease from left to right over a period due to the contraction of the atom because the effective nuclear force increases on the electrons.
Variation in group: The atomic radius generally increases while going down the group due to the addition of a new energy level or shell (shell that causes the size of atoms to shrink across the period). However, atomic radii tend to increase diagonally because the number of electrons has a greater effect than that of the bulky nucleus. For example, lithium (145 picometers) has a smaller atomic radius than magnesium (150 picometers) (Chem.libretexts.org).
Significant: Many aspects of chemistry, such as physical and chemical properties, can be determined using atomic radii. The periodic table is extremely useful in estimating atomic radius and displays a lot of patterns (internet, Calbreath, Baxter.’ Chemistry for non-major’)
- 2. Ionization energy: it is the amount of energy required to remove electron from outer most shell of a gaseous atom. Greater the ionisation energy, the more difficult it is to remove electron and vice versa. It is measured in joules and electron volts.
Variation in period: ionization energy increases because the atomic radius decreases due addition of electron in the same shell thereby increase in effective nuclear charge and it becomes harder to remove electron.
Variation in group: the ionization energy decreases because of addition of extra shell results valence electron to farther away from the nucleus which experiences a weak force of attraction.
Significant: ionization is significant because it may be used to forecast or predict the strength of chemical bonds.
- 3. Electron affinity: it is the amount of energy release due to the addition of extra electron in (lee, 1996)the outer most shell of gaseous atom. The nonmetals have more positive electron affinity than metals. Atoms, such as Group 7 elements, whose anions are more stable than neutral atoms have a higher. The electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative values. Chlorine has the highest electron affinity while mercury has the lowest. However it is expressed in units of KJmol.
Variation in period: it increases across a period due to the filling of the valence shell of the atom, the nuclear force also increases hence the electron gain enthalpy increases. For example, within the same period, a Group-17 atom releases more energy than a Group-1 atom upon gaining an electron because due to the addition of the electron it creates a filled valence shell and therefore is more stable.
• Variation in group: the electron affinity decreases as the electron is being added increasingly the atomic size causing a decrease in the electron gain enthalpy. Less tightly bound and therefore closer in energy to a free electron.
Electron affinity= 1Atomic Size
- 4. Electronegativity: it is tendency of an atom to accept the electron in a compound.
Variation in period: As we move across the period, the nuclear charge increases and the atomic size decreases as a result the value of electronegativity increases from left to right in a period. Generally, metals show a lower value of electronegativity as compared to non-metals.
Variation in group: as we move down the group, the atomic number thus, nuclear charge also increases but the effect of the increase in nuclear charge is overcome by the addition of one shell. Therefore, the value of electronegativity decreases as we move down the group. For example, in the halogen group as we move down the group from F>Cl>Br>I>At, the electronegativity value decreases.
Significant: electronegativity is important as it determines the nature of the bond.
- 5. Metallic and non-metallic properties: Metals loses electrons to form cations, whereas nonmetals gains electrons to form anions.
Metallic properties increases as we move down the group because the nuclear force of attraction decreases between nuclei and outermost electron which cause the outermost electrons to be loosely bound and thus able toto conduct heat and electricity.
Across the period, from left to right, the increasing attraction between the nuclei and outer most electron causes the metallic character to decreases.
Non-metallic property increases as we move from left to right across a period and decreases down the group due to the same reason due to an increase in nuclear attractive force. Therefore, metals are ductile and nonmetals are not.